The Basics of Chemistrys2.bitdl.ir/Ebook/Chemistry/Myers - The Basics of Chemistry (Greenwood... · Colligative Properties 131 Reactions of Solutions 134 Colloids and Suspensions

The Basics of Chemistry

Richard Myers

Basics of the Hard SciencesRobert E. Krebs, Series Editor

GREENWOOD PRESSWestport, Connecticut • London

Library of Congress Cataloging-in-Publication Data

Myers, Richard, 1951–The Basics of chemistry / Richard Myers.

p. cm.––(Basics of the hard sciences)Includes bibliographical references and index.ISBN 0–313–31664–3 (alk. paper)

1. Chemistry. I. Title. II. Series.QD33.2.2.M94 2003540––dc21 2002028436

British Library Cataloging in Publication Data is available.

Copyright © 2003 by Richard Myers

All rights reserved. No portion of this book may be reproduced, by any process or technique, without the express written consent of the publisher.

Library of Congress Catalog Card Number: 2002028436ISBN: 0–313–31664–3

First published in 2003

Greenwood Press, 88 Post Road West, Westport, CT 06881An imprint of Greenwood Publishing Group, Inc.www.greenwood.com

Printed in the United States of America

The paper used in this book complies with the Permanent Paper Standard issued by the National Information Standards Organization (Z39.48–1984).

10 9 8 7 6 5 4 3 2 1

Every reasonable effort has been made to trace the owners of copyrighted materials in this book, but in someinstances this has proven impossible. The author and publisher will be glad to receive information leading tomore complete acknowledgments in subsequent printings of the book and in the meantime extend theirapologies for any omissions.

To Chris.For all your support and companionship during our thirty years.

Contents

Preface xiii

Acknowledgments xv

Chapter 1 Introduction: Chemistry and Its Divisions 1Introduction 1The Scientific Enterprise 2What Is Chemistry? 3Divisions of Chemistry 4

Chapter 2 A Brief History of Chemistry 7Introduction 7Early History 7Greek Science 9Alchemy 11Chemistry in the Middle Ages 13

Chapter 3 The Birth of Modern Chemistry 17Introduction 17The Beginnings of Modern Chemistry 17The Phlogiston Theory 19Pneumatic Chemistry 20Lavoisier 25

Chapter 4 The Atom 31Introduction 31The Law of Definite Proportions 31Dalton and the Birth of the Atomic Theory 33Gay-Lussac and Avogadro’s Hypothesis 34The Divisible Atom 35Atomic Structure 38

Quantum Mechanics 40A Modern View of the Atom 44Atomic Mass, Atomic Numbers, and Isotopes 45Ions 45Summary 46

Chapter 5 Atoms Combined 49Introduction 49Mixtures, Compounds, and Molecules 49Naming Elements 50Naming Compounds 51Writing Chemical Formulas 54Chemical Reactions 54Types of Reactions 55From Molecules to Moles 56

Chapter 6 Elements and the Periodic Table 61Introduction 61Discovery of the Elements 61The Modern Periodic Table 64The Main Group Elements 64d and f Block Elements 67Metals, Metalloids, and Nonmetals 67Periodic Problems 67Summary 69

Chapter 7 Chemical Bonding 71Introduction and History 71Dot Formulas, the Octet Rule, and Ionic Bonds 74Covalent Bonds 76Electronegativity and the Polar Covalent Bond 76Polar Molecules and Hydrogen Bonds 78Bond Energies 79The Metallic Bond 80Bonding and Molecular Geometry 80Modern Bonding Theory 82

Chapter 8 Intermolecular Forces and the Solid and Liquid States 85Introduction 85Characteristics of Solids, Liquids, and Gases 86Intermolecular Forces 86Crystalline Solids 90Amorphous Solids 92Liquids 93Material Science 95

viii Contents

Contents ix

Chapter 9 Gases 99Introduction 99Pressure 99The Gas Laws 102The Ideal Gas Law 106Partial Pressure and Vapor Pressure 107Applications of the Gas Laws 109

Chapter 10 Phase Changes and Thermochemistry 113Introduction 113Changes of State 113Energy and Phase Changes 113Heating Curves 115Calorimetry 118Heat of Reaction 120Applications of Heat of Reactions 122

Chapter 11 Solutions 125Introduction 125The Solution Process 126Electrolytes 128Concentration of Solutions 129Solubility of Gases 130Colligative Properties 131Reactions of Solutions 134Colloids and Suspensions 136

Chapter 12 Kinetics and Equilibrium 139Introduction 139Reaction Rates 139The Collision Theory 140Factors Affecting Reaction Rates 142Reaction Mechanisms and Catalysis 144Chemical Equilibrium 147Le Châtelier’s Principle 149The Equilibrium Constant 151Ammonia and the Haber Process 152

Chapter 13 Acids and Bases 155Introduction and History 155Chemical Definitions of Acids and Bases 156Strengths of Acids and Bases 159Acid Concentration 160pH and pOH 161Neutralization Reactions 163Buffers 166Summary 169

x Contents

Chapter 14 Electrochemistry 171Introduction 171History 171Oxidation and Reduction 176Oxidation Numbers 177Electrochemical Cells 179Electrolysis 184Batteries 185Corrosion 189Electroplating 190Commercial Applications and Electrorefining 191Summary 193

Chapter 15 Organic Chemistry 195Introduction and History 195Organic Compounds 198Aliphatic Hydrocarbons 199Aromatic Hydrocarbons––Benzene 205Functional Groups and Organic Families 207Petroleum 216Organic Products 219Summary 220

Chapter 16 Biochemistry 221Introduction 221Carbohydrates 221Lipids 226Proteins 229Nucleic Acids 233Biotechnology 236

Chapter 17 Nuclear Chemistry 241Introduction 241Nuclear Stability and Radioactivity 241Radioactive Decay and Nuclear Reactions 243Half-Life and Radiometric Dating 244Nuclear Binding Energy––Fission and Fusion 246Fusion and Stellar Evolution 251Transmutation 252Nuclear Medicine 254Radiation Units and Detection 255Biological Effects of Radiation 257Summary 259

Contents xi

Chapter 18 Environmental Chemistry 261Introduction 261Stratospheric Ozone Depletion 262Acid Rain 266The Greenhouse Effect 270Water Quality 273Air Pollution 278Pesticides 281Summary 287

Chapter 19 The Chemical Industry 289Introduction 289Acids and Alkalis 290Explosives 292Synthetic Dyes 293Synthetic Fibers and Plastics 297Synthetic Rubber 299Petrochemicals 301Industry Histories 302Summary 307

Chapter 20 Chemistry Experiments 309Introduction 309The Scientific Method 309Chemistry Experimentation 312Chemical Activities 313

Chapter 21 A Future in Chemistry 327Introduction 327Chemist’s Job Description and Training 327Careers in Chemistry 330Professional Development and the American Chemical Society 332

Glossary 335

Brief Timeline of Chemistry 351

Nobel Laureates in Chemistry 355

Table of the Elements 361

Selected Bibliography 363

Author Index 365

Subject Index 369

Preface

Perhaps not everyone is a student of chemistry in the formal sense of being enrolled in a chem-istry course, but everyone uses chemistry on a daily basis. Every time a package label is readto check ingredients, food is tasted to determine whether it needs spices added, or mixtures areprepared, chemistry is practiced. Without even knowing it, all of us have accumulated a basicunderstanding of chemistry by observing the world around us. Observations such as salt melt-ing ice, milk souring, leaves turning color, iron rusting, and wood burning demonstrate thatchanges are constantly occurring. Because chemistry is essentially the study of change andwe are constantly observing these changes, we all are students of chemistry.

The Basics of Chemistry is written for students beginning a formal study of chemistry. Thesereaders are primarily high school and college students enrolled in their first chemistry course.In addition to these students, individuals who are not enrolled in a chemistry course but wouldlike a general overview of the subject should find this book helpful. Teachers of all grades mayuse The Basics of Chemistry as a general reference on the subject.

The Basics of Chemistry is a general reference book that presents the basic scientific con-cepts of chemistry in addition to providing information on several related subjects. Chaptersprogress through several areas. The first several chapters focus on chemistry’s roots as a mod-ern science. Although the first chapters focus heavily on the historical development of chem-istry, most of the other chapters give a historical overview of the chapter’s content. The heartof this book is devoted to explaining basic chemistry concepts. In these chapters, I review thecontent found in a beginning chemistry course. Subjects such as nomenclature, chemical bond-ing, acids and bases, equilibrium, kinetics, solutions, and gases are covered. Building on thesegeneral concepts, the chapters that follow explore several subdivisions of chemistry and pre-sent additional concepts especially important in these subdivisions. Chapters on organic chem-istry, biochemistry, electrochemistry, nuclear chemistry, and environmental chemistry areincluded. Chapter 19 presents an overview of the chemical industry and includes informationon industrial chemistry. Chapter 20 presents the experimental method and twenty chemistryactivities. The final chapter provides an overview of the chemistry profession, possible careersin chemistry, and chemical education.

This book is enhanced by other features. It includes a glossary of chemistry terms. Glossaryterms are typed in boldface when they first appear in the text. The year of birth and year of

xiv Preface

death of important individuals are given in parentheses following the first appearance of a per-son’s name. Ample illustrations, tables, and analogies are used to clarify basic chemistry con-cepts and highlight important chemical information. The chapter on the chemistry industryprovides background information on the founding of a number of well-known chemical com-panies.

The most difficult task in writing a book called The Basics of Chemistry is deciding whatmaterial to include in the book. What may be “the basics” for one person may not be foranother. Rather than present a list of basic facts, I present chemistry from a broad perspec-tive. An introductory book on such a broad subject should give the student a glimpse of thediscipline, provide basic concepts and information, and offer a foundation for further study.In this respect, an introductory text is like wiping the moisture from a fogged window. Manyobjects become clear through the transparent opening, but the rest of the glass remains foggy.Objects viewed through the fog are unclear. Even after wiping away all of the moisture, onlya small portion of the outside world can be seen. My hope is that The Basics of Chemistry pro-vides the student with a small, cleared opening into the world of chemistry. Once a glimpseis gained through this opening, the student can continue to enlarge the opening. After clear-ing the window, then it’s time to move to the next window.

Acknowledgments

Many individuals helped in the preparations of this book. Jennifer Rochelle provided timelytyping whenever needed. Sean Tape and the team at Impressions Book and Journal Servicesdid an exceptional job of editing the manuscript. Greenwood Press initiated this book as partof a series of general reference books in the sciences. I would like to thank the staff at Green-wood for the opportunity to write this book. Greenwood’s Debby Adams provided guidanceand a firm hand at the wheel as she shepherded the project through production. Rae Déjur’sillustrations demonstrate that one picture is truly worth a thousand words, and in some casesmuch more. A special thanks goes to Robert Krebs who mentored me throughout this projectby providing encouragement and valuable feedback.

A number of chemical structures were drawn by ACD/ChemSketch version 4.55 developedby Advanced Chemical Development, Inc. For more information please check the ACD Website at www.acdlabs.com.

1

Introduction: Chemistry and Its Divisions

IntroductionWhat do you think of when the word “chem-istry” is mentioned? For many, chemistryprojects images of a crazed scientist dressedin a white lab coat mixing a bubbly, spew-ing, stinking concoction. The absentmindedprofessor is the classic stereotype of achemist. The original Disney movie (TheAbsentminded Professor) released in 1961tells the story of a chemistry professor whoaccidentally creates “flubber,” a substancewith gravity-defying levitative properties. Afew years later in 1965, chemists producedwhat seemed to be real-world flubber. It wasand continues to be marketed as the super-ball®. A more modern version of a chemistis FBI agent Stanley Goodspeed portrayedby Nicholas Cage in the movie The Rock.Goodspeed is the FBI’s point man on chem-ical weapons. In the movie, he secretly pen-etrates Alcatraz Island, which is held by agroup of ultrapatriotic military extremists.Once on Alcatraz, he must beat a deadlineto disarm rockets loaded with a deadlychemical weapon aimed at San Francisco.

Hollywood’s portrayal of chemistrystretches the imagination, but probably no

more than that of a person living at thebeginning of the twentieth century whocould foresee life at the start of the twenty-first century. A quick look around is all ittakes to see how chemistry has transformedour modern world. The plastic disposableballpoint pen I use to write with has been inexistence only since 1950. Your clothes maybe made of natural materials, but chancesare you are wearing plenty of syntheticssuch as nylon, dacron, and polyurethane.You drink tap water that was chlorinated inthe last few hours and will be chlorinatedagain at the treatment plant before it contin-ues its journey through the water cycle. Ourvehicles burn refined oil, and catalytic con-verters reduce the amount of pollutionentering the atmosphere. Just opening amedicine cabinet gives ample evidence ofthe advances in chemistry during the lastone hundred years. Cosmetics, soaps, med-icines, and cleansers are all products of themodern chemical industry.

Much of the twentieth century has beencharacterized by the use and abuse of chem-istry. There is no doubt that advances inchemistry and chemical technology haveimproved and extended life, but there is also

2 Introduction: Chemistry and Its Divisions

a price for living in a modern chemical soci-ety. Environmental costs such as ozonedepletion, acid rain, water pollution, air pol-lution, and hazardous wastes are part of theprice we pay for the benefits of living in achemical society. New drugs and medicineshave eradicated disease, and advances inchemistry have allowed possible fertilizersand other agricultural advances to increasethe world’s food supply. These develop-ments have led to a booming world popu-lation currently at six billion and growing.To sustain a growing world population ofover six billion means converting raw mate-rials into products. In recent years, we haverecognized the need to “reduce, reuse, andrecycle.” We know that one of the funda-mental principles of chemistry is that mat-ter is not created or destroyed, and chemicalprocesses themselves play a major role inrecovering raw materials from used prod-ucts. Many items once considered waste arenow recycled supplementing raw materialsobtained from the Earth. As more countriessuch as China develop their industrial infra-structure, the demand for energy and mate-rial resources is bound to increase.

The laws of thermodynamics tell us thatthe conversion of raw materials into finishedgoods requires energy and ultimately resultsin an ever-increasing amount of heatreleased into the environment. This wasteheat is accompanied by carbon dioxide pro-duced from the burning of fossil fuels.There seems to be a consensus among sci-entists that our planet’s temperature isincreasing due to the greenhouse effect,although a minority of scientists do notbelieve that our planet’s temperature isincreasing. Many scientists also feel thatincreased temperatures are part of naturalfluctuations and that humans have littleimpact on the recent trend of elevatedglobal temperatures. International effortsare now underway to curb carbon dioxide

emissions worldwide, such as revegetation,curbing slash and burn practices, and se-questering carbon dioxide in the sea, arestalled by disagreements between developedand underdeveloped countries, as well asEurope and the United States.

The Scientific EnterpriseSo how do we strike a balance between

the benefits of better living through chem-istry and the costs associated with thislifestyle? One answer to this question isknowledge. Greater knowledge equips uswith the tools to make informed decisionswhen we live in a free democratic society.Chemical knowledge is based on a long his-tory of chemical research. The methodolog-ical procedure to attain scientific knowledgeis often defined as the scientific method. Theclassical view of the scientific method pre-sents science as a neat step-by-step proce-dure starting with formulating a question,posing a hypothesis, performing a con-trolled experiment to address the hypothe-sis, and drawing conclusions from theexperiment. Although this view is accuratein an idealized sense, in reality, true scienceis filled with pitfalls, blind alleys, dead ends,false starts, fortuitous accidents, and humanidiosyncrasies. Books on science often givethe impression that science proceeds in alinear, uniform fashion. More often, scien-tific progress follows a circuitous route withperiods of accelerated discoveries and long,relatively dormant periods.

The theories and facts we cling to todayare only an experiment away from revision.As you read these words, chemical researchis occurring in thousands of labs across theglobe. The modern scientific enterprise isvastly different from that of the quintessen-tial scientist working in isolation in a smalllaboratory. Modern research is characterizedby research teams led by scientists who spe-

Introduction: Chemistry and Its Divisions 3

cialize in very specific fields. Teamwork isessential as research teams across the worldattempt to answer fundamental questions.These teams work both cooperatively andcompetitively with other teams working onsimilar problems. Work is continuallyreviewed, critiqued, revised, and publishedin scientific journals. The advent of theInternet has led to both a wider audienceand a faster distribution for research results.Much of today’s research takes place in uni-versities and is supported by governmentalagencies that distribute millions of dollarsannually to those teams whose work isdeemed worthy of support. In addition tothis research, numerous private companiesoperate their own private research labs todevelop products and processes to give theman advantage over their competitors. Chem-ical research is both basic and applied. Basicresearch addresses the fundamental laws ofnature and is undertaken to advance the the-oretical foundation of chemistry. Appliedresearch is performed with a specific appli-cation in mind, for example, better productsor to improve the quality of life. Researchshould not be viewed in terms of a basic orapplied research dichotomy. Most researchcontains elements of both basic and appliedresearch.

As new knowledge is gained throughresearch, additional questions are raised andnew areas of research are continuallyexposed in a neverending process. Becauseknowledge is power, as we embark on thetwenty-first century, a knowledge of chem-istry is essential to shape a sustainablefuture. Chemical knowledge may lead to asimple decision to recycle or choose oneproduct over another. Perhaps this knowl-edge helps you make a decision to under-take a certain medical treatment or vote foror against an issue. Chemical knowledgecan help guide you in any number of choicesyou make today. More important, this

knowledge will help you with decisionsconcerning novel areas we can only imaginebut know will develop as the twenty-firstcentury unfolds.

What Is Chemistry?Chemistry is the branch of science that

deals with the composition and structure ofmatter and the changes that matter under-goes. Matter is anything that has mass andoccupies spaces, which means just aboutanything you consider. This book, your bodyand the air you breathe are all examples ofmatter. Matter is simply the stuff that makesup our universe.

Chemistry, like all branches of science,is a method that attempts to simplify andorganize. Every object we might consider isa separate piece of matter, and matter can beclassified in any number of ways. One sim-ple classification scheme for matter is basedon the three states of matter: solid, liquid,and gas (four if we include plasma).Another classification scheme, and one thatis fundamental to chemistry, is classifyingmatter by chemical composition. Ratherthan speaking of water as a form of matter,we can speak of water composed of approx-imately 11% hydrogen and 89% oxygen bymass. Similarly, air is a mixture of mattercontaining approximately 78% nitrogen,21% oxygen, a little less than 1% argon byvolume, and an assortment of other gases insmall percentages.

Much of the history of chemistry hasconcerned itself with determining the ulti-mate composition of matter. The Greeksconsidered all matter to be composed of dif-ferent combinations of earth, air, fire, andwater. It is only during the last two hundredyears that the modern idea of chemical ele-ments developed and only in the last onehundred years that we have determined thatelements themselves are composed of


protons, neutrons, and electrons. Hence, wecan think of all matter to be composed of 91 naturally occurring elements, and protons,neutrons, and electrons are the buildingblocks that make up the elements. Chem-istry not only deals with the composition ofmatter, but how the pieces of matter fittogether, that is, the structure of matter. Thestructure of matter can have a major impacton its properties. For example, carbon canexist as a two-dimensional flat structure, ora soft, black graphite, or in the pyramidalthree-dimensional tetrahedron shape weknow as diamond (Figure 1.1).

So far, we have noted that chemistryinvolves the composition and structure ofmatter. But beyond knowing what the com-position of matter is and how the pieces fittogether, chemistry is about how matterchanges from one substance to another.When we refer to chemical change, we meanthat the composition of matter has changed.A chemical change occurs when a substanceor substances change into other substances.

For example, when hydrogen and oxygen arebrought together under the right conditions,they can combine and change into water.When the carbon in paper reacts with oxygenin the air during combustion, it changes intocarbon dioxide. The key when consideringwhether a chemical change has taken place isto ask the question: do I end up with some-thing different from what I started? Chemi-cal changes result in the production ofsomething entirely different from the origi-nal substances. The element sodium is ahighly reactive metal that reacts violentlywith water. Chlorine is a highly toxic gas. Yet,put sodium and chlorine together under theright conditions and you have table salt.

Not all changes are necessarily chemi-cal changes. Changes may also be physicalchanges. When water changes from ice toliquid to steam, change has certainly takenplace, but not a chemical change. Water inits solid ice state is H2O. After melting andvaporizing, it is still H2O. The substancewater has not changed. The physical statehas changed, but not the substance. Simi-larly, when sugar dissolves in water, nochemical change takes place. We start withsugar and water and end with sugar andwater. Again, only a physical change occurs,because we still have sugar and water afterthe sugar dissolves. To summarize, chem-istry can be thought of as the science of thestudy of matter: what matter is made of, howit is made, and the changes in matter fromone substance into another.

Divisions of ChemistryThe definition given in the last section is

very general. Although all chemists areinvolved in the study of matter, the field is sobroad that it helps to divide chemistry intodivisions. These divisions characterize dif-ferent aspects of the study of chemistryusing some common feature. Chemistry may

Figure 1.1Diamond and Graphite Structure (Rae Déjur)

Introduction: Chemistry and Its Divisions 5

be divided into the two very large divisionsof inorganic and organic chemistry.

Inorganic chemistry involves thestudy of the chemical elements, their com-pounds, and reactions excluding those thatcontain carbon as the principal element.Organic chemistry is the study of the composition and the reaction of carbon-containing compounds. The distinctionbetween inorganic and organic results fromthe days when chemical compounds wereclassified according to their origin. Inor-ganic means not derived from life. Inor-ganic substances were thought to originateonly from mineral, nonliving sources. Con-versely, organic compounds were thought tohave come from living plant or animalsources. All organic substances contain car-bon, and it was once believed that these car-bon substances could originate only from aliving source. We know today that manyorganic compounds are synthesized frominorganic, nonliving sources. Some exam-ples include polypropylene, acrylics, andnylon used in clothing; polystyrene (styro-foam) used for insulation; and ethylene gly-col (antifreeze). Even though the originaldistinction separating inorganic and organicdoes not hold, the terms are still used to distinguish two broad areas of chemistry.For our purposes, we can think of organicchemistry as simply the chemistry of car-bon compounds and inorganic chemistry asthe study of all other compounds. It shouldbe noted, though, that certain carbon com-pounds, most notably carbon dioxide andcarbonates, are classified as inorganic.

In addition to the large divisions oforganic and inorganic chemistry, severalother large divisions of chemistry exist. Bio-chemistry is the study of chemical substancesassociated with living organisms. Bio-chemists study the nature of biological sub-stances. The study of DNA, viruses, andimmune systems are examples of biochemi-

cal research. Physical chemistry deals withthe structure of matter and how energy affectsmatter. Physical chemists are concerned withthe physical characteristics of matter. Areasof research in physical chemistry mightinclude how chemicals absorb light or howmuch energy is released or absorbed when achemical reaction occurs. Closely related tophysical chemistry is the study of nuclearchemistry. Nuclear chemistry focuses on thestudy of atomic nuclei, nuclear fission reac-tions, and nuclear fusion reactions. Thenucleus is the positive central portion of anatom composed of positively charged protonsand uncharged neutrons. Fission reactionsinvolve the physical splitting of heavier ele-ments into smaller elements, and fusioninvolves lighter elements physically combin-ing to form heavier elements, for example,the fusion of hydrogen to form helium. Ana-lytical chemistry deals with techniques usedto identify and quantify the composition ofmatter. Analytical chemists, as the nameimplies, analyze substances for their content.An analytical chemist might perform taskssuch as determining the sugar content of afruit juice, the amount of pollutant in a waterbody, or the purity of a drug. Environmentalchemistry focuses on the occurence of natu-ral and synthetic substances in the environ-ment and their impact on the environment.The study of the chemistry of pollution is amajor area of concern for environmentalchemists. Problems involving water pollu-tion, air pollution, solid waste, hazardousmaterials, and toxic substances involve astudy of environmental chemistry. Each ofthe major areas of chemistry mentionedabove, however, are not rigid divisions.

The nature of chemistry is such that thefields of study in many areas naturally crossover into other areas. A biochemist naturallyworks in the area of organic chemistry, andan environmental chemist who studies radi-ation is concerned with nuclear chemistry.


Many branches of chemistry exist in con-junction with other areas of science. Forexample, the study of the chemistry of rocksand minerals falls under the category ofgeochemistry. The combination of medicineand chemistry leads to medicinal chemistry.Other divisions include agricultural chem-istry, food chemistry, petroleum chemistry,soil chemistry, and polymer chemistry.

As chemical knowledge continues toexpand, we are certain to develop new areasof study in chemistry. Just recently, chemistshave performed simple computer functionsusing chemical reactions rather than per-forming these functions electronically.Chemical reactions that perform computeroperations bring computers closer to model-ing human thought processes. Our thoughtsresult from chemical reactions that occur inour brains. As chemists work at the forefrontof this area of technology and others, new

interdisciplinary areas will certainly develop.Those new areas add to existing branches ofchemistry and at the same time create newareas of study.

In the chapters to follow, the basic con-cepts that guide chemists in their study ofmatter are examined. In addition to intro-ducing these basic concepts, the historicalfoundation of chemical thought, as well asnew areas of chemical research, are ex-plored. This knowledge is intended to givethe reader a clearer view of the wonderfulworld of chemistry that is all around us; andwith this, a foundation to build upon. Allprogress by scientists is dependent on thework of those who precede them. In a letterto Robert Hooke, Isaac Newton stated, “If Ihave seen farther than others, it is becauseI was standing on the shoulders of giants.”The goal of this book is to help you see alittle farther.

IntroductionChemistry has always existed. The forma-tion of the Earth and the development oflife involved innumerable chemical pro-cesses. Humans, as part of the naturalworld, are part of these natural processes.As humans evolved and observed chemicalprocesses constantly occurring all aroundthem, they attempted to control these pro-cesses in their daily existence. For example,combustion in the form of fire could beused to clear land, light the darkness, cookfood, and provide protection. Eventually,humans learned that fire could be use toharden clays and melt ores to refine metals.Other chemical discoveries were used toimprove living conditions. In this respect,our ancient ancestors were no different thanmodern humans. Chemical technology hasbeen used throughout history to improveour lives. As a modern science, chemistryhas only existed for two hundred years.Modern chemistry developed out of a tra-dition of chemical technology appliedthroughout history. This chapter examinessome of that history, and the followingchapter focuses on the events that estab-lished chemistry as a modern science.

Early HistoryChemistry is often called the central sci-

ence. It derives this name because of itsimportance to all the other sciences.Although chemistry did not exist as a mod-ern science until two hundred years ago,humans have used chemistry from prehis-toric times. Evidence of the early uses ofchemistry includes cave paintings datingfrom 25,000 B.C. It is not hard to imagineour ancestors obtaining natural pigments bysqueezing berries or mixing crushed rockswith water to produce different colors to useas paints. Similar processes could have beenused to obtain dyes for clothes or for bodydecorations. Fragrances obtained fromflower extracts or fats rendered from ani-mals were the prehistoric version of the cos-metic and repellant industry. Evidence forthe use of fermentation to produce wineand beer dates from the earliest civilizations.Fermentation involves the conversion ofglucose to ethyl alcohol (ethanol) in thepresence of yeast:

2

A Brief History of Chemistry

r

8 A Brief History of Chemistry

One of the earliest uses of fire was in pot-tery making. Early humans undoubtedlyobserved that when clay was heated itswater was driven out and a hard rock sub-stance remained. Our ancestors made clayimplements and art fixtures by heating theirwork in open pit fires, and clay implementsfrom 20,000 years ago have been found.

Early human civilizations used stone,bone, and wood for objects. Approximatelyten thousand years ago, metals first ap-peared. The first metals used were thosefound in their native form, or in a pure,uncombined state. Most metals today areacquired from an ore containing the metalin combination with other elements such asoxygen. The existence of native metals israre, and only a few metals exist in nativeform. Iron and nickel were available in lim-ited supply from meteorites. The first met-als utilized widely by humans were copper,silver, and gold. Pure nuggets of these met-als were pounded, in a process known as“cold hammering,” with stones into variousshapes used for weapons, jewelry, art, andvarious domestic implements. Eventually,smiths discovered if a metal was heated itcould be shaped more easily. The heatingprocess is known as annealing. Because thesupply of native metals was limited, metalitems symbolized wealth and status forthose who possessed them.

The discovery that metals could beobtained from ore-bearing minerals signaledthe true start of the metal age. The first metalto be extracted from its ore was copper. Thetechnology used for firing clay could bedirectly applied to the smelting of metals.Kilns and furnaces were adapted withtroughs to capture metals obtained duringthe smelting process. Smelting involvedheating the metal-bearing ore with wood orcharcoal. The carbon from the charcoalcombined with the oxygen in the ore sepa-rating the metal from the oxygen. During

the smelting process, the sulfur, which wasoften associated with metal ores, was drivenoff as sulfur dioxide. Sulfur, with its char-acteristic odor and yellow fumes that arereleased when it burns, came to be closelyassociated with the transformation of met-als and subsequently played a key role inalchemical reactions. The earliest evidenceof the smelting of copper dates from around6000 B.C.

By 4000 B.C., smelters were combiningcopper and tin to produce the alloy bronze.An alloy is simply a mixture of metals. Thebronze produced in ancient times consistedmostly of copper mixed with between 1%and 10% tin. Because the production ofbronze required a source for both copper andtin, major trade centers and trade routesdeveloped to mine and produce bronze. Oneof the major sources of tin from the BronzeAge until recent times was the Cornwallregion in southwest England. The Romansand other groups settled the area andexported tin to other parts of Europe. Bronzewas the first metal used over wide geographicregions. Advances in kilns, annealing, alloy-ing, oxidation, and smelting continued toadvance the craft of metalsmithing as the pro-duction of bronze flourished.

Iron, which requires a higher tempera-ture to smelt than does bronze, is found inrelics dating from approximately 1500 B.C.Early evidence of the use of iron was pro-vided upon the opening of King Tut’s tombwhen an iron dagger that dated from 1350B.C. was found buried with the king. Thefirst irons produced were of inferior quality.Repeated heating and hammering (temper-ing) were needed to produce an iron ofacceptable quality. A major advancementmade in iron production was the use of bel-lows to obtain temperatures high enough(approximately 1,500°C) to smelt iron,although even irons created with the highertemperatures and the tempering process

A Brief History of Chemistry 9

were inferior to bronze. It was only wheniron was heated with carbon that an ironwith superior strength qualities wasobtained. This process foreshadowed themodern steel industry. Steel is iron to whichcarbon has been added.

The metal age brought about a numberof advances in chemical technology. Yet, thesmelting of metals and their incorporationinto products involved more art and craftthan science. Smiths were thought to pos-sess magical powers to be able to obtainrefined metals from rock. Civilizations andtheir rulers closely guarded their tradesecrets. After all, metals were the strategicweapons of ancient civilizations. Still,progress made in metallurgy had a directimpact on the development of chemistry.Advances in heating devices such as fur-naces and kilns, fuels to fire the kilns, andprocesses such as mixing, extracting, andmining were directly applicable to latteradvances in chemistry.

Greek ScienceThe technologies of early civilizations

used chemical technology empirically. Thatis, different substances and methods werediscovered through experimentation. Ap-plied science took precedence over any the-oretical understanding of the general lawsand principles governing chemical pro-cesses. As early as 600 B.C., Greek philoso-phers sought a more basic understanding ofmatter. Anaximander of Miletus (circa610–545 B.C.) believed air was the primarysubstance of matter and that all other sub-stances came from air. For example, fire wasa form of thin air, and earth was thick air.Heraclitus of Ephesus (544–484 B.C.) be-lieved fire was the primary substance andviewed reality in terms of change. Matterwas always in a state of flux, and nature wasalways in a process of dissolving and

reforming. His ideas on change are summa-rized in his saying: “you cannot step twiceinto the same river.” Heraclitus’ philosophyalso considered reality in terms of opposites.The way up and the way down were one andthe same depending on one’s perspective.

The Eleatic school of Greek philosophywas centered in the southern Italian city-state of Elea, which flourished in the sixthand fifth centuries B.C. The Eleatics directlyopposed the ideas of Heraclitus and believedthe universe did not change. They feltknowledge acquired through sensory per-ception was faulty because reality was dis-torted by our senses. One of the leadingEleatics was Parmenides (circa 515–450B.C.), who viewed the world as a unity andbelieved change was an illusion. To maketheir ideas about the absence of change con-form to observations, the Eleatics expandedupon Anaximander’s idea of one primarysubstance. Rather than considering air as theprimary substance, the Eleatic philosopherEmpedocles of Agrigentum (circa 495–435B.C.) proposed four primary elements: air,earth, fire, and water.

In response to the Eleatics, the atomistschool led by Leucippus (circa 450–370B.C.) and his student Democritus (circa460–370 B.C.) developed. The atomistsrefuted the idea that change was an illusionand developed a philosophy that supportedthe sensory observations of the physicalworld. The atomist school proposed that theworld consisted of an infinite void (vacuum)filled with atoms. According to Democritus,“nothing exists except atoms and emptyspace, everything else is opinion.” Atomswere eternal, indivisible, invisible, minutesolids that comprised the physical world.Atoms were homogeneous but existed indifferent sizes and shapes. The idea of atomsdeveloped on logical grounds. If a piece ofgold was continually divided into smallerand smaller units, the gold would eventually


reach a point where it was indivisible. Theultimate unit of gold, or any other substance,was an atom. According to Democritus, an infinite number of atoms moved throughthe void and interacted with each other. The observable physical world was made of aggregates of atoms that had collided and coalesced. The properties of matterdepended on the size and shape of theatoms. Change was explained in terms ofthe motion of the atoms through the void orhow atoms arranged themselves. Aristotle(384–322 B.C.) rejected the idea of atoms,but Democritus’ atomist philosophy waslater adopted by Epicurus (341–270 B.C.)and the Roman poet Lucretius (circa 95–55B.C.). Almost two thousand years would passuntil atoms were resurrected by Dalton inthe development of modern chemistry.

Of all ancient philosophers, Aristotlehad the greatest influence on scientificthought. Aristotle’s writings and teachingestablished a unified set of principles thatcould be used to explain natural phenome-non. Although today we would consider hisscience rather foolish, it was based onsound, logical arguments that incorporatedmuch of the philosophical thought of histime. The teleological nature of Aristotle’sexplanations led to the subsequent accep-tance by the Church in western Europe. Thiscontributed to Aristotle’s two thousand yearlongevity in Western thought.

Aristotle’s Meterologica synthesized hisideas on matter and chemistry. Aristotlebelieved that four qualities could be used toexplain natural processes. The four qualitieswere hot, cold, dry, and moist. Hot and coldwere active qualities. Hence, the addition orsubtraction of heat leads to the transforma-tion of things. Moist and dry were the resultof the action of the active qualities. Onlyfour possible combinations of these quali-ties could exist in substances. These werehot and moist, hot and dry, cold and moist,

and cold and dry. Opposite qualities couldnot coexist. The combinations of hot andcold and moist and dry were impossible inAristotle’s system. The four allowable com-binations determined the four basic ele-ments. Air possessed the qualities hot andmoist, fire possessed hot and dry, earth pos-sessed dry and cold, and water possessedwet and cold. Aristotle’s system is summa-rized in Figure 2.1, which is found in anumber of ancient chemical writings.

Using the four qualities of matter andfour elements as a starting point, Aristotledeveloped logical explanations to explainnumerous natural observations. Both theproperties of matter and the changes in mat-ter could be explained using Aristotle’s theory.

Aristotle explained the process of boil-ing as a combination of moisture and heat. Ifheat is added to a substance that containsmoisture, the heat draws the moisture out.The process results in the substance divid-ing into two parts. When the moistureleaves, the substance becomes thicker,whereas the separated moisture is lighterand rises. Aristotle used this type of rea-soning to explain numerous physical andchemical processes including evaporation,

Figure 2.1Aristotle’s Diagram on Matter and Change


decaying, condensation, melting, drying,putrefaction, and solvation. Although Aris-totle’s philosophy explained much, it didhave its shortcomings. A simple examplehelps to illustrate how problems arise withany scientific theory. According to Aristotle,hard substances were hard because they hadan abundance of earth and possessed thequalities of dry and cold. Aristotle claimedearth was heavy and so it moved down. Softsubstances would not contain as much earth,but contained more water and air. Thesesubstances would not be attracted asstrongly downward. Now consider ice andwater. Ice is dry and cold compared to water,which is moist and cold. We would expectice to have more of the element earth and astronger downward attraction than water.Yet, hard solid ice floats in water. Eventhough many other problems existed withAristotle’s theory, it provided reasonableexplanations for many things.

In closing this section, remember thatAristotle rejected the concept of atoms.Aristotle could not accept the idea of a voidspace and believed that “nature abhors avacuum.” Furthermore, Aristotle did notconsider internal structure. Substances con-tained their qualities and elements as ahomogenous mixture. An Aristotelianwould explain the reaction of hydrogen gasand oxygen gas to produce liquid water as

Air � Air Water

This reaction shows how one basic elementcould convert directly to another. An entirelynew element is produced from an originalelement using Aristotle’s system. A modernchemist would write the reaction:

� O 2H O

In this reaction, water is formed from thegases hydrogen and oxygen. We end up with

water, but we still have hydrogen and oxy-gen. Rather than producing new elements,the elements hydrogen and oxygen combineand rearrange themselves to form water.

AlchemyThe teachings of Aristotle and events in

Egypt, India, China, and Mesopotamia stim-ulated the practice of alchemy. The earlyroots of alchemy, which was practiced fornearly two thousand years, are difficult totrace because much of the practice wasshrouded in mystery and transmitted by oraltradition. The two main goals of the al-chemists were to produce gold from basemetals and to develop potions that wouldconfer health and even immortality. Thealchemists were crafts people who com-bined serious experimentation with astrol-ogy, incantations, and magic in hopes offinding the philosopher’s stone.

The philosopher’s stone or materialprima was a substance (a powder, tincture,or stone) that was more pure than gold itself.It was thought that by using a minute quan-tity of the philosopher’s stone a base metalcould be elevated to gold through analchemical process. Besides perfecting com-mon substances into gold, the philosopher’sstone was thought to convey immortality,cure all common ills, and cleanse the spiri-tual soul. The philosopher’s stone wassought by all and possessed by few. Itspower conveyed perfection and the ability toreach a pure state to whoever held the stone.The origination of the idea of a philoso-pher’s stone is unclear, but reference is madeto the idea in several ancient cultures. Anelixir of immortality is present in Indianwritings as early as 1000 B.C. Alchemicalpractice in China several centuries B.C.through A.D. 500 referred to a “pill ofimmortality” and the use of gold to conferimmortality. The Chinese desire for gold, as


opposed to that in the West, was strictly con-cerned with physical and spiritual healthrather than acquiring wealth. Hermetic phi-losophy and the writings of Hermes (apseudonym) incorporated both Egyptianand Greek mythology in the first few cen-turies A.D. and mentioned a solidified spiritderived from the Sun capable of perfectingall things. Ancient ideas on the philosopher’sstone were undoubtedly passed among cul-tures. Ideas from India and China couldhave easily found their way into medievalEurope through Arab cultures.

Alchemy’s foundation was based on theteachings of Aristotle. Aristotle held thatrocks and minerals were alive and grew inthe interior of the Earth. Like humans, min-erals attempt to reach a state of perfectionthrough the growth process. Perfection forminerals was reached when they ripenedinto gold. Based on these premises, thealchemist sought to accelerate the ripeningprocess for metals by subjecting them to aseries of physical and chemical processes.One typical series might include heating themetal with sulfide to remove impurities.Then a starter seed of gold was added to themetal. After the seed was added, the metalwas treated with arsenic sulfide. This treat-ment resulted in whitening of the metal,which could be interpreted as a productionof silver, or a stage halfway to the perfectgold stage. Finally, the whitened metal couldbe treated with polysulfides to produce thecharacteristic yellow gold color. Of course,the process could take weeks or months, andit all had to be performed under the rightsign of the stars.

The early center of alchemy was theintellectual capital of ancient Greece,Alexandria. Very little remains of the origi-nal alchemical manuscripts from ancientGreece. The rise of Christianity and con-cerns about disrupting the economy eventu-ally led the Roman Emperor Diocletian to

outlaw alchemy in A.D. 296. Diocletian’sedict resulted in the destruction of much of the knowledge accumulated by the al-chemists. Two original works that did sur-vive are known as the Leyden Papyrus X andthe Stockholm Papyrus. These two relatedworks were discovered in a tomb of a buriedEgyptian in ancient Thebes. The papyri datefrom the end of the third century, although itis believed the papyri are copies of an orig-inal written sometime between A.D. 100 and300. They are written in Greek and consistof a series of recipes for alchemical prepa-rations. The recipes include instructions formaking alloys, determining the purity ofmetals, making fake gold and silver, pro-ducing imitation gemstones, and dyeing textiles.

The rise of the Arab Empire shiftedadvances in learning to the East. The ArabEmpire at its peak spanned from its easternborders with China and India all the wayacross northern Africa into southern Spain.The Arabs translated and dispersed theworks of the Greeks, and for the next onethousand years learning and advances in sci-ence, medicine, astronomy, and mathemat-ics were made by the Arabs. The Arabinfluence can be seen in the word “alchemy”itself. The word “al” is the Arab prefix for“the.” Adding this prefix to “chyma,” Greekfor melting of metals, gives “al chyma.” Theterm “alchemy” first appeared in the fourthcentury A.D. Another term to appear was “aliksirs.” Iksirs were coloring agents used asan ingredient in the procedure to producegold. Today, the term “elixir” is a reminderof its alchemical root.

The period between A.D. 700 and 1100was the peak of the Arab Empire. The vastgeographic expanse of the Arab influencemeant that a great variety of materials wereavailable and that there was ample opportu-nity for the cross-fertilization of knowledge.Several individuals from this period had a


lasting influence on the development of thechemical arts. Jabir ibn Hayyan (Latin nameGeber, 721–815) accepted many of the ideasof Aristotle but also modified Aristotle’sideas. Jabir proposed two primary exhala-tions. One exhalation was smoky andbelieved to contain small particles of earth.The other was vaporous and contained smallparticles of water. The principal smokyexhalation was sulfur, and the primaryvaporous exhalation was mercury. Jabirbelieved pure sulfur and mercury combinedin different proportions to produce all met-als. Most of the writings attributed to Jabirare actually a compilation from numerousArab alchemists and early Europeanalchemists written over several hundredyears. These writings described processesfor producing acids, metals, dyes, inks, andglass.

Al-Razi (Rhazes, 854–925) was a Per-sian who studied in Baghdad. Al-Razi wroteextensively on medicine, philosophy, astron-omy, and alchemy, but he was primarily aphysician. Al-Razi was less mystical thanhis contemporary alchemists and classifiedchemicals by their origin. According to Al-Razi, chemicals came from either animals,plants, and minerals or were derived fromother chemicals. Al-Razi wrote The Com-prehensive Book, which was an enormousmedical encyclopedia that synthesized med-ical practices of ancient Greeks, Syrians,Arabs, and Persians. Al-Razi was the firstperson known to describe the disease small-pox. Most of his alchemical writings havebeen lost, but Al-Razi believed in the atomicnature of matter. Al-Razi took a systematicapproach to science and rejected the idea ofdivine intervention. His rational methodsand descriptions were more consistent withmodern science than most individuals of histime. Ali al Husayn ibn Sina (Avicenna,980–1037) was another Persian physicianwhose voluminous works, including The

Canon of Medicine, guided the practice ofmedicine for 500 years after his death. Avi-cenna rejected the idea that a base metalcould be transformed into gold. Avicennaclaimed correctly that diseases were spreadthrough air and water. Much of Avicenna’steachings questioned the status quo andteachings of Aristotle.

By Avicenna’s time, around 1000, theArab Empire was in decline from both inter-nal and external forces. Factions of theIslamic faith battled one another. A generalintolerance of science pervaded Arab cul-ture, and scientists were not free to publishtheir ideas. Christian Crusaders from theWest and Mongol invaders from the Eastexerted pressure on the Arabic world. As tra-ditional Arab regions were recaptured byEuropeans, the classical knowledge that hadbeen preserved and advanced by the Arabsinfluenced European thinking. Major Arablearning centers, such as Toledo in Spain,provided works to rekindle European sci-ence. From the twelfth century, majoradvances in the chemical arts shifted fromArab lands to western Europe.

Chemistry in the MiddleAges

Scholars, especially clergy from theCatholic Church, were primarily responsi-ble for translating Arabic, Syrian, and Per-sian writings into Latin. The translators notonly reintroduced traditional Aristotelianphilosophy, but they also explained numer-ous technological advances made during theperiod in which the Arabs flourished. Trans-lated works brought to light advances in thepractice of distillation, glass making, filtra-tion, calcination, sublimation, and crystal-lization. From these translations it wasapparent that the Arabs and Persians hadadvanced chemical knowledge significantly,especially during the period from A.D. 700


to 1100. Contributions from the Arab andPersian regions followed the work of Jabir.Jabir developed chemistry by focusing onpractical applications using the method ofexperimentation.

Arab and Persian work in the area ofchemistry was principally conducted byindividuals practicing medicine. In additionto preparing medicines, experimenterssought improvements in various crafts suchas refining metals, preparing inks, dyeingfabrics, waterproofing materials, and tan-ning leather. Modern distillation originatedwith the Arabs and was improved by Avi-cenna, who invented the refrigerated coil.Advanced distillation techniques spreadthroughout the Middle East, and the Arabsused the refrigerated coil to distill ethanolfrom fermented sugar. The ethanol was usedas a solvent for extracting plant oils, whichwere substituted for animal fats in prepara-tions. Because essential oils were consid-ered the basis of odors and flavors, anyimprovement in extraction processes werevalued. The Arabs also incorporated otherindustries into their empire as theyexpanded. The manufacture of glass wasdominated by Syria and Alexandria untilA.D. 850. The Arabs conquered these areasand incorporated and improved the art ofglass making. Glass was used as a buildingmaterial and aesthetically in ornamentationand religious designs. After the fall of theArab Empire, the center of glass makingmoved west to Venice. Besides technicaladvances in processes and apparatus, theArabs had developed and improved thepurity of substances such as alcohols, acids,and gunpowder, which were not available tothe Europeans.

Several Europeans, while not al-chemists themselves, contributed to settingthe stage for modern chemistry. AlbertusMagnus (1200–1280), also known as Albertthe Great, played an important role in intro-

ducing Greek and Arabic science and phi-losophy in the Middle Ages. Albert pro-duced numerous commentaries on theworks of Aristotle. These commentarieshelped establish the value of studying natu-ral philosophy along with the philosophicallogic of Aristotle. Albert had a major influ-ence on his most famous student, ThomasAquinas (1225–1274). Aquinas reconciledAristotelian logic and Christian teaching.Albert was declared a saint in 1931 and isconsidered the patron saint of all who studyscience.

Roger Bacon (1214–1294) was anEnglish contemporary of Albert Magnus.Bacon was a Franciscan clergy who taughtat Oxford and conducted studies in alchemy,optics, and astronomy. He conducted majorstudies on gunpowder. In his Opus Majus(Major Work), Opus Minus (Minor Work),and Opus Tertium, Bacon argued that thestudy of the natural world should be basedon observation, measurement, and experi-mentation. Bacon proposed that the univer-sity curriculum should be revised to includemathematics, language, alchemy, and exper-imental science. Because of his teachings,Bacon often had difficulties with his supe-riors, although today, we accept many ofBacon’s ideas as the foundation of modernscience.

Another influential figure was the Swissself-proclaimed physician Paracelsus(1493–1541). Paracelsus’ real name wasPhillippus Aureolus Theophrastus Bombas-tus von Hohenheim. The name “Paracelsus”was a nickname he chose for himself, mean-ing “superior to Celsus.” Celsus was a second-century Greek poet and physician.Paracelsus was an itinerant scholar. He trav-eled throughout Europe undertaking bothformal studies and absorbing folk remedies.At various times in his life, he worked as asurgeon (which at that time was a low-classmedical craft), a physician to miners, and a


medical admistrator. Paracelsus applied theprinciples of alchemy to medicine. As such,he rejected the traditional medical practicesof his time, which were based on the ancientprinciples of Galen (circa 129–199).Paracelsus adopted the ideas of Aristotle andapplied these to medicine. He believed thatdisease was the result of an outside agentinvading the body. According to Paracelsus,disease was the result of seeds being spreadbetween individuals, and it was the physi-cian’s job to find a specific cure to negateeach disease seed. Paracelsus experimentedand searched for remedies by distillingmaterials to produce purified medicines. Inaddition to preparing an assortment of reme-dies, Paracelsus knew it was important todetermine the dosage necessary to producemaximum benefit. Paracelsus used mercurysalts to cure syphilis and as disinfectants inworking against infections. Although muchof his approach involved a systematic exper-imental search for cures, Paracelsus was a mystic who also used astrology in his practice.

Paracelsus attacked the establishedmedical community of his time. His criti-cism was harsh, and he alienated many ofhis established fellow physicians with hisbombastic writings. Yet, many of his prac-tices were sound, and more importantly,those physicians who did ascribe to his prac-tices achieved better results than when usingtraditional methods. Paracelsus’ works werenot published and distributed widely untilsome twenty years after his death. By theend of the sixteenth century, Paracelsus’methods had attracted a number of follow-ers, and his methods formed the new basisof standard medical practice. Paracelsus’followers became known as the iatrochem-ical physicians. They used a variety of drugsand treatments that depended on specificdosages of medicine prepared with specificpurity. His ideas on chemical purity and for-

mulation also found use in other crafts suchas metallurgy and soap making. Chemistsbecame known as dealers in medicinaldrugs who gained knowledge about prepa-rations from skilled masters (Figure 2.2).Although Paracelsus was an alchemist, hiswork led to the term “alchemy” beingrestricted to the practice of trying to convertbase metal into gold. Through the work ofParacelsus and his followers, the term“chymia” was now reserved for the study ofmatter.

Several other individuals made impor-tant contributions at about the same time asParacelsus, further establishing chymia as ascience apart from alchemy. Vannoccio

Figure 2.2Woodcut print of chemical lecture from1653. Small furnace sits on table and variousglassware used for distillation rests onshelves above stove in background. Imagefrom Edgar Fahs Smith Collection, Univer-sity of Pennsylvania Library.


Biringuccio (1480–1539) wrote Pirotech-nia in 1540. This work was a basic textwritten in Italian on the science of metal-lurgy. Biringuccio’s work provided tech-niques for processing metal, and he rejectedthe alchemists’ idea of transmutation ofmetals. A similar work focusing on the pro-cessing of metals was Concerning Metalswritten by Georgius Agricola (1494–1555).Andreas Libavius (1540–1616) critiquedParacelsus. His work Alchemia published in 1597 was a comprehensive review ofchemical knowledge available at the time.Libavius’ work organized chemistry bysummarizing the methods, operations,chemicals, and properties of chemicals.Libavius entitled the section of Alchemiadescribing the properties of substancesChymia. The term “chymia” eventually

became the term “chemistry.” Libavius con-tributed to the establishment of chemistryas a unique discipline deserving study in itsown right.

By the start of the seventeenth century,the stage had been set for transformation inthe study of chemistry. Up to this time,chemistry played a central role in the artsand crafts, but as a science, it had made onlymodest advances from the days of Aristotle.This started to change as individualsadopted a scientific approach and subjectedthe ideas of Aristotle to experimental test-ing. By 1700, the walls of unquestionedauthority concerning scientific thought hadbeen shattered by individuals such as Gali-leo, Isaac Newton, and Francis Bacon. Thetime was now ripe for rapid changes in thestudy of chemistry.

Introduction

Through the sixteenth century, the chemicalarts had advanced due to developments inmetallurgy, medicine, and alchemy. Physi-cians, metal smiths, tanners, textile workers,soap makers, and a host of other crafts peo-ple and artisans employed chemical knowl-edge to carry out their trades. Althoughalchemy and the teaching of the ancientGreeks persisted well into the nineteenthcentury, increasingly individuals started toconduct science by performing experiments.Natural philosophers critically examinedideas such as Aristotle’s four-element theoryand proposed alternative explanations todescribe the behavior of matter. Chemistrywas now ready to establish its own identity.Throughout Europe during the two hundred-year period from 1600 to 1800, naturalphilosophers posed questions, experi-mented, published their findings, and modi-fied their ideas until the basic principles ofchemistry started to take shape. The story ofthe developments leading up to the lateeighteenth-century chemical revolutionmust be told through the individuals respon-sible for this revolution.

The Beginnings of ModernChemistry

Jan Baptista van Helmont (1577–1644)was one of the early iatrochemists who fol-lowed the practices of Paracelsus. Van Hel-mont developed the idea of a gas, which hedistinguished from ordinary air. He used theterm “spiritus sylvestrius” to describe thegas produced during the combustion pro-cess, and he realized that this same gas alsowas produced during fermentation and whenacids reacted with sea shells. To van Hel-mont, a gas was a substance that could notbe retained in a vessel and was invisible. VanHelmont considered air and water as thebasic elements of matter, and so he neverassociated different gases as new substances.Because he could condense water vapor butnot ordinary air, he used the Greek word forchaos, “khaos.” The Greek “kh” is phoneti-cally pronounced as “g” in Flemish, and theresult was the English word gas. To van Hel-mont, a gas was modified water, and waterwas the basis for chemical change. In one ofhis experiments van Helmont attempted toprove that plants were transformed water.Van Helmont planted a willow in a measured

3

The Birth of Modern Chemistry

18 The Birth of Modern Chemistry

weight of soil. Reweighing the soil and wil-low after five years of watering, van Helmontdiscovered no change in the weight of thesoil, but the willow had gained approxi-mately 80 kilograms of weight. From this,van Helmont concluded the water had beenconverted into plant mass.

The most influential figure of seven-teenth-century chemistry was Robert Boyle(1627–1691). Boyle was born in Ireland to awealthy English family; his father was theEarl of Cork. The Boyle’s family wealthallowed Robert to pursue studies in naturalphilosophy, free from worry about incometo support himself. While still in his teens,Boyle traveled to Italy and was in Florencewhen Galileo died. Boyle was greatly influ-enced by Galileo’s approach to science, andthroughout his life based his approach toscience on an experimental, mathematical-mechanical study of nature. Because of civilunrest in his native Ireland, Boyle returnedto Oxford where he was a member of thegroup known as the Invisible College. Thisgroup was an informal, independent associ-ation of thinkers devoted to developing newideas in science and philosophy.

Boyle, like many scholars of his day,studied and published in a number of areasincluding theology, philosophy, science, andpolitical thought. In the area of chemistry,Boyle, in the tradition of van Helmont, stud-ied gases. Aided by his assistant RobertHooke, Boyle used a vacuum pump to con-duct experiments in which he discovered airwas necessary for life, sound does not travelin a vacuum, and that the volume of a gas isinversely proportional to pressure. This lastdiscovery is one of the basic gas laws, andtoday is known as Boyle’s Law (see Chap-ter 9). Boyle applied his work on gases to astudy of the atmosphere and determined thedensity of air, and how atmospheric pressurechanges with elevation.

During the seventeenth century, anatomic view of nature was employed to

explain the physical properties of matter.Daniel Sennert (1572–1627) proposed fourdifferent types of atoms corresponding to theancient elements of earth, air, fire, and water.Boyle took an atomic view of nature believ-ing that matter consisted of corpuscles (anidea borrowed from Descartes), which wereparticles of different size and shape. Likevan Helmont, Boyle considered air to be onesubstance, but considered gases to be vari-ous varieties of air rather than a differentform of water. Boyle’s studies includedappreciable work on chemical analysis. Inthis area, Boyle rejected the idea that waterwas an element. In one of his many experi-ments, he boiled water in a flask for manydays, adding more water to replace waterthat boiled away. Boyle discovered sedimentin the flask after the boiling process, and hesuggested this finding might indicate waterwas a compound. He speculated that the sed-iment may have come from the glass andsuggested weighing the glass before andafter the boiling process to determine thesource of the sediment. Boyle never repeatedthe experiment to determine if the sedimentcame from the glass. One hundred years laterAntoine Lavoisier (1743–1794) repeatedBoyle’s experiment and demonstrated thatindeed the sediment came from the glassflask and had not been derived from thewater. Boyle’s lasting contribution to chem-istry was his book The Sceptical Chymistpublished in 1661. In this work, Boylesought to expose the inconsistencies andparadoxes of chemical theory of his day.Using the same approach that Galileo usedin his Dialogue, Boyle presented his ideas inthe form of a conversation between severalindividuals. One person speaks for Aristotle,another for Paracelsus, and one who ques-tions both systems by providing new expla-nations for the behavior of matter. Boylerefuted Paracelsus’ idea that salt, sulfur, andmercury were basic substances. He alsoadvanced the corpuscular theory of matter

The Birth of Modern Chemistry 19

and used this theory to differentiate mixturesfrom compounds. Boyle also performedexperiments on combustion with RobertHooke. Boyle and Hooke showed that com-bustion of charcoal and sulfur required thepresence of air, and that these substancescould ignite in a vacuum when mixed withsaltpeter (potassium nitrate, KNO3). Hence,Boyle concluded there must be somethingcommon to both air and saltpeter (which wenow know was oxygen). Boyle also did workon acids and bases and he developed indica-tors to distinguish between them.

John Mayow (1643–1679), a contem-porary of Boyle, did important work oncombustion and respiration. Mayow iso-lated gases emitted from combustion andrespiration by using a siphon to capture thegases in a vessel underwater. Mayow dis-covered that only part of a volume of air isutilized during combustion and respiration.Mayow called this part nitro-aerial particles.The work, conducted in the 1600s, pointedthe way for advances that would take placein chemical theory in the 1700s. The studyof gases, respiration, combustion, and calci-nation were areas that sparked interest.While we are familiar with the process ofcombustion, to understand chemical devel-opments in the 1700s, it is important tounderstand the process of calcination. Cal-cination is a process in which a metal isheated without melting. During the proce-dure, the metal’s volatile componentsescape. Calcination is often employed as aninitial step in the extraction of a metal fromits ore. The material left after a substancehas undergone the calcination process isknown as a calx. Combustion and calcina-tion have been known since ancient times.Combustion was considered to be a processin which something was removed from asubstance. The ancients thought that fire wasliberated when a substance burns. Paracel-sus associated sulfur with combustion, andBoyle proposed special fire particles to

explain combustion and calcination. In thisline of thinking, fire was the primary ele-ment in the combustion/calcination processand air played only a secondary role. Airwas the media that absorbed and transportedthe fire away from a substance during com-bustion. Using this interpretation, a candlegoes out when a jar is placed over it becausethe air in the jar becomes saturated with fire.Because the air cannot absorb any more fire,the candle goes out.

Using the ancient idea of fire and itsmodification over the ages, reasonableexplanations were developed for chemicalprocesses. Yet, a fundamental fact continuedto plague individuals working in chemistry.Combusted substances weighed less aftercombustion because fire had left the sub-stance. This fact was consistent with thefour-element theory. Conversely, when met-als were heated during calcination, theygained weight. If fire was supposed to leavea material, why did calcinified metals gainweight? As more emphasis was placed onquantitative methods in the 1700s, this prob-lem grew increasingly troublesome.

The Phlogiston TheoryThe phlogiston theory was an attempt to

formulate a comprehensive explanation forchemical observations made up to the startof the eighteenth century. John Becher(1635–1682) is generally given credit forlaying the foundation of the phlogiston theory. Born in Germany, Becher was a self-educated scientist, economist, and busi-nessman. Becher spent his life serving in theadministrations of European royalty, pro-moting different industrial projects, and pur-suing various other business schemes. Henever remained in place long, quickly wear-ing out his welcome with his unfulfilledpromises. Included in the latter were a per-petual motion machine and a contraption tochange sand into gold. Becher traveled


throughout Europe, always keeping one stepahead of his enemies and detractors.Although considered a charlatan by many,Becher conducted and published serious sci-entific studies. Becher proposed five funda-mental elements: air, water, fusible earth,fatty earth, and fluid earth. Fatty earth wascombustible earth, fusible earth did notburn, and fluid earth was distillable.Becher’s three earths were analogous toParacelsus’ sulfur (fatty earth), mercury(fluid earth), and salt (fusible earth) princi-ples of matter. Becher’s work summarizedmany of the ideas of his predecessors andincluded many ideas that the Chinese heldon chemical processes.

Becher did not develop his theory com-pletely; this was left to his principal disciple,the German physician George Stahl(1660–1734). Stahl changed Becher’s termfor fatty earth (terra secunda) to phlogistonand popularized the phlogiston theory. Phlo-giston is derived from the Greek word phlo-gistos meaning flammable. Stahl acceptedphlogiston as a real substance that wasresponsible for combustion. Phlogiston hadno color or taste. Materials that burnedalmost to completion, such as coal, werebelieved to be almost completely phlogiston.Substances that did not burn possessed littleor no phlogiston. Although we might bequick to dismiss the phlogiston theory today,we must remember that it was the acceptedtheory for explaining a wide range of phe-nomenon by the mid-1700s. The phlogistontheory made sense of a number of chemicalobservations and among these were:

• Combustibles lose weight when they burnbecause they lose phlogiston.

• Charcoal almost burns completely becauseit is almost entirely phlogiston.

• A flame is extinguished because the sur-rounding air becomes saturated with phlo-giston.

• Animals die in airtight spaces because theair becomes saturated with phlogiston.

• Metal calxes turn to metal when heatedwith charcoal because phlogiston transfersfrom the charcoal to calx.

While the phlogiston theory explained anumber of observations, the problem of cal-cination still remained. Why did metalsincrease in weight when heated? After all, ifphlogiston left the metal during calcination,the metal should weigh less. To accommo-date the phlogiston theory, different solu-tions were proposed to make the theorywork. This is precisely the way scienceworks. A theory develops and claimswidespread acceptance. As scientists use thetheory, observations and experimental re-sults occasionally arise that are not consis-tent with the theory. Scientists are thenforced to modify the theory in some fashionto explain their results. In the case of the cal-cination problem, some individuals pro-posed different types of phlogiston. Onetype had negative weight; therefore, when itwas liberated from the substance, the sub-stance weighed more. The idea of negativeweight may sound a bit far-fetched, but wemust remember people were still consider-ing only a few basic elements at that time.It’s easy for us to reason a balloon stays sus-pended in air because it contains helium,and the balloon sinks as it loses helium. Inessence, negative-weight phlogiston was theeighteenth-century equivalent to helium. Itwas an attempt to resolve a major inconsis-tency in accepted phlogiston theory, but asfurther experimentation on gases ensued, anew theory developed to take its place.

Pneumatic ChemistryThe stage was now set for a critical

examination of phlogiston theory and thebirth of modern chemistry. By the mid-


1700s, the study of chemistry as a distinctdiscipline had been established. Althoughthe research of Boyle, van Helmont, Stahl,and others provided a sound foundation tobuild on, it was primarily a series of exper-iments on gases that took place in the lasthalf of the eighteenth century that createdthe science of modern chemistry. Severalmen stand out as prominent figures whocontributed to the demise of the phlogistontheory. Black, Priestley, and Cavendish con-ducted important studies in Great Britain,but it was a Frenchman who revolutionizedchemistry. Antoine Lavoisier synthesizedhis work and that of other natural philoso-phers into a modern theory of chemistry.

The pneumatic chemists did experi-ments on gases. To do this work, a methodwas required to generate and isolate thegases produced in a chemical reaction. Itwas important that the gases generated dur-ing reactions not become contaminated withair. An invention that facilitated the study ofgases was the pneumatic trough (Figure3.1). Stephen Hales (1677–1761) perfectedthis device in his studies of animal and plantphysiology. Hales’ trough was nothing morethan a gun barrel bent and sealed at one end.Hale placed a substance in the sealed endand heated it. The open end of the barrel wasdirected upward into a flask filled with waterand partially submerged in a bucket ofwater. Any gas generated when a substancewas heated would be directed through thetube and would displace the water in theflask. Using Hales’ pneumatic trough, agroup of English chemists conducted impor-tant experiments that assisted Lavoisier informulating his ideas.

Joseph Black (1728–1799) conductedan important series of experiments duringwork on his doctoral dissertation inmedicine. Black was searching for a mate-rial to dissolve kidney stones. He chosemagnesia alba (magnesium carbonate), but

he found it could not be used for his primarytopic on kidney stones so he turned hisattention to stomach acidity. His studies ofmagnesia alba were based on a number ofexperiments on carbonates. Black observedthat magnesia alba reacted with acid to pro-duce a gas and a salt. Black also workedwith limestone (calcium carbonate, CaCO3)and quicklime (calcium oxide, CaO) at thesame time, and so he thought that magnesiaalba, which behaved similarly to limestone,might be another form of limestone. Blacksubjected magnesia alba to calcination andfound that it lost a substantial amount ofweight during calcination. The residue leftafter calcination did not behave at all likequicklime. Because quicklime is obtainedwhen limestone is calcined, Black con-cluded magnesium alba was not a form oflimestone. Black continued his investigationby treating both the magnesia alba and thecalcined magnesia alba with acid. He foundthat magnesia alba and acid produced a saltand gas, while the calcined magnesia alba

Figure 3.1Diagram of Pneumatic Trough (Rae Déjur)


produced the same salt but not the gas.Blacks two reactions were

acid � magnesia alba salt � gasacid � calcined magnesia alba salt

Black explained his results by determiningthat calcination of magnesia alba gives off agas, and this gas is the same gas producedwhen magnesia alba is treated with acid. Itshould be remembered that Black inter-preted his results based on phlogiston the-ory. According to phlogiston theory, themagnesia alba should have gained weightwhen it was calcinated due to phlogistonleaving the fire and flowing into the magne-sia alba. Therefore, during the calcination ofmagnesia alba, gas was given off and phlo-giston was gained. Black then devised amethod to determine the weight of the gasevolved from magnesia alba. Hales’ pneu-matic trough would not work, because thegas dissolved in water. Therefore, Blackweighed the acid and magnesia alba beforeand after they reacted. Black measured thedifference in weights as the weight of gasevolved. Once he did this, Black found thatthe change in weight was close to the weightloss when magnesia alba was calcined. Thisresult indicated that there was no weightgain due to phlogiston entering the magne-sia alba. Black could explain the changesusing conservation of mass without resort-ing to a gain of phlogiston. The changes inweight in magnesia alba could be explainedby the loss of a gas during both calcinationand reaction with an acid.

Black decided that the gas given off inthe reaction between magnesia alba and acidwas similar to one described by van Hel-mont. He coined the term “fixed air” to indi-cate that this gas was fixed or trapped inmagnesia alba. Black also recognized thatfixed air was the same gas produced in res-piration, combustion, and fermentation.Today, we know Black’s fixed air was car-

bon dioxide. Black published his work in1756 in England, but it took a number ofyears for word of his study to spread.Black’s work inspired other British scientiststo study gases. He had demonstrated thatgases had to be accounted for in chemicalreactions, especially when looking atchanges in weight. Black also showed thatchemical reactions could be explained with-out resorting to phlogiston. According to thephlogiston theory, limestone, when heated,absorbed phlogiston to become quicklime.Black showed that the production of quick-lime could be explained by the loss of fixedair from limestone. Phlogiston was notneeded to explain the process, althoughBlack never abandoned the phlogiston theory.

Joseph Black, upon completion of hisstudy on magnesia alba, received his medi-cal degree. He published very little after thisstudy. Black taught at universities in Glas-gow and Edinburgh, and continued to dosolid research, which he presented in hislectures. One area in which he did importantwork was heat and the latent heat of steam.His work in this area inspired one of his stu-dents James Watt to apply Black’s ideas inmaking improvements to the steam engine.

Similar to Black, Daniel Rutherford(1749–1819) studied gases for his medicaldegree dissertation. Rutherford found thatcommon air contained a part that supportedrespiration and a part that did not. Initially,Rutherford assumed the part that did notsupport respiration was contaminated byfixed air. Rutherford experimented andremoved the fixed air, and he discovered theuncontaminated air still did not support lifeor combustion. Rutherford assumed the gashe had isolated was ordinary air saturatedwith phlogiston; hence, it was phlogisticatedair, which he referred to as noxious air.What Rutherford had isolated was nitrogen,and he is given credit for its discovery.


Another important figure in the chemi-cal revolution was Joseph Priestley(1733–1804). Priestley was a minister bytraining who earned his living at varioustimes as a pastor, tutor, and school adminis-trator. He had no formal science training,and throughout his life he was branded adissident and radical. Early in his scientificstudies, Priestley met Benjamin Franklinwho sparked his interest in electricity. Oneyear after meeting Benjamin Franklin,Priestley wrote The History of Electricity.Priestley’s first investigations in chemistrydealt with the gas produced in a brewerynear his parish in Leeds. Priestley discov-ered he could dissolve the gas in water toproduce a pleasant tasting beverage. Priest-ley had actually produced soda water, theequivalent of sparkling water. Priestley’sprocess for producing carbonated water wasmade commercially feasible in the 1790s byJacob Schweppe (1740–1821), the founderof the company that still bears his name.Priestley’s work on electricity and gasesresulted in his election to the French Acad-emy of Sciences in 1772, and he was alsohonored by the Royal Society in 1773.

In 1772, the second Earl of Shelburneemployed Priestley as an advisor and per-sonal secretary. This position gave Priestleyample time and freedom to continue hisstudies on gases. Priestley used a strongmagnifying glass to heat substances trappedunder glass, and in this manner he was ableto liberate gases from substances. He alsoemployed a pneumatic trough using mercuryinstead of water. In this manner, Priestleywas able to isolate fixed air (CO2) and anumber of other gases. He produced nitrousoxide, N2O, which he called dephlogisti-cated nitrous air. Priestley noted dephlo-gisticated nitrous air caused people to laugh,and thus, this gas became known as laugh-ing gas. Nitrous oxide was used years laterfor anesthesia, and it is currently used as a

propellant in cans of whipped cream. Priest-ley also isolated alkaline air (ammonia, NH3)vitriolic acid air (sulfur dioxide, SO2), heavyinflammable air (carbon monoxide, CO), andnitrous acid air (nitrogen dioxide, NO2). Hisresearch on gases was summarized in aseries of volumes published between 1774and 1786 entitled Experiments and Obser-vations on Different Kinds of Air. Priestley’smost famous studies concerned the produc-tion of oxygen (Figure 3.2). Priestley heatedthe calx of mercury (mercuric oxide, HgO)and collected the gas. He observed that acandle burned brightly in the gas. At first,Priestley assumed the gas was dephlogisti-cated nitrous air (N2O) because of results hehad obtained previously. When Priestley fur-ther examined the gas and its solubility prop-erties, he decided it was not dephlogisticatednitrous air but an entirely new gas. Henoticed that mice lived longer in the gas, andthe gas did not lose its ability to supportcombustion as quickly as dephlogisticatednitrous oxide. Priestley came to realize hewas dealing with a gas similar to commonair but with a greater ability to support com-bustion and respiration. We know that Priest-ley was dealing with oxygen, but Priestleywas a strong believer in phlogiston theory.He explained his findings in terms of phlogiston. Priestley assumed oxygen wasdephlogisticated air, common air in whichphlogiston had been removed. To Priestley,dephlogisticated air was air that had the abil-ity to absorb a substantial amount of phlo-giston. Priestley was conservative in theinterpretation of his results, which was amajor difference between Priestley andLavoisier. Additionally, Lavoisier empha-sized quantitative measurements in his studyof gases while Priestley placed secondaryimportance on quantitative relationships.

Priestley left the employment of LordShelburne in 1780 having made his majorscientific discoveries. Priestley’s political


and religious ideas continued to cause himtrouble within England. He had supportedthe colonists in the Revolutionary War andsupported the French Revolution. Hisunorthodox Unitarian views ran againstthose of the Church of England. In 1782, hepublished a book entitled History of Cor-ruptions of Christianity, which was bannedby the Church of England. Eventually, thesituation got so bad for Priestley that hishome and church were burned by an angrymob in 1791. Priestley was forced to moveto the United States in 1794, where he livedthe last decade of his life in Pennsylvania.In 1874, on the 100th anniversary of Priest-

ley’s discovery of oxygen, a celebration washeld, and the event marked the founding ofthe American Chemical Society. The high-est honor this society gives is the PriestleyAward.

While Joseph Priestley generally re-ceives credit for the discovery of oxygen, itwas Carl Sheele (1742–1786) who probablydeserves credit as the first to isolate oxygen.Scheele was a Swede who became inter-ested in chemistry during his apprenticeshipas an apothecary. Scheele acquired his ownpharmacy and during the process researchedthe production of different medicines.Scheele realized that common air contained

Figure 3.2Apparatus used by Priestley for investigations on gases. The tub is Priestley’s pneumatic trough.From Experiments on Different Kinds of Air. Image courtesy of School of Chemical Sciences,University of Illinois at Urbana-Champaign.


a component that was responsible for com-bustion and termed this component fire air.Scheele succeeded in isolating fire air byreacting nitric acid (HNO3) and potassiumhydroxide (KOH) to obtain potassiumnitrate (KNO3) and then heated the potas-sium nitrate and used a salt to absorb thenitrogen oxides from the gas produced. Theprocess resulted in fire air. Unfortunately, bythe time Scheele published his results in1777, Priestley’s results were well known. Inaddition to his work on gases, Scheele madenumerous other discoveries in the field ofchemistry. He is credited with the discoveryof a number of organic acids including tar-taric, oxalic, uric, lactic, and citric. Scheelealso discovered numerous acids such ashydrofluoric, hydrocyanic, and arsenic. Thefact that Scheele lived far from the centersof scientific activity and died at the age of44 contributed to his lack of recognition.Scheele was a firm believer in phlogistontheory and explained all his findings withinthe phlogiston framework.

Another of the English pneumaticchemists who made important discoverieswas Henry Cavendish (1731–1810). Foryears, it was known that gas was producedwhen metals reacted with acid. Cavendishstudied this gas and in 1766 used the term“inflammable air” to describe the gas.Cavendish’s inflammable air was hydrogen.Because of its explosive nature, Cavendishconsidered inflammable air to be pure phlo-giston. Cavendish was able to demonstratewater was a compound by combininginflammable air (hydrogen) and dephlogis-ticated air (oxygen) and using a spark toignite the mixture to produce water.

By 1780, a number of different types ofairs had been isolated and discoveries madethat raised serious questions about thenature of matter. Air, earth, and water wereshown to be compounds and mixtures ratherthan the basic elements from ancient times.

Fire was associated with the mysteriousmaterial known as phlogiston. This phlogis-ton could not be detected with the sensesand had either positive, negative, or noweight depending on how experiments wereinterpreted. The discoveries of Boyle, Black,Cavendish, Scheele, Priestley, and other pio-neer chemists were interpreted using thephlogiston theory. Yet, phlogiston theory lefttoo many questions unanswered. Lavoisier,like his contemporaries, was initially a con-firmed phlogistonist, but his interpretationof his own work and that of others lead himto abandon phlogiston. Lavoisier’s work fol-lowed Newton’s revelations in physics andpreceded Darwin’s theory of evolution andestablished the modern science of chemistryat the end of the eighteenth century.

LavoisierAntoine Laurent Lavoisier was a mem-

ber of the French nobility; his father was awealthy Parisian lawyer. Educated at thebest French schools, Lavoisier developed adeep interest in science. Lavoisier showedearly promise in geology and astronomy, buthe also was interested in chemistry, whichhe studied under Giullame Francois Rouelle(1703–1770). Lavoisier was wealthy enoughto fund his own scientific studies, but he alsowas involved in significant governmentwork and business dealings. In 1768,Lavoisier purchased a position in a privateFrench tax-collecting firm known as FermeGénérale. It was through this firm thatLavoisier met his wife, Marie-Anne, whoassisted Lavoisier in his scientific endeavorsthroughout his life.

Lavoisier’s serious work in chemistrystarted around 1770. Lavoisier conductedstudies on the combustion of sulfur andphosphorous and focused on the weightincrease during the combustion process.Lavoisier’s work on combustion led him to


believe that a weight gain may accompanyall combustion and calcination reactions.Lavoisier related his ideas on air being fixedin substances during combustion and calci-nation in a note dated November 1, 1772, tothe French Academy.

In 1774, Lavoisier’s work OpusclesPhysiques et Chymiques (Physical andChemical Tracts) appeared in print. Thisbook presented Lavoisier’s findings and wasbased largely on repeating and expandingupon the work of Joseph Black. Opusclesindicated that Lavoisier believed calcinationand combustion involved combination withsome component of air. Initially, Lavoisierthought fixed air was the gas involved withcalcination and combustion. At this point,Lavoisier was still unclear about the gasinvolved in combustion and calcination, andhe turned to the work of other naturalphilosophers to point the way. This was typ-ical of Lavoisier. Much of Lavoisier’s labo-ratory work was based on repeating thework of other scientists. Using this methodenabled Lavoisier to observe a wide rangeof chemical observations and tie thesetogether without resorting to phlogiston.After publication of Opuscles, Lavoisier metwith Priestley in Paris during the fall of1774. Lavoisier also was familiar with thework of French chemist Pierre Bayen(1725–1798); Bayen also rejected the ideaof phlogiston. Bayen, like Priestley, hadreported that the calx of mercury whenheated produced fixed air in the presence ofcarbon and another gas when carbon wasnot present. Lavoisier also received a letterfrom Scheele on his work in the fall of 1774.

Lavoisier resumed his work on com-bustion and calxes in the spring of 1775. Atthis time, Lavoisier wrote his initial paperentitled On the Nature of Principles whichCombine with Metals During Calcinationand Increases Their Weight. A revised ver-sion of this paper did not appear until 1778.

In this paper, Lavoisier said air consists oftwo elastic fluids and distinguished betweenactive and inactive components of air.Lavoisier’s reactive air was the same asPriestley’s dephlogisticated air (oxygen).Lavoisier slighted Priestley by not givingPriestley due credit for helping him interprethis work. Priestley felt Lavoisier was claim-ing credit for his work, but Lavoisier wasequally adamant in staking claim to an orig-inal interpretation of the role of dephlogis-ticated air in combustion and calcination.This claim irritated Priestley immensely andresulted in a bitter rivalry between the twofor the rest of their lives.

In 1775, Lavoisier was appointed Chiefof the Royal Gunpowder Administration. Inthis position, Lavoisier reformulated Frenchgunpowders and improved their quality tobecome the best in Europe. At his arsenallab, Lavoisier continued his chemical exper-iments. Lavoisier focused much of his atten-tion on acids. He was the first to use theterm “oxygine” in September of 1777, aword derived from the Greek terms oxys andgen meaning acid producing. Lavoisierbelieved incorrectly oxygine was a basiccomponent in all acids.

Another problem Lavoisier wrestledwith was the reaction of dephlogisticated airand inflammable air (oxygen and hydrogen)to produce water. Lavoisier repeatedCavendish’s experiments. He also per-formed an experiment where he passedsteam through a heated iron rifle barrelobtaining hydrogen and iron calx (rust). Bydecomposing water into hydrogen and oxy-gen, Lavoisier claimed credit for demon-strating water was a compound. Again,others such as Cavendish, Priestley, andJames Watt had previously shown water wasa compound, but it was Lavoisier whoexplained the true nature of water withoutusing phlogiston. Lavoisier conductednumerous experiments with different gases,


isolating them and studying their character-istics (Figure 3.3). By 1780, Lavoisier waswell into developing his new philosophy toexplain chemical reactions. He publishedReflections on Phlogiston in 1783. In thiswork he attacked the phlogiston theory andintroduced his own oxygen theory to explainchemical reactions.

Lavoisier’s ideas met initial resistancefrom the established chemical community.After all, phlogiston had been used for acentury to explain chemistry, and naturalphilosophers had explained their work basedon phlogiston theory. The leading chemistsof the day, including Richard Kirwan(1733–1812), Priestley, Cavendish, and Tor-bern Bergman (1735–1784), not onlyrejected Lavoisier’s ideas, but they also wereantagonistic toward him. This was not to say

that Lavoisier did not have support. Otherchemists had questioned phlogiston, butthey could not put forth an acceptable alter-native theory. Lavoisier’s strongest discipleswere several of his countrymen. The Frenchchemists Claude Louis Berthollet (1749–1822), Guyton de Morveau (1737–1816),and Antoine Francois de Fourçroy(1755–1809) joined Lavoisier in promotinghis new chemistry. In addition to Lavoisier’snew theoretical explanation, he and his col-leagues proposed a new system of nomen-clature. Led by de Morveau, the Frenchassociates proposed a new system of nam-ing chemicals based on their elemental com-position. At the time substances were namedusing archaic names passed down throughthe ages. Names such as fuming liquor ofLibavius (tin tetrachloride, SnCl4), Epsom

Figure 3.3Lavoisier conducting human respiration experiment while Madame Lavoisier takes notes.Image from Edgar Fahs Smith Collection, University of Pennsylvania Library.


salt (magnesium sulfate, MgSO4), azote(nitrogen), and acid of salt (hydrochloricacid, HCl) gave no indication of the truecomposition of substances. In their Methodsof Chemical Nomenclature, published in1787, the French chemists proposed a sys-tem consisting of Latin and Greek namesusing prefixes and suffixes to distinguish thecomposition of substances. Using the newnomenclature, chemists now had a languagethat told them what a substance was madeof and pointed the way on how to make thesubstance. The supporters of phlogistonwere reluctant to accept the system and ini-tially rejected the new nomenclature.

Lavoisier summarized his ideas devel-oped over the previous twenty years in hisseminal 1789 book Traité Élémentaire deChimie (Elements of Chemistry). This workpresented his findings on gases and the roleof heat in chemical reactions. He explainedhis oxygen theory and how this theory wassuperior to phlogiston theory. Lavoisierestablished the concept of a chemical ele-ment as a substance that could not be bro-ken down by chemical means or made fromother chemicals. Lavoisier also presented atable of thirty-three elements. The thirty-three elements mistakenly included lightand caloric (heat). Lavoisier put forth themodern concept of a chemical reaction, theimportance of quantitative measurement,and the principle of conservation of mass.The final part of Lavoisier’s book presentedchemical methods, a sort of cookbook forperforming experiments.

Lavoisier’s Elements of Chemistry syn-thesized and explained in a coherent man-ner the chemistry of his day. The work wasreadily accepted by many as a substantialimprovement over the phlogiston theory.Young chemists quickly adopted Lavoisier’sideas and new nomenclature rather than try-ing to fit their work within a phlogistonframework. Upon reading Lavoisier’s book,

many of the “old guard” became converts.Lavoisier’s ideas spread quickly to othercountries. English, Spanish, German, Ital-ian, and Dutch translations of Elements ofChemistry were made within a few years.Furthermore, Lavoisier actively promotedhis new chemical system in lectures,reviews, and by lobbying fellow chemists.

Lavoisier’s greatest ally in his campaignto promote the new chemical system wasMarie-Anne Lavoisier. In one episode,Marie-Anne staged a scene where she pub-licly burned the works of Stahl and otherbooks promoting phlogiston. Marie-Anneserved Antoine in several important capaci-ties. As a gifted linguist and artist, she trans-lated English works and illustrated hispublications. She served as a personal sec-retary transcribing notes, editing papers, andassisting in his laboratory experiments.

Lavoisier, by all accounts, was not ahumble person. Some suspect he took creditfor the work of others, was arrogant, and didlittle original work. While true in somerespects, Lavoisier nevertheless synthesizedthe vast wealth of chemical knowledge thathad accumulated by the end of the eigh-teenth century. Lavoisier never discoveredan element or isolated a new gas, but whathe accomplished was far more important.Lavoisier established chemistry as a modernscience and laid the foundation for the quan-titative experimental methods of chemistry.Lavoisier also served France throughout hislife as a public servant. His work on Frenchgunpowder has already been mentioned, butLavoisier also worked on reform in agricul-ture, the measurement system (he was anearly proponent of the metric system), penalsystems, and medicine. Unfortunately,Lavoisier fell victim to the Reign of Terrorof the French Revolution. As a partner ofFerme Générale, he was associated with thisdespised tax-collecting agency. BecauseLavoisier was a member of the French

nobility, he received little sympathy from therevolutionaries. Although the case againsthim was weak, Lavoisier was convicted andexecuted by the guillotine on May 8, 1794.

By the end of the eighteenth century,chemistry had established itself as a modernscience. Conservation of mass and theaccounting of reactants and products in achemical balance sheet, although establishedearly in the century, was perfected byLavoisier as a unifying principle in chemistry.Lavoisier used conservation of mass to drawconclusions, devise new experiments, andtransform ideas in chemistry. In this manner,Lavoisier established a method for examiningchemical reactions that applies to this day.Like Lavoisier, modern chemists work withina theoretical framework and are guided bygeneral principles as they conduct experi-ments. Experimental results collectively areused to lend support or ultimately refuteaccepted theories. Lavoisier’s experiments ledhim away from the phlogiston theory, and in

so doing created another framework thatguided chemical thought. Lavoisier estab-lished modern chemical reasoning and amodel for modern chemists. Lavoisier’sexperiments often required the modificationof equipment or fabrication of new equip-ment. Likewise, advances in modern scienceare dependent on advances in equipment andtechnology. Lavoisier also stressed the impor-tance of language in presenting the results ofscience. Chemical nomenclature today ishighly specific with the names of substancesportraying their chemical composition usinga precise language. Finally, Lavoisier demon-strated how chemical knowledge could beused to better society and promote industry.Lavoisier applied chemistry to real-worldproblems such as improving agriculture,manufacturing gunpowder, or developinglighter-than-air balloons. Most modernchemists operate in just this manner, apply-ing chemical principles to a endless multitudeof practical problems.


IntroductionAs the eighteenth century drew to a close,Lavoisier’s new system of chemistry hadwon the day. Mysterious phlogiston was nolonger needed to interpret chemical phe-nomenon, yet phlogiston still had a fewstaunch adherents such as Joseph Priestley.Priestley persisted in his belief of phlogis-ton until his death in 1804. By this time,however, a whole new school of chemistswas busy wrestling with problems raisedby Lavoisier’s new chemistry. Foremostamong these problems was determining thecorrect composition of substances andassigning the correct relative masses to thechemical elements. The relative mass refersto the ratio of the mass of one element toanother, for example, how many timesheavier is oxygen than hydrogen. The com-position and mass problems were closelyrelated. Without knowing the exact com-position of a substance, it was impossibleto determine accurately the mass of indi-vidual elements. Likewise, inaccurate rela-tive masses of the elements brought intoquestion the exact composition of a sub-stance. The mass/composition dilemmawas critical to advancing chemistry and

drew the attention of chemists at the startof the nineteenth century. Atoms wereadopted and gradually accepted during thefirst part of the century. By the end of the century, the atomic model was one of the fundamental theories in science andfundamental particles comprising the atombegan to be discovered.

The Law of Definite Proportions

The problem of chemical compositionwas directly related to that of chemical com-bination. Combining chemicals to producenew substances had been a goal of individ-uals throughout history. Lavoisier’s workemphasized the quantitative study of chem-ical combination. Chemists sought to deter-mine the proportion in which chemicalscombined, and how much of one substanceit took to react with another. Rather thancombining substances using the alchemist’strial and error method, chemists sought todetermine the specific ratio of chemicalsinvolved in chemical reactions.

Jermias Benjamin Richter (1762–1807) was one of the first chemists to focus

4

The Atom

32 The Atom

attention on the relative weights of the ele-ments. Richter conducted experiments inwhich he determined how much acid ofvarious types it took to neutralize bases.Richter was searching for the underlyingmathematical relationships inherent inchemical reactions. Richter used measure-ments, math, and observations to decipherchemical reactions. He coined the word“stoichiometry,” which is still used todescribe mass balances associated withchemical reactions. Richter based his cal-culated atomic weights on equivalences, orhow much of one substance would reactwith another. A specific amount of sulfuricacid was used as a reference. Ernest Gott-fried Fischer (1754–1831) brought atten-tion to Richter’s work when he publishedhis German translations of Berthollet’sstudy on chemical combinations.

According to Lavoisier’s discipleBerthollet, when elements combined toform compounds, the compounds containedvariable proportions of elements withindefined limits. Berthollet believed the com-position of a compound depended on howmuch of each element was present when thecompound was formed. If salt happened tobe produced by a process with amplesodium and little chlorine, then it would beenriched in sodium. Because of Berthollet’sstature as a chemist, his views and ideas car-ried great weight at the turn of the nine-teenth century. He had based his ideas onyears of study. Berthollet had accompaniedNapoleon to Egypt and noticed that in saltlakes sodium chloride reacted with lime-stone (calcium carbonate) deposits to givesodium carbonate and calcium chloride.Berthollet knew from his lab work in Francethat solutions of calcium chloride andsodium carbonate reacted to give calciumcarbonate and sodium chloride. He nowobserved the reverse reaction occurring inEgypt’s salt lakes.

This observation provided evidence toadvance the idea of reversible reactions.Berthollet analyzed many substances andfound their composition did vary within alimited range strengthening his convictionthat substances did not display constantcomposition. Joseph Louis Proust (1754–1826) opposed Berthollet’s views on thecomposition of substances. Proust’s chemi-cal education began as an apprentice to hisapothecary father in Angers in westernFrance. Subsequently, Proust moved to Parisand left for Spain during the French Revo-lution. In Madrid, Proust held several aca-demic positions and conducted studies in awell-equipped laboratory. Through his anal-yses of substances, Proust found that com-pounds had a fixed and invariable chemicalcomposition. Proust based his findingslargely on careful studies of various metalcompounds of sulfates, carbonates, andoxides. Proust discovered that laboratory-prepared copper carbonate was no differentthan copper carbonate found in nature.Proust expressed his views clearly on thecomposition of compounds:

The properties of true compounds are invari-able, as is the ratio of their constituents. Betweenpole and pole, they are found identical in thesetwo respects; their appearance may vary owingto the manner of aggregation, but their proper-ties never.

Proust put forth his views starting in 1799and for the next ten years attacked Berthol-

Lake Reaction

2NaCl � CaCO r Na CO � CaCl(aq) 3(s) 2 3(s) 2(aq)

sodium chloride � calcium carbonate

r sodium carbonate � calcium chloride

Lab Reaction

CaCl � Na CO r CaCO � 2NaCl2(aq) 2 3(aq) 3(s) (aq)

calcium chloride � sodium chloride

r calcium carbonate � sodium chloride

The Atom 33

let’s ideas on variable composition. Proustdemonstrated that Berthollet’s results werebased on analyzing substances with impuri-ties and pointed out several other flaws inBerthollet’s methods. Berthollet counteredwith his own explanations, but by 1808 headopted Proust’s view. Proust’s Law, alsoknow as the Law of Definite Proportions,was one of the cornerstones in the develop-ment of modern chemistry. The Law of Def-inite Proportions meant that compounds hada fixed chemical composition that could beused to help determine the relative massesof elements making up the compound.

Dalton and the Birth of theAtomic Theory

John Dalton (1766–1844) was a Quakerand largely self-educated schoolteacher whodeveloped an interest in chemistry as aresult of his work in meteorology. His earlyscientific work included essays on meteoro-logical observations, equipment, and color-blindness (Daltonism). As a result of hisinterest in meteorology and the atmosphere,Dalton turned his attention to a study ofgases. He discovered that the vapor pressureof water increases with temperature. Daltonconsidered air to be a mixture of gasesrather than a chemical compound, and heformulated the concept of partial pressure.Dalton determined that the total pressure ofair was the sum of the individual pressuresexerted by individual gases in a mixture.Furthermore, the pressure exerted by eachgas was independent of the other gases pres-ent in the mixture.

One question that Dalton consideredand which guided his formulation of theatomic theory was why the gases in airformed a homogenous mixture and did notseparate according to density. Dalton knewthat water vapor was lighter than nitrogen,and nitrogen was lighter than oxygen. Why

then didn’t these gases form stratified, orseparate, layers in the atmosphere? The rea-son for this according to Dalton was thatgases diffused in each other. According toDalton, the forces between different gas par-ticles were responsible for the mixing ofgases in the atmosphere. Like gas particlesrepelled each other, and unlike gas particleswere neutral with respect to each other.

In attempting to explain the behavior ofgases, Dalton experimented on the solubil-ity of gases in water. Dalton’s work on sol-ubility was done with his close colleagueWilliam Henry (1774–1836). Dalton wasconvinced that the different solubility ofgases in water was due to the weight of theultimate particle of each gas, heavier parti-cles being more soluble than lighter parti-cles. Dalton needed to know the relativeweights of the different gases to support hisideas on gas solubility. This need led him todevelop a table of comparative weights ofultimate gas particles. In 1803, he presentedhis table of relative weights (Table 4.1).Over the next five years, he continued todevelop his ideas on atomism and work onatomic weights.

As a result of his work on relativeweights, Dalton formulated the Law ofMultiple Proportions, which states thatwhen elements combine to form more thanone compound, then the ratio of the massesof elements in the compounds are smallwhole number ratios of each other. Forexample, the elements carbon and oxygenform the two compounds carbon monoxide(CO) and carbon dioxide (CO2). The ratio of

Table 4.1

34 The Atom

carbon to oxygen masses using Dalton’sTable would be

Dalton’s Law of Multiple Proportions meantthat two elements combine in simple wholenumber ratios. Dalton believed that com-pounds found in nature would be simplecombinations. Hence, knowing that hydro-gen combines with oxygen to give water,Dalton’s formula for water would consist of1 H and 1 O. Its formula would be HO usingmodern nomenclature. Both Proust’s Law ofDefinite Proportions and Dalton’s Law ofMultiple Proportions are outcomes of anatomic view of nature. In 1808 Dalton pub-lished his table of relative atomic weightsalong with his ideas on atomism in A NewSystem of Chemical Philosophy.

Because Dalton did not know the chem-ical formula for compounds, he assumed thegreatest simplicity. This worked fine forsome compounds such as CO or NO, butintroduced error for other compounds, forexample, assuming water was HO. Never-theless, Dalton’s ideas laid the foundationfor the modern atomic theory. Dalton’s ideasbriefly summarized are:

1. Elements are made of tiny indivisible par-ticles called atoms.

2. Elements are characterized by a uniquemass specific to that element.

3. Atoms join together in simple whole num-ber ratios to make compounds.

4. During chemical reactions, atoms rear-range themselves to form new substances.

5. Atoms maintain their identity during thecourse of a chemical reaction.

Dalton’s work on relative weights, mul-tiple proportions, and the atomic theory didnot have an immediate effect on chemists ofhis day. Dalton’s ideas did provide a frame-work for determining the empirical formulaof compounds, but his table of relativeweights was not accurate enough to giveconsistent results. Many scientists stilldebated the existence of atoms in the secondhalf of the nineteenth century. Still, little bylittle, the atomic theory was adopted bychemists as a valid model for the basicstructure of matter. While Dalton continuedhis life as a humble tutor in Manchester,other chemists used Dalton’s ideas to estab-lish the atomic theory. Foremost amongthese was Jöns Jacob Berzelius (1779–1848) of Sweden, the foremost chemicalauthority of the first half of the nineteenthcentury.

Gay-Lussac and Avogadro’sHypothesis

Dalton’s work provided a system forrepresenting chemical reactions, but inev-itably, conflicts arose when trying to resolveDalton’s idea on chemical combination withexperimental evidence. According to Dal-ton, one volume of nitrogen gas combinedwith one volume of oxygen to give one vol-ume of nitrous gas (nitric oxide). Daltonreferred to combination of atoms as com-pound atoms. Using Dalton’s symbols, thisreaction would be represented as

But when one volume of nitrogen reactedwith one volume of oxygen, the result wastwo volumes of nitrous oxide, not one.Joseph Louis Gay-Lussac (1778–1850) hadexperimentally demonstrated that two vol-umes of nitrous gas result from combiningone volume of nitrogen with one volume of

CO CO2

C 4.3 C 4.3� � 0.78 � � 0.39

O 5.5 O 11.0

C 0.78 2in CO compared to CO � �2O 0.39 1

The Atom 35

oxygen. Dalton could not use atomic theoryto explain Gay-Lussac’s results. Berzeliusknew Gay-Lussac’s experiments were cor-rect and supported Gay-Lussac’s hypothesisthat volumes combined just like atoms. Thismeant equal volumes of gases containedequal numbers of atoms. If Gay-Lussac wascorrect, then volumes could be used todetermine the correct formulas. All onewould have to do is determine the ratio ofcombining volumes. This would be similarto the practice of using the ratio of combin-ing masses. Employing gas volumes gavechemists another method for determiningatomic masses. It must be remembered thatat the time different formulas were obtaineddepending on the atomic mass used. Forexample, the formula for water could beHO, H2O, HO2, or H2O2 depending onwhich values were used for the masses ofhydrogen and oxygen.

One way to resolve the problem of com-bining volumes was to split the combiningatoms in two before the combination. Butatoms were supposed to be indivisible, andso this was not an acceptable solution.Berzelius urged Dalton to reexamine hisideas in light of Gay-Lussac’s findings, butDalton would not accept Gay-Lussac’sresults. In 1811, Amedeo Avogadro (1776–1856) resolved the problem using both Dalton and Gay-Lussac’s ideas. Avogadrohypothesized that equal volumes of gasescontain equal numbers of molecules asopposed to atoms. Avogadro took the basicunit of a gas as a molecule not an atom. Thismeant when nitrogen and oxygen combined,the reaction could be represented as

Avogadro had no experimental evidence toback his claim, yet his hypothesis solved theproblem. It took fifty years for the scientific

community to accept Avogadro’s idea.Today, we take for granted that commongases such as hydrogen, oxygen, and nitro-gen exist as molecules rather than atoms intheir natural state.

The Divisible AtomDalton proposed that atoms are the

basic building blocks of matter in the early1800s. It is one thing to claim the existenceof atoms, but it is another to develop theconcept until it is readily accepted. A cen-tury after Lavoisier’s chemical revolutionthe true nature of the atom began to unfold.Although the question about the true natureof the atom concerned chemists, physicistsled the way in the study of the atom. Evi-dence accumulated during the 1800s indi-cated that Dalton’s atoms might not beindivisible particles. As the twentieth cen-tury unfolded, so did the mystery of theatom.

One area of research that raised ques-tions about the atom involved experimentsconducted with gas discharge tubes. Gasdischarge tubes are sealed glass tubes con-taining two metal electrodes at oppositeends of the enclosed tube (similar to smallfluorescent lights). During operation, themetal electrodes are connected to a highvoltage power source. The negativelycharged electrode is called the cathode, andthe positively charged electrode is theanode. Before a gas discharge tube issealed, most of the air is pumped out of it.Alternatively, the tube may contain a gassuch as hydrogen or nitrogen under very lowpressure. During the last half of the nine-teenth century, numerous improvementswere made in gas discharge tubes.

William Crookes (1832–1919) ob-served a faint greenish glow and beam usinghis own discharge tubes called Crookestube. Crookes noted that the beam in his

36 The Atom

tubes originated from the cathode, andhence, the beams became known as cathoderays. The tubes themselves took on the namecathode-ray tubes (Figure 4.1). Crookesconducted numerous experiments with cathode-ray tubes. He observed that cathoderays moved in straight lines, but they couldbe deflected by a magnet. From this work,Crookes concluded cathode rays were somesort of particle, but other researches be-lieved cathode rays were a form of light.

Part of the problem involved in inter-preting cathode rays resulted from theambiguous results obtained by variousresearchers. For example, cathode rays werebent in a magnetic field supporting the par-ticle view of cathode rays, yet when sub-jected to an electric field, cathode rays werenot deflected. The fact that cathode rayswere not deflected in an electric field sup-ported the view that cathode rays were aform of electromagnetic radiation or light.The true nature of cathode rays was ulti-mately determined by a group of physicistsled by J. J. Thomson (1856–1940) at theCavendish Laboratory at Cambridge Uni-versity.

Joseph John Thomson was appointedhead of the prestigious Cavendish Labora-tory (endowed by the Cavendish family inhonor of Henry Cavendish) at the age of 28.

Thomson was a brilliant mathematician andexperimenter who tackled the problem ofcathode rays at the end of the nineteenthcentury. Thomson’s initial experiments ledhim to believe cathode rays were actuallynegatively charged particles. Using cathode-ray tubes in which the air had been evacu-ated until a very low pressure was attained(� 1/200 atmosphere), Thomson demon-strated cathode-rays could indeed bedeflected by an electric field. The failure ofprevious researchers to evacuate their dis-charge tubes to such low pressure had pre-vented the cathode-ray beams from beingdeflected in an electric field. Thomson wasconvinced that cathode rays were particlesor in his terms “corpuscles,” but he still hadto determine the magnitude of the chargeand mass of his corpuscle to solidify hisarguments.

Thomson and his colleagues at Caven-dish conducted numerous experiments oncathode-rays. They used different gases in their tubes, varied the voltage, and em-ployed different metals as electrodes. As theyears went by and Thomson accumulatedmore and more data, he was able to calcu-late a charge to mass ratio for the corpuscle.Thomson’s results startled the scientificcommunity. His results showed that thecharge to mass ratio (e/m) was one thousandtimes greater than the accepted value forhydrogen ions (hydrogen ions are just protons).

At the time, hydrogen ions were the small-est particle known to exist. Thomson’sresults could be interpreted as either thecharge of the corpuscle being one thousandtimes greater than the hydrogen ion or themass of the corpuscle being 1/1,000 that of

Figure 4.1Cathode-Ray Tube (Rae Déjur)

Thomson's Corpuscle Hydrogen Ion

Charge 1000e e�

Mass m m

The Atom 37

the hydrogen ion. Thomson suspected thecharge of the corpuscle was equal and oppo-site to that of the hydrogen ion, and there-fore, assumed the corpuscle had a mass of1/1,000 of a hydrogen ion. Thomson pub-lished his results in 1897. Thomson’s cor-puscles eventually came to be calledelectrons based on a term coined by GeorgeStoney (1826–1911) to designate the funda-mental unit of electricity.

While Thomson’s work proved the exis-tence of electrons, his studies still left openthe question of the exact charge and mass ofthe electron. Robert Andrew Millikan(1868–1953) tackled this problem at theUniversity of Chicago by constructing aninstrument to measure the mass and chargeof the electron (Figure 4.2). Millikan’sinstrument consisted of two parallel brassplates separated by about one centimeter. Apinhole-size opening drilled into the topplate allowed oil droplets sprayed from anatomizer to fall between the two plates. Theopening between the two plates was brightlyilluminated and a microscopic eyepieceallowed Millikan to observe individual oildroplets between the two plates. Millikancalculated the mass of oil droplets by mea-suring how long it took an oil droplet tomove between the two uncharged plated.Next, Millikan charged the top brass platepositive and the bottom plate negative. Healso used an x-ray source to negatively

charge the oil droplets when they enteredthe space between the plates. By varying thecharge, Millikan could cause the oil dropletsto rise or stay suspended between the plates.Using his instrument and knowing the massof the oil droplet enabled Millikan to calcu-late the charge of the electron, negative. In1913 Millikan determined that the charge ofan electron was �1.591 � 10�19 coulomb.This is within 1% of the currently acceptedvalue of �1.602177 � 10�19 coulomb. WithMillikan’s value for charge, the electron’smass could be determined. The value waseven smaller than given by Thomson andfound to be 1/1,835 that of the hydrogenatom. Millikan’s work conducted around1910 gave definition to Thomson’s electron.The electron was the first subatomic parti-cle to be discovered, and it opened the doorto the search for other subatomic particles.

In their studies with cathode rays,researchers observed different rays travelingin the opposite direction of cathode rays. In1907, Thomson confirmed the rays carrieda positive charge and had variable massdepending on the gas present in the cathode-ray tube. Thomson and others found the pos-itive rays were as heavy or heavier thanhydrogen atoms. In 1914, Ernest Rutherford(1871–1937) proposed that the positive rayswere composed of a particle of positivecharge as massive as the hydrogen atom.Subsequent studies on the interaction ofalpha particles with matter demonstratedthat the fundamental positive particle wasthe proton. By 1919, Rutherford was cred-ited with identifying the proton as the sec-ond fundamental particle.

The last of the three fundamental parti-cles is the neutron. Experimenters in theearly 1930s bombarded elements with alphaparticles. One type of particle produced hadthe same mass of the proton, but carried nocharge. James Chadwick (1891–1974), incollaboration with Rutherford, conducted

Figure 4.2Picture of Millikan’s Instrument (Rae Déjur)

38 The Atom

experiments to confirm the existence of theneutral particle called the neutron. Chad-wick is credited with the discovery of theneutron in 1932.

Atomic StructureWith the discovery of the neutron, a

basic atomic model consisting of three fun-damental subatomic particles was complete.While discoveries were being made on thecomposition of atoms, models were ad-vanced concerning the structure of the atom.John Dalton considered atoms to be solidspheres. In the mid-1800s, one theory con-sidered atoms to consist of vortices in anethereal continuous fluid. J.J. Thomson putforth the idea that the negative electrons ofatoms were embedded in a positive sphere.Thomson’s model was likened to raisins representing electrons in a positive blob ofpudding; the model was termed the “plum-pudding” model as shown in Figure 4.3. Thenegative and positive parts of the atom inThomson’s model canceled each other,resulting in a neutral atom.

A clearer picture of the atom began toemerge toward the end of the first decade of the twentieth century. This picture was

greatly aided by several significant discov-eries in physics. In 1895 while observing theglow produced by cathode rays from asealed Crookes tube, Wilhelm ConradRoentgen (1845–1923) noticed that nearbycrystals of barium platinocyanide glowed.Barium platinocyanide was a material usedto detect cathode rays, but in this case, thecathode rays should have been blocked bythe glass and an enclosure containing theCrookes tube. Roentgen also discovered thata photographic plate inside a desk wasexposed and contained an image of a keyresting on top of the desk. Evidently what-ever had caused the barium platinocyanideto glow could pass through the wooden deskbut not through the metal key. Roentgencoined the term “x-rays” for his newly dis-covered rays. He was awarded the firstNobel Prize in physics for his discovery in1901.

Roentgen’s discovery of x-rays stimu-lated great interest in this new form of radi-ation worldwide. Antoine Henri Becquerel(1852–1908) accidentally discovered theprocess of radioactivity while he wasstudying x-rays. Radioactivity involves thespontaneous disintegration of unstableatomic nuclei. Becquerel had stored ura-nium salts on top of photographic plates ina dark drawer. When Becquerel retrieved theplates, he noticed the plates containedimages made by the uranium salts. Bec-querel’s initial discovery in 1896 was furtherdeveloped by Marie Curie (1867–1934) andPierre Curie (1859–1906). Marie Curiecoined the word “radioactive” to describethe emission from uranium.

Three main forms of radioactive decayinvolve the emission of alpha particles, betaparticles, and gamma rays. An alpha parti-cle is equivalent to the nucleus of a heliumatom. Beta particles are nothing more thanelectrons. Gamma rays are a form of elec-tromagnetic radiation.

Figure 4.3Thomson proposed the atom consisted ofnegative electrons embedded in a positivepool, like raisins in plum pudding.

The Atom 39

Physicists and chemists doing pioneerwork on atomic structure employedradioactive substances to probe matter. Byexamining how radiation interacted withmatter, these researchers developed atomicmodels to explain their observations.Rutherford along with Johannes WilhelmGeiger (1882–1945), creator of the firstGeiger counter in 1908 to measure radia-tion, and Ernest Marsden (1889–1970) car-ried out their famous gold foil experimentthat greatly advanced our concept of theatom. Rutherford’s experimental set-up isshown in Figure 4.4.

Rutherford used foils of gold, plat-inum, tin, and copper and employed polonium as a source of alpha particles.According to Thomson’s plum puddingmodel, alpha particles would pass throughthe gold with little or no deflection. Theuniform distribution of charge in Thom-son’s model would tend to balance out thetotal deflection of a positive alpha particleas it passed through the atom. Rutherford,Geiger, and Marsden found a significant

number of alpha particles that experiencedlarge deflections of greater than 90°. Thisresult surprised Rutherford who years laterdescribed the phenomenon of widely scat-tered alpha particles “as if you had fired a15-inch shell at a piece of tissue paper andit came back and hit you.” To account forthese experimental results, Rutherford pro-posed a new atomic model in which theatom’s mass was concentrated in a smallpositively charged nucleus. The electronshovered around the nucleus at a relativelygreat distance. To place the size of Ruther-ford’s atom in perspective, consider thenucleus to be the size of the period at theend of this sentence, the electrons wouldcircle the nucleus at distances of severalmeters. If the nucleus was the size of aping-pong ball, the electrons would beabout one kilometer away. Rutherford’splanetary model represented a miniaturesolar system with the electrons rotatingaround the nucleus like the planets aroundthe sun. Most of the atom’s volume inRutherford’s model consisted of emptyspace, and this space explained the resultsfrom the gold foil experiment. Most alphaparticles passed through the foil with littleor no deflection. This occurred because thenet force on the alpha particle was close tozero as it passed through the foil. Thewidely scattered alpha particles resultedwhen the positively charged alpha particlesapproached the nucleus of a gold atom.Because like charges repel each other, thealpha particle would be deflected or evenback-scattered from the foil.

Table 4.2Forms of Radiation

Figure 4.4Rutherford’s Set-Up (Rae Déjur)

40 The Atom

Niels Bohr (1885–1962) advancedRutherford’s model by stipulating that anatom’s electrons did not occupy just anyposition around the nucleus, but theyoccupied specific orbitals to give a stableconfiguration. Bohr based his ideas on astudy of the spectrum for hydrogen. Aspectrum results when light is separatedinto its component colors by a prism.When visible light passes through a prism,a continuous spectrum similar to a rain-bow results. Light from a hydrogen dis-charge tube does not produce a continuousrainbow, but it gives a discontinuous pat-tern of lines with broad black areas sepa-rating the lines. Bohr proposed specificelectron energy levels or orbitals toaccount for the lines produced in hydro-gen’s emission spectrum. In Bohr’s model,when the hydrogen in the tube becameenergized by applying a voltage, electronsjumped from a lower energy level to ahigher level. As they moved back down,energy was released in the form of visiblelight to produce the characteristic lines ofhydrogen in its spectrum. Figure 4.5 dis-plays this graphically. While Bohr’s modelsuccessfully explained hydrogen’s linespectrum, it could not account for thespectra obtained for gases with more thanone electron such as that for helium.Nonetheless, Bohr had shown that matterwas discontinuous and introduced the ideaof the quantum nature of the atom.

Quantum MechanicsWorld War I delayed advances in basic

research on the atom for several years as sci-entists turned their attention to nationalsecurity. The end of WWI brought renewedinterest and further development on atomictheory in the 1920s. Louis de Broglie(1892–1977) from France introduced theidea of wave properties of matter in his doc-toral thesis in 1924. The Austrian EdwinSchrödinger (1887–1961) formulated wavemechanics to explain the behavior of elec-trons in terms of wave properties. WernerHeisenberg (1901–1976) from Germanydeveloped an atomic theory using matrixmathematics. Schrödinger and Heisenberg’swork between 1925 and 1928 contributed tothe formulation of quantum mechanics. Inquantum mechanics, the position and energyof electrons are described in terms of prob-ability. Rather than saying an electron islocated at a specific position, quantum num-bers are used to describe each electron in anatom. The numbers give the probability of finding an electron in some region surrounding the nucleus. The quantum-mechanical model is a highly mathematicaland nonintuitive picture of the atom. Ratherthan thinking of electrons orbiting thenucleus, the quantum mechanical pictureplaces electrons in clouds of differentshapes and sizes around the nucleus. Theseclouds are referred to as orbitals. Severalorbital are shown in Figure 4.6.

Figure 4.5Electrons Changing Energy Level in theBohr Atom

Figure 4.6Representation of Several Different Orbitals

The Atom 41

In the quantum-mechanical model of theatom, electrons surround the nucleus inorbitals, which are in regions called shells.Shells can be thought of as a group of orbitals,although the first shell contains only oneorbital. The number of electrons surroundingthe nucleus equals the number of protons inthe nucleus forming a neutral atom. An atomicorbital is a volume of space surrounding thenucleus where there is a certain probability offinding the electron. The electrons surround-ing the nucleus of a stable atom are located inthe orbitals closest to the nucleus. The orbitalclosest to the nucleus has the lowest energylevel, and the energy levels increase the far-ther the orbital is from the nucleus. The loca-tion of electrons in atoms should always bethought of in terms of electrons occupying thelowest energy level available. An analogy canbe drawn to the gravitational potential energyof an object. If you hold an object such as thisbook above the floor, the book has a certainamount of gravitational potential energy withrespect to the floor. When you let go of thebook, it falls to the floor where its potentialenergy is zero. The books rests on the floor ina stable position. If we pick the book up, weused work to increase its potential energy andit no longer rests in a stable position. In a sim-ilar fashion we can think of electrons as rest-ing in stable shells close to the nucleus wherethe energy is minimized.

Orbitals can be characterized by quan-tum numbers. One way to think of quantumnumbers is as the addresses of electrons sur-rounding the nucleus. Just as an addressallows us to know where we are likely tofind a person, quantum numbers give usinformation about the location of electrons.We cannot be sure a person will be home attheir address, and we can never be exactlysure an electron will be located within theorbital. The address of electrons in an atomis given by using quantum numbers and isreferred to as the electron configuration. The

electron configuration of each element pro-vides valuable information about the ele-ment’s chemical properties.

Four quantum numbers characterize theelectrons in an atom. These are:

1. The principal quantum number symbol-ized by the letter n. The principal quantumnumber tells which shell the electron is inand can take on integral values startingwith 1. The higher the principal quantumnumber, the farther the orbital is from thenucleus. A higher principal quantum num-ber also indicates a higher energy level.Letters may also be used to designate theshell. Orbitals within the same shell havethe same principal quantum number.

2. The angular momentum quantum numberis symbolized by the letter l. The angularmomentum quantum numbers gives theshape of an orbital. Values for l depend onthe principal quantum number and canassume values from 0 to n–1. When l � 0,the shape of the orbital is spherical, and when l � 1, the shape is a three-dimensional figure eight (see representa-tions in Figure 4.6). As the value of lincreases, more complex orbital shapesresult. Letter values are traditionally usedto designate l values.

The letter designation originated fromspectral lines obtained from the elements.Lines were characterized as sharp, princi-pal, or diffuse (s,p,d). Lines after d werelabeled alphabetically starting from theletter f.

3. The magnetic quantum number symbol-ized by ml. The magnetic quantum num-

Principal Quantum Numbersn 1 2 3 4 5 6

Shell K L M N O PEnergy r

l value 0 1 2 3 4Letter s p d f g

Energy r

42 The Atom

ber gives the orientation of the orbital inspace. Allowable values for ml are integervalues from �l to �1, including 0. Hence,when l � 0, then ml can only be 0, andwhen ml � 1, ml can be �1, 0, or �1.

4. The spin quantum number symbolized byms. The spin quantum number tells howthe electron spins around its own axis. Itsvalue can be �1⁄

2or �1⁄

2, indicating a spin

in either a clockwise or counterclockwisedirection.

Using the above definitions for the fourquantum numbers, we can list what combi-nations of quantum numbers are possible. Abasic rule when working with quantumnumbers is that no two electrons in the sameatom can have an identical set of quantumnumbers. This rule is known as the PauliExclusion Principle named after WolfgangPauli (1900–1958). For example, when n �1, l and ml can be only 0 and ms can be �1⁄

2

or �1⁄2. This means the K shell can hold a

maximum of two electrons. The two elec-trons would have quantum numbers of1,0,0,�1⁄

2and 1,0,0,�1⁄

2, respectively. We

see that the opposite spin of the two elec-trons in the K orbital means the electrons donot violate the Pauli Exclusion Principle.Possible values for quantum numbers andthe maximum number of electrons eachorbital can hold are given in Table 4.3 andshown in Figure 4.7.

Using quantum numbers, electron con-figurations can be written for the elements.To write electron configurations we need toremember that electrons assume the moststable configuration by occupying the low-est energy levels. We also must rememberthat no two electrons can have the same fourquantum numbers. Let’s start with the sim-plest element hydrogen and write severalconfigurations to demonstrate the process.Hydrogen has only one electron. This elec-tron will most likely be found in the K (n �1) shell. The l value must also be 0. We don’t

need to consider the ml and ms since the for-mer is 0 and the latter can arbitrarily beassigned either plus or minus 1/2. The elec-tron configuration for hydrogen is written as

Helium has two electrons. Both elec-trons have n values of 1, l values of 0, and ml

values of 0, but they must have oppositespins to conform to the Pauli ExclusionPrinciple. Therefore, the electron configura-tion of He is 1s2. Lithium with an atomicnumber of 3 has three electrons. The firsttwo electrons occupy the 1s orbital just likehelium. Because the 1s orbital in the K shellcan hold only two electrons, the third elec-tron must go into the next lowest energylevel available in the L shell. This is wherethe n value is 2 and the l value is 0. The elec-tron configuration of lithium is thus 1s2 2s1.As the atomic number increases movingthrough the periodic table, each element hasone more electron than the previous ele-ment. Each additional electron goes into theorbital with the lowest energy. The electronconfiguration for the first twenty elements isgiven in Table 4.4.

An examination of the electron config-uration of the first 18 elements indicates thatthe addition of each electron follows a reg-

Figure 4.7Maximum Numbers of Electrons perOrbital

The Atom 43

ular pattern. One would expect that the elec-tron configuration of element 19, potassium,would be 1s2 2s22p6 3s23p63d1, but Table 4.4indicates the 19th electron goes into the 4sorbital rather than 3d. The same is true forcalcium with the 20th electron also fallinginto the 4s orbital rather than 3d. It is notuntil element 21, scandium, that the 3dbegins to fill up. The reason for this and sev-

eral other irregularities has to do with elec-tron repulsion. It is easier for the electrons togo into the 4th shell rather than the partiallyfilled 3rd shell.

The quantum mechanical model andthe electron configurations of the elementsprovide the basis for explaining manyaspects of chemistry. Particularly importantare the electrons in the outermost orbital of

Table 4.3Quantum Numbers

44 The Atom

an element. These electrons, known as thevalence electrons, are responsible for thechemical properties elements display, bond-ing, the periodic table, and many chemicalprinciples. The role the valence electronsplay will be discussed more fully whenthese topics are addressed.

A Modern View of the AtomSince the beginning of the twentieth cen-

tury when Thomson identified the electron aspart of the atom, hundreds of other subatomicparticles have been identified. Modern parti-cle accelerators and atom smashers haveenabled scientists to catalog a wide variety ofparticles each characterized by their uniquemass, charge, and spin. Murray Gell-Mann(1929–) and George Zweig (1937–) indepen-dently proposed the existence of quarks in1963 as the basic constituents of matter. Theterm “quark” was coined by Gell-Mann whoobtained it from a line in James Joyce’s 1939novel Finnegans Wake: “Three quarks forMuster Mark!” According to the quark the-ory, small hypothesized particles call quarkscombine to produce subatomic particles.

Quarks carry a fractional charge of �1/3 or�2⁄

3. Six “flavors” or types of quarks make

up all subatomic particles. Each flavor ofquark can be further classified as having oneof three colors. These are not colors or flavorsas commonly thought of, but part of a classi-fication scheme used to explain how matterbehaves. The language of quarks makes themseem like some creation of fantasy, but thequark theory can be used to explain manyproperties of subatomic particles. For exam-ple, a proton can be considered to be made oftwo up quarks and a down quark, and a neu-tron of two down quarks and an up quark(Figure 4.8). Quark flavors and charges aregiven in Table 4.5.

Table 4.4Electron Configurations of the First 20 Elements

Figure 4.8Structure of Protons and Neutrons

Table 4.5Flavor and Charge of Quarks

The Atom 45

Electrons are not thought to be made ofquarks, but of another class of elementaryparticles called leptons. Leptons do notseem to be made of other particles and areeither neutral or carry a unit charge.

Atomic Mass, Atomic Numbers, and Isotopes

The most basic unit of a chemical ele-ment that can undergo chemical change isan atom. Atoms of any element are identi-fied by the number of protons and neutronsin the nucleus. The number of protons in thenucleus of an element is given by theatomic number. Hydrogen has one protonin its nucleus so its atomic number is one.The atomic number of carbon is six, becauseeach carbon atom contains six protons in itsnucleus. Besides protons, the nucleus con-tains neutrons. The number of protons plusthe number of neutrons is the mass numberof an element. A standard method of sym-bolizing an element is to write the elementswith the mass number written as a super-script and the atomic number as a subscript.Carbon-12 would be written as

An element’s atomic number is con-stant, but most elements have varying massnumbers. For example, all hydrogen atomshave an atomic number of one, and almostall hydrogen atoms have a mass number ofone. This means most hydrogen atoms haveno neutrons. Some hydrogen atoms havemass numbers of two or three, with one andtwo neutrons, respectively. Different formsof the same element that have different massnumbers are known as isotopes. Three iso-topes of hydrogen are symbolized as

Isotopes of elements are identified bytheir mass numbers. Hence, the isotope C-14 is the form of carbon that containseight neutrons. Isotopes of an element havesimilar chemical properties. Therefore,when we eat a bowl of cereal and incorpo-rate carbon in our bodies, we are assimilat-ing C-14 (and other carbon isotopes) alongwith the common form of carbon, C-12.

The mass number gives the total num-ber of protons and neutrons in an atom of anelement, but it does not convey the absolutemass of the atom. To work with the massesof elements, we use comparative masses.Initially, Dalton and the other pioneers ofthe atomic theory used the lightest elementhydrogen and compared masses of other ele-ments to hydrogen. The modern system usesC-12 as the standard and defines one atomicmass unit (amu) as 1/12 the mass of one C-12 atom. One amu is approximately 1.66 �10–24g. This standard means the masses ofindividual protons and neutrons are slightlymore than 1 amu as shown in Table 4.6.

The atomic mass of an element is actu-ally the average mass of atoms of that elementin atomic mass units. Average mass must beused because atoms of any elements exist asdifferent isotopes with different masses. Thus,the atomic mass of carbon is slightly higherthan 12 with a value of 12.011. This isbecause most carbon (99%) exists as C-12,but heavier forms of carbon also exist.

IonsMass numbers and atomic masses focus

on the nucleus of the atom. Little has been

12C6

2 31H H H1 1 1

Hydrogen Deuterium Tritium

Table 4.6Masses of Fundamental Atomic Particles

46 The Atom

said about electrons because the mass ofelectrons is negligible compared to the massof protons and neutrons. Electrons have onlyabout 1/2,000 the mass of protons and neu-trons. For an atom to remain electricallyneutral, the number of electrons must equalthe number of protons. When a neutral atomgains or loses electrons, a charged particlecalled an ion results. This process is knownas ionization. Positive ions are referred toas cations, and negative ions are callanions. Atoms gain or lose one or moreelectrons to become ions. Ions are writtenusing the elements symbol and writing thecharge using a superscript. A simple equa-tion can be written to symbolize the ioniza-tion process. For example, when lithiumloses an electron to form Li�, the equation is

Li Li� � e–

Several other equations representing theionization process are

F � e– F–

Ca Ca2 � 2e–

O � 2e– O2–

Atoms do not gain or lose electrons in ahaphazard fashion. Chemical reactionsinvolve the loss and gain of electrons. Infact, the chemical behavior of all substancesis dictated by how the substances’ valenceelectrons interact when the substances arebrought together. This subject is introducedin the next chapter in a preliminary discus-sion of compounds, and the concepts formthe basis of our discussion on bonding.

SummaryThe resurrection of Democritus’atom by

Dalton at the start of the nineteenth-century was one of those seminal events inthe history of science. Dalton’s atom, whilenot immediately accepted, nonetheless pro-

vided an organizing framework to guidechemical thought for the next one hundredyears. Dalton’s atomic theory followedclosely on the heels of Lavoisier’s work, andthe work of these two scientists lit the paththat led to the development of modern chem-istry. The atomic theory eventually becamethe touchstone by which new chemical theo-ries were judged, and it played a fundamentalrole in the development of new branches ofchemistry such as organic chemistry, bio-chemistry, and nuclear chemistry. The atomictheory and the problem of atomic weightswas at the heart of the development of theperiodic table in the mid-1800s.

As the twentieth-century unfolded, theinternal structure of the atom began tounfold, and as it did, immediate practicaluses were made of this knowledge. Roent-gen’s x-rays were applied to medical diag-nosis almost immediately after theirdiscovery in 1895. While humans havealways pondered the ultimate nature of mat-ter, theoretical knowledge also provided thehope of harnessing the atom’s energy. Thishope was realized in December of 1942when Enrico Fermi (1901–1954) created thefirst controlled chain reaction beneath thefootball stadium at the University ofChicago. This experiment led three yearslater to the use of atomic weapons to endWorld War II. Nuclear fission was initiallyviewed as a new source of energy with vastpotential. The promise of nuclear power hasyet to be realized as nuclear accidents occur-ring at Three Mile Island, Pennsylvania, in1979 and Chernobyl, Ukrainia, in 1986raised serious questions in the public’s mindabout the safety of nuclear power. Currentresearch on fusion may eventually lead to itbeing employed as an energy source. Morewill be said about this topic in the chapteron nuclear chemistry (Chapter 17).

While the hope of obtaining energyfrom atoms has diminished in recent years,

The Atom 47

there are numerous other applications basedon the atomic nature of matter. With simpleatoms and molecules as starting blocks,organic synthesis led to the development ofthe modern chemical industry. Over theyears, a host of materials have been pro-duced that we naturally accept as part ofeveryday life, for example, teflon, polyvinylchloride (pvc), polyester, and nylon. Thisprocess continues as witnessed by the cre-ation of the relatively new discipline ofmaterial science. Advances in medicines andpharmaceuticals are yet another aspect oforganic synthesis that owes its existence toa fundamental knowledge of atomic theory.

The atomic theory is so rooted intoday’s science that we readily accept theexistence of atoms. Even though few of ushave observed them directly, we have faiththat everything around us is made of them.Yet we continue to probe the structure of theatom. It is interesting that we continue toseek greater knowledge about the atom andits smallest constituents in order to under-stand the largest entity we can think of—theuniverse. Just as Democritus’ quest was tobuild a philosophical system to explainnature, we continue to probe the atom inhope of understanding the universe and ourplace in it.

IntroductionAtoms rarely exist as individual units.Atoms combine with each other to producethe familiar substances of everyday life.Chemistry is largely the study of how atomscombine to form all the different forms ofmatter. The reason atoms combine involvesthe subject of chemical bonding, which isexplored in Chapter 7. In this chapter, thegrouping of atoms into different types ofcompounds is examined. In the first half ofthe chapter, chemical nomenclature is dis-cussed. Some of the basic rules for namingcompounds are presented. Atoms combineand are rearranged through chemical reac-tions. The last half of the chapter examinesthe basic process of chemical reactions andclassifies several different types of reactions.

Mixtures, Compounds, andMolecules

An element is a pure substance that can-not be broken down chemically. The smallestunit of an element is an atom. When sub-stances combine and retain their originalidentity, the combination is called a mixture.In a mixture, two or more substances are

brought together and no chemical reactiontakes place. Mixtures may consist of sub-stances in the same or different physicalstates. For example, bronze is an alloy con-sisting of the two solid metals, copper and tin.The most common solutions are mixtures ofa solid in a liquid. Mixtures may be classifiedas homogenous or heterogeneous. Homoge-nous mixtures consist of a uniform distribu-tion of the substances throughout the mixture.Air is a homogenous mixture of nitrogen,oxygen, argon and other trace gases. Hetero-geneous mixtures do not have a uniform con-sistency. A bottle of salad dressing consistingof separate layers of oil and vinegar is anexample of a heterogeneous mixture.

Substances resulting from the chemicalcombination of two or more elements infixed proportions are called compounds. Theelements in compounds are chemicallybonded to each other. Most elements do notexist in their free native state but are foundin combination with other elements aschemical compounds. Compounds mayexist as ionic compounds or molecular com-pounds. In ionic compounds, the constituentelements exist as ions. Ions are atoms orgroups of atoms that carry a charge by virtueof losing or gaining one or more electrons.

5

Atoms Combined

50 Atoms Combined

Ionic compounds result from the combina-tion of a positive ion known as a cation anda negative ion called an anion. Salt is anionic compound in which sodium cationsand chloride anions chemically combined.Molecular compounds contain discretemolecular units. Molecular units or mole-cules are the smallest unit of a molecularcompound. Atoms in a compound are heldtogether by covalent bonds. Bonds dictatehow atoms are held together in a compoundor molecule, but for now, just think of ioniccompounds as compounds composed ofions, and molecular compounds as com-pounds composed of molecules. Sugar,water, and carbon dioxide are examples ofmolecular compounds.

While a compound consists of two ormore elements, a molecule consists of twoor more atoms. The water molecule is com-posed of two atoms of hydrogen and oneatom of oxygen. Not all molecules will contain different elements. The oxygenmolecule is made of two oxygen atoms, andthe ozone molecule consists of three oxygenatoms.

Naming ElementsJust as the modern science of chemistry

developed from the ancient arts andalchemy, so did the language of chemistryprogress from ancient roots. During ancienttimes, humans knew of seven metals andeach of these was associated with the sevenknown celestial bodies and the seven daysof the week. The characteristics of the metaland celestial bodies were thought to berelated, and astrological symbols repre-sented an early form of naming chemicals.Gold was associated with the glowing sun,and a circle, considered to be the most per-fect shape, was used to represent it. Silverwas related to and represented by the moon.Saturn, the most distant planet known at the

time, moved slowly across the sky and wasassociated with the heavy metal lead. Theseven metals and their associated celestialbody and ancient symbols are shown inTable 5.1.

As the ancients and alchemists discov-ered more substances, an increasing numberof symbols were required to represent thesubstances. Because different civilizationsused different symbols for the same sub-stance, confusion resulted and there was nocommon language to transfer chemicalknowledge.

This situation persisted up to the nine-teenth century. Even Lavoisier, who put somuch effort in constructing an unambiguouschemical language, used pictures to repre-sent elements. Lavoisier and his colleaguesused letters enclosed in a circle, short

Table 5.1Metal, Celestial Bodies, and AncientSymbols

Atoms Combined 51

straight lines, and semicircles to representsubstances. Hydrogen, sulfur, carbon, andphosphorus were represented as , , , ,respectively.

Dalton too would not divorce himselffrom the use of pictures to represent ele-ments and compounds. Dalton resorted to aseries of circles containing various patternsand letters to represent substances. A sam-ple of Dalton’s nomenclature is shown inTable 5.2.

In 1814, the eminent Swedish chemistJacob Berzelius (1779–1848) discarded theold sign language of chemical symbols andproposed a new system based on the initialletter of the element. Berzelius, in compil-ing the Swedish Pharmacopia, used the ini-tial letter of its Latin name to symbolize anelement. If two elements had the same firstletter, Berzelius would include a second let-ter that the two elements did not have incommon. Using Berzelius’ system, hydro-gen, oxygen, nitrogen, carbon, phosphorus,and sulfur became H, O, N, C, P, and S,respectively. Berzelius’ system utilized tra-ditional Latin names for existing chemicals,and this explains why some of our modernsymbols seem unrelated to its English name.For example, gold comes from aurum (Au),

sodium from natrium (Na), and potassiumfrom kalium (K).

During the nineteenth and twentiethcenturies, as new elements were identified,the discoverer received the honor of namingthe element. Different trends in assigningnames developed at different times. Ele-ment names were based on mythologicalfigures, celestial bodies, color, chemicalproperties, geographical areas, minerals,derived names, and people. Table 5.3 givesthe derivation of names and symbols for thecommon elements.

Naming CompoundsBerzelius’ method of assigning letters to

represent elements was also applied to com-pounds. Berzelius used superscripts todenote the number of atoms in a compound.Thus, water would be H2O and carbon diox-ide CO2. Later these superscripts werechanged to the current practice of designat-ing the number of atoms using subscripts.The absence of a subscript implies the sub-script one. Using symbols we can write thechemical formula for ammonia:

The rule for naming compoundsdepends on the type of compound. For ioniccompounds consisting of two elements(binary compounds), we start by naming thecation element. After the cation element isnamed, the stem of the anion is used withthe ending “ide” added to the stem.

Anionoxygen oxidenitrogen nitridechlorine chloridesulfur sulfide

Table 5.2Dalton’s Chemical Symbols

N H3

F1 atom of 3 atoms ofNitrogen Hydrogen

F

52 Atoms Combined

Examples of ionic compounds namedusing this system follow:

NaCl sodium chlorideCaF2 calcium fluorideBaO barium oxideNaI sodium iodide

Two or more atoms may combine toform a polyatomic ion. Common poly-atomic ions are listed in Table 5.4.

The names of polyatomic ions may beused directly in compounds that contain

Table 5.3Symbols and Name Derivations for Common Elements

Table 5.4Polyatomic Ions

Atoms Combined 53

them. Hence, NaOH is sodium hydroxide,CaCO3 is calcium bicarbonate, andBa(NO3)2 is barium nitrate.

Naming molecular compounds usesGreek prefixes to indicate the number of

atoms of each element in the compound or molecule. Prefixes are given in Table 5.5.

Prefixes precede each element to indi-cate the number of atoms in the molecularcompound. The stem of the second elementis used with the “ide” suffix. The prefix“mon” is dropped for the initial element,that is, if no prefix is given, it is assumed theprefix is one. Examples of molecular com-pounds are:

CO2 carbon dioxideCO carbon monoxideCCl4 carbon tetrachlorideN2O4 dinitrogen tetroxide

Some compounds are known by theircommon name rather than their systematicnames. Many of these are common house-hold chemicals. Table 5.6 summarizes someof these common chemicals.

Table 5.5Greek Prefixes

Table 5.6Common Chemicals

54 Atoms Combined

Writing Chemical FormulasKnowing the names of the elements and

a few basic rules allows us to name simplecompounds given the chemical formula. Wealso can reverse the process. That is, if weknow the name of the compound, we shouldbe able to write the chemical formula. Theprocess is straightforward for molecularcompounds because prefixes are included inthe names. Hence, the formula for sulfurdioxide is SO2 and carbon monoxide is CO.

Writing the formula for ionic com-pounds requires us to know the charge of thecation and anion making up the ionic com-pound. Information on the charge of com-mon ions can be obtained from the periodictable. More will be said about this in Chap-ter 7, but for now, a few basic rules will helpus write the formulas for simple ionic com-pounds:

1. Elements in the first column (Group 1) ofthe periodic table form ions with a �1charge.

2. Elements in the second column (Group 2)of the periodic table form ions with a �2charge.

3. Generally speaking, elements in Group 13(the column starting with the element B)form ions with a �3 charge.

4. Oxygen and sulfur generally form ionswith a –2 charge.

5. The halogens, Group 17 (the column withF at the top), form ions with a –1 charge.

Using these rules and the fact that acompound carries no net charge allow us towrite the compound’s formula. This is illus-trated for the compound magnesium chlo-ride. Magnesium is a Group 2 element, so itforms Mg2� ions. Chloride is a halogen andforms Cl– ions. For the compound magne-sium chloride to be neutral, two chlorideions must combine with one magnesium ionto give MgCl2. An easy way of determining

the formula on an ionic compound is towrite the elements with their charges overthem and then crisscross the charges tomake subscripts.

Notice that two choride ions, each with a –1charge, are necessary to balance the �2charge of the magnesium ion. If the sub-scripts have a common factor, then thatcommon factor should be divided into bothsubscripts. Therefore, the formula for bar-ium oxide would be BaO.

The same method can be applied when apolyatomic ion is one of the ions. For exam-ple, the formula for sodium sulfate would beNa2SO4.

Chemical ReactionsChemical changes take place through

chemical reactions. Chemical symbols areused to write chemical reactions in a chem-ical equation. For example, when a mixtureof hydrogen gas (H2) and oxygen gas (O2)ignites, the hydrogen and oxygen combineto form water. A chemical equation express-ing the combination of hydrogen and oxy-gen to form water is

Let’s examine the chemical equation abovemore closely. Hydrogen and oxygen are writ-ten as molecules rather than atoms. Hydro-gen, oxygen, and a number of other gases(Cl2, N2, Br2, I2) exist in their natural state as

Atoms Combined 55

diatomic (two atoms) molecules. The reasonfor this diatomic form becomes apparentafter we discuss electron configuration andbonding. The plus sign in the equationmeans “reacts with.” The equation states amolecule of hydrogen reacts with a moleculeof oxygen to yield a water molecule. Hydro-gen and oxygen are called reactants. Reac-tants appear on the left side of a chemicalequation. Water is the product in the reac-tion; products appear on the right side ofchemical equations. The standard practicefor writing a chemical equation is to writethe individual reactants on the left side sep-arated by a plus sign and use an arrow point-ing toward the product(s) on the right side.The circles drawn below the equation areused to represent the atoms. The circles indi-cate this chemical equation is not balanced.One of the basic principles of chemistry thatmust be reflected in the chemical equation isconservation of mass. An accounting of theatoms in the formation of water shows animbalance in the number of oxygen atoms:

The process by which we make a chemicalequation conform to conservation of mass iscalled balancing the equation. A chemicalequation is balanced by placing whole num-ber coefficients in front of reactants and/orproducts to make the number of atoms ofeach element equal on both sides of the equa-tion. In our example, this is done by placinga 2 in front of both hydrogen and water

By balancing the equation, we see that thenumber of hydrogen and oxygen atoms isequal on both sides of the equation:

The balanced equation now reads twomolecules of hydrogen react with a moleculeof oxygen to yield two molecules of water.

Additional information may be includedwhen writing a chemical equation. One com-mon practice is to indicate the physical stateof substances by use of the subscripts (s), (l),(g), (aq) to indicate solid, liquid, gas, oraqueous (dissolved in water), respectively.Hence, the combustion of carbon to formcarbon dioxide would be represented as:

C(s) � O2(g) CO2(g)

A small symbol over the arrow may be usedto indicate conditions necessary for thereaction to take place. For example, a �(Greek delta) over the arrow indicates thatheat is required for the reaction, and an h�indicates ultraviolet radiation is needed forthe reaction to take place.

Types of ReactionsThe reactions of hydrogen and oxygen

to form water and carbon and oxygen toyield carbon dioxide are examples of com-bination reactions. A combination or syn-thesis reaction results when two or moresubstances unite to form a compound. Manyother types of reactions exist. Three othercommon types of reactions are decompo-sition, single replacement, and doublereplacement.

Decomposition is the opposite of com-bination. Decomposition takes place whena compound breaks down into two or moresimpler substances. An electric current canbe used to decompose water into hydrogenand oxygen:

Reactants Products

2 atoms of hydrogen 2 atoms of hydrogen2 atoms of oxygen 1 atom of oxygen

Reactants Products

4 atoms of hydrogen 4 atoms of hydrogen2 atoms of oxygen 2 atoms of oxygen

56 Atoms Combined

Decomposition reactions are used to obtainpure metals from their ore.

In a single replacement reaction, oneelement replaces another element in a com-pound. An example of a single replacementis when sodium is added to water. The reac-tion of sodium and water is represented bythe following chemical equation:

2Na(s) � 2H2O(l) 2NaOH(aq) � H2(g)

Sodium displaces the hydrogen in the watermolecule. The displaced hydrogen is liber-ated from the water resulting in the forma-tion of hydrogen gas and sodium hydroxide.

Double replacement involves two ele-ments exchanging places in a reaction. Acommon type of double replacement reac-tion occurs when a strong acid reacts with astrong base. The reaction of hydrochloricacid and sodium hydroxide would be anexample of a double replacement reaction:

HCl(aq) � NaOH(aq) NaCl(aq) � H2O(l)

In the above reaction, hydrogen and sodiumexchange places. This reaction results in theacid and base neutralizing each other.

From Molecules to MolesOur discussion of balancing chemical

equations has focused on accounting foratoms and molecules in both reactants andproducts. Although a balanced equation tellsus how many atoms of an element it maytake to form so many molecules, it isimpractical to speak in these terms for com-mon applications. For example, if we wantto know how much carbon dioxide is pro-duced when some vinegar is added to bak-

ing soda, a practical unit for the carbondioxide would be liters or milligrams. Indus-trial chemists often measure quantities ingallons and tons. Atoms and molecules arejust too small to be convenient for practicalapplications. The unit used to define theamount of a substance in a chemical reac-tion is the mole (mol). A mole is the quan-tity of a substance that contains as manyparticles (atoms, molecules, ions) as thereare atoms in exactly 12 grams of carbon-12.

The use of carbon-12 as a standard wasadopted in 1961 by international agreement.Originally, Dalton had assigned an atomicmass of 1 for hydrogen, and the atomic massesof all other elements were based on this value.Using a value of 1 for hydrogen gave anatomic mass of 15.9 for oxygen. Because oxy-gen was typically used in determining howelements combined, it was decided it mademore sense to base atomic masses using anoxygen standard. Oxygen was assigned avalue of 16; this designation meant hydrogenhad an atomic mass of slightly greater than 1.A complication developed, though, becauseoxygen consists of three isotopes. Differentsamples of oxygen containing slightly differ-ent ratios of the isotopes would have differentmasses, which presented problems becausephysicists were using oxygen samples todetermine atomic masses. In 1961, physicistsand chemists agreed to use the isotope carbon-12 as the standard for atomic mass measure-ments. Carbon-12 with six protons and sixneutrons was assigned a value of 12.0000 fordetermining atomic masses.

What exactly does it mean to have asmany atoms as there are in 12 grams of C-12?The number is known as Avogadro’s num-ber, and its accepted value is 6.022 � 1023.Therefore, from the definition of a mole wecan say there are 6.02 � 1023 atoms in a moleof C-12. The term Avogadro’s number wascoined by Jean Baptiste Perrin (1870–1942)who was the first to experimentally determine

currentelectric

r2H O 2H � O2 (l) 2(g) 2(g)

Atoms Combined 57

an accurate value of this constant in 1908.One mole of any substance contains approxi-mately 6.02 � 1023 particles of the substance.One mole of hydrogen atoms is equal to 6.02� 1023 hydrogen atoms, one mole of hydro-gen molecules equals 6.02 � 1023 moleculesof hydrogen, and one mole of toothpicks is6.02 � 1023 toothpicks.

The fact that both the mole and the massof an element are based on carbon-12enables us to relate mole and mass. A molarmass is defined as the mass in grams of onemole of a substance, and it can be obtaineddirectly from an element’s atomic mass. Wecan use the elements hydrogen and nitrogento illustrate this concept. Periodic tableentries for both elements are shown below.The whole number above the element is theatomic number and gives the number of pro-tons in the nucleus. The number below theelement’s symbol is the molar mass (as wellas the atomic mass):

The entries indicate one mole of hydrogenatom equals 1.00794 gram and one mole of nitrogen atoms equal 14.0067 grams.Because each hydrogen molecule containstwo hydrogen atoms, a mole of hydrogenmolecules contains two moles of hydro-gen atoms. Therefore, one mole of hy-drogen molecules, H2, would equal 2 �1.00794 � 2.01588 grams. In a similar fash-ion, the molar mass of ammonia, NH3 canbe calculated by totaling the molar massesof constituent elements:

NH3 N 1 � 14.0067 � 14.0067H3 3 � 1.00794 � �3.0238

17.0305

Therefore, one mole of ammonia contains6.022 � 1023 NH3 molecules and has a massof 17.0305 grams; conversely, 17.0305grams of ammonia equals one mole ofammonia.

Coefficients in a balanced chemicalequation are typically taken to representmoles. For the reaction of hydrogen andnitrogen to form ammonia, the equation

N2(g) � 3H2(g) 2NH3(g)

can be read as one mole of nitrogenmolecules reacts with three moles of hy-drogen molecules to yield two moles ofammonia molecules. Knowing the moles of reactants and product(s) in the balancedequation allows us to express mass relation-ships. The ammonia reaction expressed interms of mass tells us that approximately 28grams of nitrogen reacts with 6 grams ofhydrogen to yield 34 grams of ammonia.

Having a balanced chemical equationand knowing the relationship between massand moles allows us to predict how muchreactant is necessary to yield a certainamount of product. This knowledge hasimportant applications in industrial chem-istry, environmental chemistry, nutrition,and in any situation where reactions takeplace. The balanced equation is a recipe fora chemical reaction. Just as it is necessaryto know the amount of eggs, flour, sugar,and salt to bake a cake, we need to know theamount of ingredients that go into a chemi-cal reaction. The balanced chemical equa-tion gives the quantities of differentreactants that are required to produce a spe-cific amount of product.

We can illustrate this with a practicalproblem having environmental implications.The greenhouse effect refers to the heatingof the Earth due to the trapping of infraredradiation by certain gases. A principalgreenhouse gas is carbon dioxide, CO2.

58 Atoms Combined

International efforts in recent years haveattempted to limit the amount of carbondioxide produced by individual countries.Because carbon dioxide is a principal prod-uct when fossil fuel is burned, any limitationon carbon dioxide has far-reaching effectson a nation’s economy. Automobiles are achief producer of CO2 in the United States.How much CO2 is produced by automo-biles? Let’s calculate the amount of CO2

produced when a gallon of gas is burned.We can simplify the problem a bit by assum-ing the gasoline is octane and it reacts com-pletely with oxygen to produce carbondioxide and water according to the equation:

One gallon of octane has a mass ofapproximately 2,600 grams. The molar massof octane is 114 grams, so 2,600 grams ofoctane is almost 23 moles of octane. Noticethat the coefficients of CO2 and C8H18 in theequation are 16 and 2, respectively. There-fore, the balanced equation tells us that forevery mole of octane burned, eight moles ofcarbon dioxide are produced. This means182 moles of CO2 result (23 � 8). Becausecarbon dioxide’s molar mass is 44 grams,the amount of CO2 produced is about 8,000grams. It is interesting to note that the ratioof carbon dioxide produced to gasoline(octane) burned is more than three to one.Translated into English units, this meansthat for every gallon (a gallon is about sixpounds) of gas burned, 18 pounds of carbondioxide are produced. Given the number ofcars in operation in this country, we see thata tremendous amount of carbon dioxide isemitted to the atmosphere every day. Indus-tries and municipalities regularly use similarcalculations to predict the amount of pollu-tion they are emitting to the atmosphere.Such calculations help them conform to theClean Air Act.

Balanced chemical reactions are criticalto the chemical industry. Plant managersmust know the amount of reactants neces-sary to yield a product to keep productionlines moving. If reactants or products are notpure, this must be taken into account. Alter-native reactions can be examined to find themost cost-effective process. For example, afertilizer may be produced using one pro-cess with sulfuric acid as a reactant oranother process that starts with ammoniumcarbonate. Working through the balancedequations for each process would helpdecide which process to adopt.

A final example of an application of bal-anced chemical equation involves the use ofone form of the breath analyzer used bypolice to determine if someone was drivingunder the influence of alcohol. The chemi-cal equation guiding the reaction whensomeone breaths into the instrument is

The equation shows that alcohol in the formof ethanol reacts with an orange solution ofpotassium dichromate to produce a greensolution of chromic sulfate. The source ofalcohol is the suspect’s breath. More alco-hol produces a greater color change. Thebreath analyzer measures this color changeand coverts this measurement into anamount of alcohol in the blood.

In many instances, the ratio of reactantsavailable is different than that given by thebalanced chemical equation. When thishappens, the reactant in the smallest rela-tive abundance is said to be limiting, whilethe other reactant is referred to as theexcess reactant. Again, using the ammoniareaction, we see the ratio of hydrogen tonitrogen is 3 to 1. If three moles of both

2C H � 25O r 16CO � 18H O8 18 2 2 2

octane � oxygen r carbon dioxide � water

3C H OH � 2K Cr O � 8H SO 3CH COOH2 5 2 2 7 2 4 3

� 2Cr (SO ) � 2K SO � 11H O2 4 3 2 4 2

ethanol � potassium � sulfuric acetic acid � chromic � potassium � waterdichromate acid sulfate sulfate

orange green

r

r

Atoms Combined 59

hydrogen and nitrogen are brought together,the amount of hydrogen will limit the reac-tion. Nitrogen will be in excess; two extramoles of nitrogen would be left over afterthe reaction is complete. An analogy mayhelp to illustrate the idea of excess and lim-iting reactants. Suppose you wanted tomake bicycles. For each bicycle we needtwo tires and one frame. Although otherparts are required, we can simplify the anal-ogy by focusing on tires and frames. If wehad twenty tires and twenty frames, wecould build ten bikes. The tires would limitthe number of bikes, and the frames wouldrepresent the excess reactant. Even if wehad one hundred frames, we could buildonly ten bikes with twenty tires. A chemi-cal reaction is analogous to building bikes.One substance can be totally consumedwhile another remains when the reactionceases.

One last point is that just because achemical equation indicates a chemicalreaction takes place, does not mean it will.

Many conditions are required for a chemi-cal reaction to proceed. Conditions such asheat, light, and pressure must be just rightfor a reaction to take place. Furthermore, thereaction may proceed very slowly. Somereactions occur in a fraction of a second,while others occur very slowly. Consider thedifference in the reaction times of gasolineigniting in a car’s cylinder versus the oxida-tion of iron to form rust. The area of chem-istry that deals with how fast reactions occuris known as kinetics (Chapter 12). Finally,not all reactions go to completion. Theamount of product produced based on thechemical equation is known as the theoreti-cal yield. The amount actually obtainedexpressed as a percent of the theoretical isthe actual yield. In summary, it’s best tothink of a chemical equation as an ideal rep-resentation of a reaction. The equation pro-vides a general picture of the reaction andenables us to do theoretical calculations, butin reality reactions deviate in many waysfrom that predicted by the equation.

IntroductionIf there were a flag that represented the sci-ence of chemistry, it would be the periodictable. The periodic table is a concise organi-zational chart of the elements. The periodictable not only summarizes important factsabout the elements, but it also incorporatesa theoretical framework for understandingthe relationships between elements. Themodern periodic table attests to human’ssearch for order and patterns in nature. Assuch, the periodic table is a dynamicblueprint for the basic building blocks of ouruniverse. This chapter examines the devel-opment of the modern periodic table andpresents information on how the modernperiodic table is organized.

Discovery of the ElementsThe organization of the elements into

the first modern periodic table in 1869depended on two important developments.One was that enough elements had to beidentified to organize the elements into ameaningful pattern. Nine elements wereknown in ancient times. These were gold,silver, iron, mercury, tin, copper, lead, car-

bon, and sulfur. Bismuth, zinc, antimony,and arsenic were isolated in the alchemicalperiod and Middle Ages. By the start of thenineteenth century, approximately thirty ele-ments had been identified, and this numberdoubled over the next sixty years. Currently,116 elements have been identified, but sci-entists at research labs throughout the worldare attempting to create more using highspeed particle accelerators. Figure 6.1 givesthe approximate dates of the discovery ofelements during different time periods.

As the number of elements increased, sodid attempts to organize them into mean-ingful relationships. Johann Döbereiner(1780–1849) discovered in 1829 that certainelements had atomic masses and propertiesthat fell approximately mid-way between themasses and properties of two other ele-ments. Döbereiner termed a set of three ele-ments a triad. Thus, chlorine, bromine, andiodine form a triad; Döbereiner proposedseveral other triads (lithium-sodium-potassium, calcium-strontium-barium). Dö-bereiner recognized that there was some sortof relationship between elements, but manyelements did not fit in any triad group, andeven those triads proposed displayed numer-ous inconsistencies.

6

Elements and the Periodic Table

62 Elements and the Periodic Table

John Newlands (1837–1898) attemptedto organize the elements by listing them inorder of increasing atomic mass. Newlandsprepared a table with elements grouped incolumns of seven elements. Several rows inNewlands’ table contained related elementsand also included Döbereiner’s triads.Because there was some evidence of char-acteristics repeating with every eighth ele-ment, Newlands proclaimed a law ofoctaves for the elements. Newlands’ tablelisted the elements in a continuous mannerand did not leave any blank spaces. Hisadherence to a rigid organizational schemeresulted in a number of inconsistencies.Other chemists ridiculed Newlands when hepresented his table at conferences and meet-ings between 1863 and 1865. In fact, therewas so much skepticism regarding his workthat he was unable to get it published. Yearslater Newlands was vindicated when othersrecognized his insight of recognizing peri-odic patterns among the elements.

The second development that led to themodern periodic table was the acceptance ofspecific atomic masses for the elements.While today we readily accept the massesgiven in a periodic table, there was much

disagreement regarding masses in the mid-nineteenth century. A major problem thatwas not recognized until that time is thatsome elements exist in their natural state asatoms, while others exist as molecules. Theelement hydrogen (H) has an atomic massof approximately 1, but hydrogen exists asa diatomic molecule (H2) with a mass ofapproximately 2. The existence of moleculesand Avogadro’s law was the key for deter-mining universally accepted atomic masses.Although Avogadro put forth his ideas in1811, the concept of molecules as a basicunit of matter was not widely accepted inthe mid-nineteenth century. The situationchanged when Stanislao Cannizzaro(1826–1910), a disciple of Avogadro’s ideas,resurrected Avogadro’s work at the FirstInternational Chemical Congress convenedin Karlsruhe, Germany, in 1860. Cannizzarovigorously defended the existence ofmolecules using Avogadro’s law as a basisfor determining the correct atomic massesof the elements. Cannizzaro circulated apaper he had written in 1858 outlining hisideas on the matter. This paper helped per-suade several prominent chemists on theutility of incorporating Avogadro’s law into

Figure 6.1Discovery Dates of the Elements

Elements and the Periodic Table 63

their thinking. Cannizzaro’s crusade and theKarlsruhe Conference ultimately resulted inuniversally accepted values for atomicmasses.

Two chemists in attendance at the Karls-ruhe were Julius Lothar Meyer (1830–1895)and Dmitri Mendeleev (1834–1907). Thesetwo independently developed the periodiclaw and constructed their own versions ofthe periodic table. Meyer based his table pri-marily on the physical properties of the ele-ments. Meyer plotted atomic volume againstthe atomic mass and noticed the periodicityin volumes of the elements. Other physicalproperties also showed periodic trends. Fig-ure 6.2 shows how the melting point of thefirst fifty-five elements rises and falls in aroughly periodic fashion as atomic numberincreases. Based on his analysis, Meyerpublished his periodic table in 1870.

Despite Meyer’s efforts, Mendeleevreceives the majority of credit for the mod-ern periodic table. Mendeleev’s tableresulted from years of work examining theproperties of the elements. Mendeleev’smethod involved writing information aboutthe elements on individual cards and thentrying to organize the cards in a logicalorder. His genius resulted from modifica-tions he introduced into his table that othershad not included. He produced his table in

1869 as part of a text he had written forchemistry. Mendeleev did not alwaysarrange the elements in strict order ofatomic masses, and he claimed that severalof the accepted atomic masses were wrong.More important, he left blank spaces in histable where he claimed undiscovered ele-ments existed. Two of the empty spaces fellright after aluminum and silicon, respec-tively. Mendeleev named these undiscoveredelements eka-aluminum and eka-silicon(“eka” is the Sanskrit word for first, so eka-aluminum was first after aluminum). Usinghis knowledge about the periodic propertiesof the elements, he predicted the chemicaland physical properties of eka-aluminumand eka-silicon. In 1875, the element gal-lium (named after ancient France, Gaul) was discovered, which turned out to beMendeleev’s eka-aluminum. The propertiesof gallium were almost identical to thoseMendeleev had predicted as seen in Table6.1. Germanium (named after Germany)discovered in 1886 turned out to beMendeleev’s eka-silicon. Scandium (namedfor Scandanavia) discovered in 1879 wasanother element Mendeleev had predictedwith his table.

Mendeleev’s periodic table (Figure 6.3)was used with little modification until wellinto the twentieth century. The top of each

Figure 6.2Periodic Trend of Melting Point


group gives the general formula for how thegroup’s elements combine with hydrogenand/or oxygen. Blank spaces appear atatomic masses of 44, 68, 72, and 100 forundiscovered elements. At the turn of thecentury, the discovery of the noble gases ledto the addition of another column. Since thetime of Mendeleev’s original table, manydifferent versions of the periodic table haveappeared. These include three-dimensional,circular, cylindrical, and triangular shapes.Yet, Mendeleev’s basic arrangement hasstood the test of time, and the periodic tableis the universally recognized icon of chem-ical knowledge.

The Modern Periodic TableElements in the modern periodic table

are arranged sequentially by atomic numberin rows and columns. Mendeleev and hiscontemporaries arranged elements accord-ing to atomic mass. In 1913, Henry Mose-ley’s (1887–1915) studies on the x-raydiffraction patterns for metals showed arelationship between the spectral lines andthe atom’s nuclear charge. Moseley’s workestablished the concept of atomic number,the number of protons in the nucleus, as thekey for determining an element’s position inthe periodic table. Rows in the periodic table

are referred to as periods, and columns arecalled groups. The periods run left to right,and the groups from top to bottom. A groupmay also comprise a chemical family. Thefirst and second groups are the alkali met-als and alkaline earth metals, respectively.The group starting with fluorine (F) is thehalogen family (halogen means salt form-ing). The last group on the far right of thetable is the noble gas family.

Groups are numbered across the top ofthe table. Unfortunately, Americans andEuropeans have different group numberingsystems that at times cause confusion. TheAmerican system uses Roman numerals fol-lowed by either the letter A or B. The Agroups are known as main group or repre-sentative elements. Those groups with theletter B comprise the transition elements.Main groups fall at both ends of the periodictable with transition elements in the center.The European system also uses Romannumerals and the letters A and B, but the Ais used for groups on the left half of the tableand B for groups on the right half. In recentyears, scientists have suggested eliminatingthe letters and using the Arabic numerals1–18 to number each group starting on theleft. We use the latter method to identifygroups.

The Main Group ElementsTo understand how elements are orga-

nized into different groups, we need to con-sider the electron configuration of elements.Specifically, it is the electron configurationof the valence electrons that characterizesgroups. Remember, the valence electronsare those electrons in the outer shell. We canuse Group 1 to illustrate how the valenceelectron configuration defines the group.The electron configurations of Group 1 ele-ments are listed on page 66.

Table 6.1Comparison of Mendeleev's Prediction andActual Properties for Eka-Aluminum(Gallium)

Mendeleev’sPrediction

ActualProperties

Atomic Mass 68 69.7Melting Point low 29.8�CDensity 5.9 g/cm3 5.9 g/cm3

Oxide X2O3 Ga2O3

Figure 6.3Mendeleev’s 1869 Periodic Table (Top) and 1870 Table (bottom). Image from Edgar Fahs SmithCollection, University of Pennsylvania Library.


H 1s1

Li 1s2 2s1

Na 1s2 2s2 2p6 3s1

K 1s2 2s2 2p6 3s2 3p6 3d10 4s1

Rb 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 4f14 5s1

Cs 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 4f14 5s2

5p6 5d10 6s1

Fr 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 4f14 5s2

5p6 5d10 6s2 6p6 7s1

We see from the electron configurations thatall Group 1 elements have a valence config-uration of ns1, where n is the period of theelement. Electron configurations for ele-ments in Groups 2 and 13–18 show a regu-lar pattern similar to that demonstratedabove. Each group would have its own dis-tinctive valence electron configuration:

Group Valence Electron Configuration

1 ns1

2 ns2

13 ns2np1

14 ns2np2

15 ns2np3

16 ns2np4

17 ns2np5

18 ns2np6

Groups 1 and 2 together are often referredto as the s-block, and Groups 13–18 are thep-block elements. These names come fromthe valence orbital being filled in eachblock. In the s-block, the s orbital is filled,while in the p-block, the p orbital is pro-gressively filled across a period.

The valence electron configurationexplains why groups define chemical fami-lies that exhibit similar chemical properties.Because it is the valence electrons thatdetermine an element’s chemical behavior,and groups have similar valence electronconfigurations, then elements in the samegroup display similar chemical behavior.The electron configuration also explains the

periodic law. We can demonstrate this byexamining how one property, the atomicradius, changes moving across a period.Lithium is the first element in the 2nd periodwith an atomic number of 3. It has an atomicradius of 152 picometers (pm). As we moveacross the 2nd period, the atomic numberincreases in increments of one. Each timethe atomic number increases one more pro-ton is added to the nucleus and one addi-tional electron is added to the 2nd electronshell. Because each electron goes into ashell where the principal quantum numberis 2, and the principal quantum numbergives the average distance from the nucleus,we would expect the radii of atoms acrossthe period to stay constant. In reality, mov-ing across the period there is a tendency forthe radii to decrease. This happens becauseas protons and electrons are added, the pos-itive nucleus exhibits a relatively strongerpull on the cloud of electrons in the 2porbital. This trend continues across theperiod until we reach sodium. At this point,an additional electron must go into the 3rdshell, and the radius displays a “jump” fromneon. After this abrupt increase in sodium’sradius occurs, a trend similar to that forperiod 2 repeats itself in period 3. This trendis demonstrated in Figure 6.4, which is agraph of atomic radii versus atomic numberfor periods 2 and 3.

The periodic trend of a decrease inatomic radii across a period is readily seenin the Figure 6.4. Other properties related toatomic radii show a similar pattern. Know-ing that the elements exhibit a general peri-odic trend allows us to predict unknownproperties for elements and aided in the dis-covery of unknown elements. The periodicnature of the elements supported the devel-opment of the quantum theory. The ele-ments show a periodic pattern in both theirproperties and electron configurations. Theperiodic trend in properties of the elements


should be interpreted as a general patternthat elements in the periodic table follow.There are many examples of deviationsfrom the periodic law. It should not beviewed as a strict rule, but as pattern thatpresents general trends among the elements.

d and f Block ElementsThe elements in Groups 3 to 12 are

called the transition elements or transitionmetals. Transition elements are characterizedby a filling of the d orbitals from left to rightacross a period. Collectively these elementsform the d-block of the periodic table. Thed-block groups do not demonstrate the con-sistent pattern displayed by main group ele-ments. In general, though, the pattern offilling the d orbital gives transition elementsa nd1–10 configuration. For example Group 3elements have nd1 electrons, and Group 12nd10 electrons, where n represents the period.

The two periods at the bottom of theperiodic table represent two special groupsthat fall between the s and d blocks. Theperiod beginning with lanthanum (La) iscalled the lanthanides. The next period begin-ning with actinium (Ac) is the actinides. Ele-ments in these two periods form the f-block.

Moving across each period the f orbital isprogressively filled; the 4f orbital is filled forthe lanthanides, and the 5f orbital is filled for the actinides. These elements are some-times referred to as the rare earths, becauseit was originally difficult to separate andidentify these elements. Rare earths are actu-ally not scarce, but the term “rare earths” isstill used for the lanthanides and actinides. Amore accurate modern term for these twoperiods are the inner transition elements.

Metals, Metalloids, and Nonmetals

Figure 6.6 summarizes different blocks,families, and areas of the periodic table.Most elements can be classified as metals.Metals are solid at room temperature, aregood conductors of heat and electricity, andform positive ions. Moving across the tablefrom left to right elements lose their metal-lic characteristics. The metalloids, alsoknown as the semi-metals, have propertiesintermediate between metals and nonmetals.Because they display characteristics of bothconductors and nonconductors, elementssuch as silicon and germanium find wideuse in the semi-conductor industry. Non-metals are found on the far right of the peri-odic table. Nonmetals are poor conductorsand are gases at room temperature.

Periodic ProblemsDon’t be surprised if you see a periodic

table that looks different than the one pre-sented in this book. As already mentioned,slightly different versions exist depending onthe conventions used to prepare the table.Some tables start the lathanides and actinidesin the f-block with cesium and thorium.Another problem is the placement of hydro-gen. Hydrogen is positioned above lithiumin Figure 6.5, which indicates it can act like

Figure 6.4Change in Atomic Radii for Periods 2 and 3


an alkali metal, but hydrogen should not beconsidered an alkali metal. It often acts as anonmetal like a halogen because, like otherhalogens, it needs one more electron toachieve the electron configuration of a noblegas. Hydrogen is a gaseous nonmetal that attime behaves like a halogen. Some tables listhydrogen separately to indicate that it is notassociated with any family.

The problem concerning the use of let-ters to identify groups has already been men-tioned. A similar but more serious probleminvolves the naming of newly discovered ele-

ments. The elements beyond uranium, withatomic numbers greater than 92, are not nat-ural. These heavy elements are synthesizedin particle accelerators. New elements arecreated by bombarding an existing heavy ele-ment with neutrons or the nucleus of anotherelement such as hydrogen, carbon, or oxy-gen. In the latter case, fusion of the nuclei ofthe heavy element and the bombarding ele-ment creates a new element. The newly cre-ated heavy element is unstable and may lastfor only a fraction of a second before itdecays. The fact that only a few atoms of a

Figure 6.5Blocks, Families, and Areas of the Periodic Table


new element are created for only a fractionof a second makes detection difficult. It iscommon for separate research labs to claimpriority of the discovery of new elements.Because elements are named by their dis-coverers, it is important to establish priorityin order to assign a name to the new element.During the past twenty years, American,Russian, and German scientists have had sev-eral disputes over discovery claims. This con-fusion has led to some elements not having auniversally accepted name. The probleminvolves the naming of the transferium ele-ments, that is, those elements whose atomicnumber is greater than 100.

IUPAC (International Union of Pure andApplied Chemistry) has attempted to settlethe controversies and determine priority withmixed success. A temporary system pro-posed to name newly discovered elements isbased on using a symbol for each of the dig-its zero through nine:

0 nil 5 pent1 un 6 hex2 bi 7 sept3 tri 8 oct4 quad 9 enn

Using this system, element 104 is namedunnilquadium (un � nil � quad � ium). Theium is added to indicate the element is ametal. Unnilquadium’s symbol is Unq. Com-bining terms to give an element’s name isjust a temporary solution until an acceptedname is determined. Stringing the termstogether, while convenient, results in ratherawkward names. One problem IUPAC facedwhen trying to decide an element’s name wastheir insistence that an element could not benamed for a living person. In the mid-1990s,Americans called element 106 seaborgium,named after Glenn Seaborg (1912–1997)who at the time was still alive. In 1997, afterSeaborg’s death, the names of elements 101to 109 were established:

101 medelevium 106 seaborgium

102 nobelium 107 bohrium

103 lawrencium 108 hassium

104 rutherfordium 109 meitnerium

105 dubnium

While this solved the dilemma for elementsup to number 109, new elements continue tobe discovered. In 1994, elements 110 and111 were identified, and in 1996, element112 was discovered. In recent years, ele-ments 113, 114, 116, and 118 have been dis-covered, and very soon all elements between104 and 118 should be identified. Elements104–118 will fill another row below theactinides, and are called the transactinides.Beyond element 118 lie the superactinideswaiting for more powerful accelerators andbetter detectors to be created so they can beidentified.

SummaryFrom ancient times, humans have pon-

dered what the universe is made of. Earlyphilosophers proposed fire, earth, water, andair either individually or in combination as thebuilding blocks of nature. Lavoisier defined an element operationally as a substance thatcannot be broken down chemically. Using this definition, the number of elements hasincreased from around 30 in Lavoisier’s timeto over 115 today. The initial search for ele-ments involved classical methods such asreplacement reactions, electrochemical sepa-ration, and chemical analysis. New methodssuch as spectroscopy greatly advanced the dis-covery of new elements during the twentiethcentury. The last half century has been markedby the synthesis of elements by humans.

Though the quest to produce more ele-ments continues, amazingly our observableuniverse is composed of a relatively fewcommon elements. Hydrogen and heliummakes up over 90% of the universe. Only six


elements make up over 96% of the Earth asshown in Table 6.2.

Over 99% of the human body is com-posed of only seven elements (Table 6.3),yet trace amounts of a number of elementsare essential for vital functions, for exam-ple, iron in hemoglobin or zinc in enzymes.

Although our world is dominated by afew common chemicals, all the natural chem-ical elements are essential to both our sur-vival and our society. We have come a longway in learning how to combine less than onehundred elements into millions of differentsubstances that enhance our lives. The list ofproducts is endless, but consider what lifewould be like without using the elements to produce fertilizers, medicines, clothing,

building materials, paper, and electronics. Weoften lose sight of the fact that elements arethe building blocks of nature, and all sub-stances are essentially elements packaged indifferent bundles. Just consider that we breathoxygen, drink hydrogen and oxygen, eat car-bon (and host of other elements), calculatewith silicon and oxygen computer chips, lightour homes with tungsten, advertise withneon, drink out of aluminum, store foods intin, disinfect with iodine and chlorine, andthe list goes on. We live in a sea of elements,and we ourselves are part of that sea.

Table 6.2Elemental Composition of Earth by Mass

Table 6.3Elemental Composition of Human Body by Mass

Introduction and HistoryOnly a small portion of matter consists ofelements in uncombined atomic states.Hydrogen, which comprises over 90% of theuniverse, exists as molecular hydrogen, H2.Likewise, the nitrogen and oxygen in the airwe breath exist as the diatomic molecules N2

and O2. Practically all substances consist ofaggregates of atoms in the form of com-pounds and molecules. The fact that matterexists as combinations of atoms leads natu-rally to several basic questions:

• What forces hold atoms together?

• Why do certain elements combine toform compounds?

• How are atoms connected to formcompounds?

The concepts of chemical bonding providethe answers to these questions and are fun-damental to the science of chemistry.

The atomists led by Leucippus andDemocritus explained chemical combina-tion by proposing that atoms consisted ofvarious shapes, which contained projec-tions. Different-shaped atoms interlockedlike pieces of a jigsaw puzzle or atoms with

hook-like projections latched onto otheratoms. By proposing atoms of various sizesand shapes, the atomist Greeks explainedthe existence of the three states of matter.Solids were composed of atoms with jaggededges, while liquid and gases were made ofparticles with smooth, rounded edges. Aris-totle rejected the idea of atoms and consid-ered matter to be continuous in nature.Adopting this viewpoint eliminated the needto address chemical combination.

Traditional theories on chemical combi-nation were based on projecting observablephenomenon at the macroscopic level to themicroscopic level. Hence, the mechanicalexplanations using size and shape persistedwell into the 1600s. Alchemical explanationsprojected human qualities on matter toaccount for specific reactions and combina-tions. Substances that combined had a natu-ral attraction or “love” for one another.Conversely, other substances “hated” eachother and did not combine.

Isaac Newton (1642–1727) was one ofthe first to speculate on the true nature ofchemical attraction. Newton believed anattractive force similar to that of gravity,magnetism, and electricity also applied tomatter at the molecular level. Other scientists

7

Chemical Bonding

72 Chemical Bonding

based their ideas of chemical attraction onNewton’s law of universal gravitation.Comte de Buffon (1701–1788) proposedchemical combination resulted from thegravitational attraction between particles ofvarious shapes. Buffon believed the attrac-tive force depended on the inverse square ofthe distance between particles, while the spe-cific shapes of particles explained while spe-cific substances combined.

While several natural philosopherssought a theoretical explanation for chemi-cal reactions and combinations, many indi-viduals were content to merely catalog thecombining properties of substances. Thestudy of how substances combined became

known as chemical affinity. Tables ofchemical affinity were developed to sum-marize different reactions. Etienne FrancoisGeoffroy (1672–1731) presented his ideason chemical affinity in 1718 (Figure 7.1).Others such as Tobern Bergman (1735–1784) and Berthollet advanced Geoffroy’sideas. Bergmann realized that chemicalaffinity depended on conditions such astemperature and quantities of reactants.

Experiments on electricity at the start ofthe nineteenth century shed light on thenature of chemical affinity. Humphrey Davy(1778–1829) used the voltaic cell inventedby Alessandro Volta (1745–1827) to isolatepure metals from their oxides. Davy used

Figure 7.1Diagram of Geoffroy’s Table of Affinities

Source: Leicester and Klickstein, 1952.

Chemical Bonding 73

electrolysis to discover a number of metalsincluding potassium, sodium, calcium, bar-ium, and magnesium. Davy’s work providedevidence that chemical affinity was some-how related to electricity. Michael Faraday(1791–1867), Davy’s assistant and succes-sor, expanded upon Davy’s studies andbelieved electrical attraction and chemicalaffinity not only were related, but also wereactually a manifestation of the same force.Berzelius proposed a model in which a com-pound was held together because it con-tained a positive part and negative part.Berzelius’ model explained affinity for sim-ple inorganic compounds, but it fell shortwhen applied to more complex compounds,especially organic compounds. Berzelius’polar model consisting of an attractionbetween positive and negative componentsof a compound introduced the concept ofthe ionic bond.

To explain the structure and composi-tion of organic molecules, Berzelius pro-posed radicals as basic building blocks.Berzelius believed radicals had oppositeelectrical charges and were held togetherby electrical attraction. According toBerzelius, organic radicals consisted ofindestructible combinations of positivehydrogen and negative carbon. AugusteLaurent (1808–1853) disputed Berzelius’ideas on organic structure and proposedthat organic molecules had a nucleus towhich radicals, single atoms, or groups ofatoms were attached. The nucleus of theorganic molecule dictated its nature. Lau-rent carried out reactions in which chlorinewas substituted for hydrogen in organiccompounds. Because both carbon andchlorine were considered negative, Lau-rent’s work refuted Berzelius’ radical theory. Laurent believed a process of sub-stitution was a better explanation for howorganic molecules were built.

Berzelius’ idea of charged radicals lostfavor to Laurent’s substitution model after

Berzelius died in the mid-1800s. Using Lau-rent’s model as a starting point, it becameapparent that a central atom serving as anucleus would only accept a certain numberof atoms or groups when it combined withother atoms. For example, oxygen wouldonly accept two atoms or radicals, nitrogenwould combine with three, and sodiumwould only combine with one. EdwardFrankland’s (1825–1899) work on organo-metallic compounds led him to propose thetheory of combining power. Frankland usedthe term “atomicity” in a paper he deliveredin 1852 to indicate the tendency of an atomto combine with a specific number of otheratoms. Frankland’s ideas on combinationwere advanced by Friedrich Kekulé(1829–1896) in the 1860s and led directlyto the concept of valence. The term“valence” (Latin for power) started to beused about this time.

By the start of the twentieth century,the stage had been set for advancing theidea of chemical affinity. Thomson’s dis-covery of the electron provided a theoret-ical foundation upon which new ideas ofaffinity could be based. Chemical affinityand valence had been based on empiricalevidence; the discovery of the electronand a clearer picture of atomic structureprovided the framework for new ideas onchemical bonding. The ionic bond involv-ing the transfer of electrons from oneatom to another was firmly established atthe start of the twentieth century, but ithad numerous shortcomings that made itdifficult to explain certain observations.Specifically, halogen molecules, such asCl2 and I2 could not be explained using apolar model based of ionic bonds. Theproblem was how could two halogens,which were considered to be negativelycharged, be attracted to each other. A sim-ple polar explanation also fell short whentrying to explain bonding in organicmolecules.

74 Chemical Bonding

Richard Abegg (1869–1910) noted thatthe noble gases were extremely stable andsuggested other atoms may gain or lose elec-trons to achieve a stable electron configura-tion similar to the noble gases. One of thepioneers in developing new concepts on affin-ity was the American chemist Gilbert NewtonLewis (1875–1946). Lewis expanded Abegg’sidea into the octet rule. Lewis had proposedas early as 1902 that electrons were arrangedin an atom around a nucleus in a cubic con-figuration as represented in Figure 7.2.

In 1913, Lewis first hinted at a nonpolarbond in which electrons were not transferredbetween atoms, but shared by atoms. Itshould be noted that at about the same timeas Lewis was developing his theories simi-lar ideas were proposed by the GermanWalther Kossel (1888–1956). Lewis pub-lished his ideas on bonding in a 1916 paper.In this paper, Lewis proposed that electronsat the corners of cubic atomic structureswere shared. According to Lewis, two cubessharing an edge represented the affinity ofhalogens. This shared edge would comprisethe single bond holding the atoms together.In this arrangement, each atom would besurrounded by eight electrons (the octetrule) and obtain the stable electron configu-ration of a noble gas (Figure 7.3).

A double bond formed when atomsshared a common face with four electronsbeing shared (Figure 7.4).

Lewis’s model established the idea ofthe nonpolar covalent bond, although hisidea of the cubic arrangement of electronshad several major flaws, for example, howto represent a triple bond in which six elec-

trons must be shared. Irving Langmuir(1875–1946) further developed Lewis’s ideaon the nonpolar bond from 1919 to 1921.The nonpolar model became known as the Lewis-Langmuir model. The Lewis-Langmuir model established the covalentbond to complement the ionic bond toexplain chemical affinity.

Dot Formulas, the OctetRule, and Ionic Bonds

Lewis and Langmuir’s ideas were beingdeveloped as the modern view of the atomwas unfolding. Knowledge of atomic struc-ture enables us to develop a more accuratepicture of chemical bonding. A chemicalbond is simply a force that holds atomstogether. Chemical bonding involves thevalence electrons, those electrons in theouter shells of atoms. One method of repre-senting an atom and its valence electrons isby using Lewis electron dot formulas. ALewis electron dot formula consists of anelement’s symbol and one dot for eachvalence electron. The representative ele-ments and noble gases contain between one

Figure 7.2Lewis’s Representation of Valence Electronsin Cubic Arrangements

Figure 7.3Lewis’s Representation of Single Bond

Figure 7.4Lewis’s Representation of Double Bond

Chemical Bonding 75

and eight valence electrons according totheir position in the periodic table. TheLewis structure for the first element in eachfamily of the representative elements andnoble gases are shown in Figure 7.5.

As mentioned previously, in 1916,Lewis noted that noble gases were particu-larly stable and did not form compounds.Lewis used these facts to formulate the octetrule. The noble gases have their outer elec-tron shell filled with eight electrons.(Helium is an exception with only two elec-trons in its outer shell.) The octet rule saysthat the most stable electron configurationof an atom occurs when that atom acquiresthe valence electron configuration of a noblegas. That is, when an atom can acquire eight(octet) electrons in its valence shell (or twofor hydrogen to become like helium).

Using the octet rule and knowing thevalence configuration of the elements ex-plains why ionic compounds form. Strongmetals such as the alkalis and alkaline earthshave a strong tendency to lose electrons.When an alkali metal loses an electron, itobtains an octet valence electron configura-tion and acquires a �1 charge. Alkaline earthmetals lose two electrons and become cationswith a �2 charge. Conversely, nonmetalstend to accept electrons. The halogens havehigh electron affinities and readily accept asingle electron to become anions with a–1charge. By accepting a single electron, halo-gens obtain an octet electron configuration.Similarly, elements such as oxygen and sulfuraccept two electrons and acquire a–2 charge.When atoms of a strong metal and nonmetalreact, electrons are transferred from the metalto nonmetal. Because the metal and nonmetal

ions produced in this process have oppositecharges, the ions are attracted to each other.This electrical attraction is the ionic bond.

We can use Lewis dot formulas to repre-sent the transfer of electrons in the formationof ionic compounds. For example, the forma-tion of the ionic compound sodium fluoride,NaF, can be represented using Lewis dot for-mulas and valence electron configurations:

Sodium donates an electron to fluorine, andin the process, sodium becomes the sodiumion, Na�, and fluorine becomes the fluorideion, F–. The net outcome of the transfer ofan electron in this case results in bothsodium and fluorine obtaining a valenceelectron configuration similar to the noblegas neon. The Na� and F– are held togetherby the electrostatic attraction between theoppositely charged ions.

More than a single electron may betransferred when an ionic bond is formed.Two examples of ionic compound formationinvolving the transfer of more than one elec-tron are the formation of calcium chlorideand magnesium oxide:

Figure 7.5Lewis’s Dot Structure of Representative Elements and Noble Gas

76 Chemical Bonding

In the first reaction, a single calcium atomdonates two electrons to two chlorine atoms.This results in the formation of a Ca2� ionand two Cl– ions. The positive ion and twonegative ions are attracted to each other. Inthe formation of magnesium oxide, magne-sium donates two electrons to oxygen result-ing in Mg2� and O2–

In summary, ionic bonds form whenthere is a transfer of electrons betweenatoms of different elements. The result ofthis transfer produces oppositely chargedions. The ions produced generally obtain thevalence electron configuration of noblegases, that is, conform to the octet rule. Theoppositely charged ions produced are heldtogether by electrostatic attraction. Thisattractive force is the ionic bond.

Covalent BondsLewis and many other chemists had rec-

ognized the shortcomings of the ionic bond.When diatomic molecules, such as H2 orCl2, were considered, there was no reasonwhy one atom should lose an electron andan identical atom should gain an electron.There had to be another explanation for howdiatomic molecules formed. We have seenhow the octet rule applies to the formationof ionic compounds by the transfer of elec-trons. This rule also helps explain the for-mation of covalent bonds when molecules(covalent compounds) form. Covalent bondsresult when atoms share electrons. Usingfluorine, F, as a representative halogen, wecan see how the octet rule applies to the for-mation of the F2 molecule. Each fluorineatom has seven valence electrons and needsone more electron to achieve the stable octetvalence configuration. If two fluorines sharea pair of electrons, then the stable octet con-figuration is achieved:

A shared pair of electrons creates a sin-gle covalent bond in F2. Often the doubledots forming the bond are replaced with ashort dash to represent the bond (F-F). Theother halogens form diatomic moleculeswith single bonds in a similar fashion.

Other diatomic molecules, such as oxy-gen and nitrogen, are also held together bycovalent bonds, but they share more than asingle pair of electrons. Oxygen, with its sixvalence electrons, needs two additional elec-trons to achieve an octet. If two oxygensshare four electrons, each oxygen atomobtains the stable valence configuration ofneon. The four shared electrons create adouble bond.

Similarly, sharing six electrons unites twonitrogen atoms. Each nitrogen obtains anoctet, and the two nitrogen atoms are unitedby a triple bond.

Electronegativity and thePolar Covalent Bond

It is easy to see from the examples in theprevious section how two identical atomscan share electrons to achieve an octet andform diatomic molecules. Because each ofour examples dealt with identical atoms, theelectrons can be considered to be sharedequally by each atom. The bond formedwhen the atoms are equally shared can bethought of as a pure covalent bond. But whathappens in covalent compounds? Remem-ber, a compound contains two different elements. When atoms of two different ele-ments are held together by covalent bonds,there is an unequal sharing of the electrons.The sharing of electrons in a covalent bondmay be compared to you and a friend shar-ing a flashlight while walking down a darkstreet. If you and your friend both held the

Chemical Bonding 77

light exactly between you, there would be anequal sharing of the light. It’s much morelikely, though, that one of you holds theflashlight and this person receives more ofthe light. The light, while shared, is notshared equally. When atoms of two differentelements share electrons, one element willhave a greater attraction for the electrons. Ameasure of an element’s attraction for ashared pair of electron is given by the ele-ment’s electronegativity. The basic conceptof electronegativity dates back to the earlynineteenth century, but it was not until 1932that Linus Pauling (1901–1994) was able tocalculate electronegativity. Pauling usedbond energies to derive a number to repre-sent an element’s attraction for a shared pairof electrons in a covalent bond. Table 7.1gives the electronegativities of the commonelements. The higher the element’s elec-tronegativity, the greater the element’sattraction for the shared electrons.

Table 7.1 demonstrates that the ele-ments with the highest electronegativitiesare located in the upper right corner of theperiodic table and electronegativities de-crease moving down and to the left. Fluo-rine has the highest electronegativity. Using

the electronegativities, we can represent thesharing of electrons in a more realistic man-ner. For example, for HCl we see the elec-tronegativity of hydrogen is 2.1 and chlorineis 3.0. Therefore, chlorine has a greaterattraction for the bonding pair of electronsin the HCl molecule. Because the sharedpair of electrons forming the covalent bondis more attracted to the chlorine in themolecule, the bond is referred to as a polarcovalent bond. Because of the unequalsharing of the electrons in HCl, the H partof the molecule tends to have a positivecharge, and the chlorine part a partial nega-tive charge. Hydrogen chloride is a polarmolecule. Whenever we have diatomicmolecules consisting of two different ele-ments, the molecule is generally polar. Itshould be remembered that the electronssurrounding atoms should not be viewed asstationary discrete points. Electrons are inconstant motion in regions best described byquantum theory. For HCl, it is more appro-priate to say that the electron pair formingthe covalent bond spends more time near thechlorine atom, or the probability of findingthe bonding pair nearer the chlorine isgreater. Figure 7.6 represents this situation.

Table 7.1Electronegativities of Elements

78 Chemical Bonding

The point to emphasize with polar covalentbonds is that atoms do not equally share thebonding electrons.

Electronegativities provide a conve-nient tool for classifying the type of bond-ing present in a compound. The closer theelectronegativities of two elements, thegreater the degree of sharing of electrons.In the extreme case for homonucleardiatomic molecules (H2, Cl2, O2, etc.) theelectronegativities are equal, and the bondis a pure covalent bond. As the differencein electronegativities increases, the bondbecomes less covalent and more ionic.While we can characterized bonds as cova-lent, polar covalent, and ionic, these termsshould be interpreted as describing bondsin a general sense. The true nature of abond should be viewed more as a contin-uum moving from ionic when the differ-ence in electronegativities is large to purecovalent when the electronegativities arethe same. Table 7.2 is useful when usingelectronegativity to characterize the natureof a bond.

Polar Molecules and Hydrogen Bonds

In the previous section, HCl wasdescribed as a polar molecule containing apolar covalent bond. The unequal sharing ofelectrons in a bond leads to what is referredto as a dipole. A dipole is symbolized usinga modified arrow pointing toward the nega-tive end of the dipole. For HCl, this can berepresented as

The dipole in HCl causes the H end of themolecule to possess a partial positive chargeand the Cl end to possess a partial negativecharge. This is denoted as

��HCl

with the character indicating a partialcharge.

The dipole created in HCl creates apolar molecule. The presence of polar cova-lent bonds does not necessarily mean themolecule will be polar. Some moleculescontain two or more dipoles that cancel togive a nonpolar molecule. For example, inCO2 the two oxygens are attached to carbonby polar covalent bonds.

The more electronegative oxygens attractthe shared electrons that form the doublecovalent bonds. Thus, two polar double

Figure 7.6The electron density is greater nearer thechlorine atom in the HCl molecule. Dotsrepresent the nuclei.

Table 7.2Percent Ionic Character Based on Difference in Electronegativity

Chemical Bonding 79

covalent bonds are present in CO2, butbecause of the symmetry of the dipoles themolecule itself is nonpolar.

When determining whether a moleculeis polar or nonpolar, it is important to con-sider the geometry of the molecule. Carbondioxide is nonpolar because it is a straightmolecule in which the dipoles balance eachother so that the center of negative chargecoincides with the center of positivecharge. Nonpolar CO2 can be contrastedwith H2O:

Water is a bent molecule. In water, the polarcovalent bonds lead to dipoles in which thecenters of positive and negative charge donot coincide. This makes water a polarmolecule.

In water, the hydrogen atoms are bondedto an oxygen atom. Oxygen has a high elec-tronegativity, which means the bond has ahigh polarity. The fact that hydrogen is a small atom also means that its positivenucleus can closely approach an oxygenatom on a neighboring water molecule. Thepartial positive charge of the small hydrogenatoms and partial negative charge of oxy-gen create a highly polar molecule, which creates a strong attraction between watermolecules. The strong attraction betweenhydrogen in one water molecule and oxygenin another water molecule is called thehydrogen bond. Hydrogen bonding is notunique to water, but occurs when hydrogenis bonded to atoms of oxygen, nitrogen, orfluorine. The fact that hydrogen is small cou-pled with the high electronegativities of oxy-gen, nitrogen, and fluorine lead to hydrogenbonding. This is why compounds of theseelements have unique physical properties,for example, high boiling points, high heatcapacity.

Bond EnergiesA bond is a force holding atoms

together. When atoms come together to formcompounds and molecules, the atoms ac-quire a more stable electron configuration.The more stable electron configurationmeans that the chemical potential energy ofthe system is lowered, and energy is released.To decompose a compound, energy must besupplied to break the chemical bond(s) hold-ing the atoms together. The energy necessaryto break a chemical bond is referred to as thebond energy. Bond energy is measured inkilojoules per mole of bonds. Average bondenergies for some common covalent bondsare given in Table 7.3.

From Table 7.3, it is seen that doublebonds are stronger than single bonds, andtriple bonds are stronger than double bonds.Table 7.3 shows that the bond energies forcovalent bonds are about several hundredkilojoules per mole. For comparison, theamount of energy associated with ionicbonds is given by the lattice energy. The lat-tice energy is the force that holds ionstogether in an ionic compound. Lattice ener-gies range from approximately 1,000 kilo-joules per mole to several thousandkilojoules per mole. This indicates that ionicbonds are roughly a magnitude strongercompared to covalent bonds. The strength ofhydrogen bonds ranges between 10 and 40kilojoules per mole, a magnitude smallerthan covalent bonds.

Table 7.3Bond Energies in kJ per Mol

80 Chemical Bonding

The Metallic BondSeveral properties characteristic of met-

als include high thermal and electrical con-ductivity, ductibility, and malleability. Theproperties of metals can be attributed to thespecial forces holding metal atoms together.In a metal, the atoms are arranged in aclosely packed structure. Rather thanvalence electrons being shared between twoadjacent atoms, the electrons move freelythrough the crystalline structure. For exam-ple, in solid sodium eight other sodiumatoms surround each sodium atom, asshown in Figure 7.7.

Each sodium atom contains a 3s1

valence electron. Metals have a tendency tolose electrons and form positive ions. Thevalence electrons loosely held in metals arepooled and belong to the crystal structure asa whole. The positive metal ions are contin-ually being formed as they lose their valenceelectrons, and these electrons are sharedamong the metal atoms. In this situation, we

say the electrons are delocalized and arefree to move throughout the crystal. Figure7.8 represents the sodium ions embedded ina sea of the pooled electrons.

The Na� ions are surrounded by a poolof delocalized electrons, which acts as the“glue” that holds the metallic atoms to-gether. This arrangement is referred to as themetallic bond. This mobile pool of electronsaccounts for the characteristic properties ofmetals. For example, because the electronsare loosely attached, rigid bonds are notformed and atoms can easily be shapedbecause the electrons move freely through-out the structure.

Bonding and MolecularGeometry

While the basic principles of bondingprovide insights on how and why atomscombine to form compounds and molecules,the information does not give us a picture ofthe structure of substances. Molecular struc-ture depends on the interaction between thevalence electrons of atoms making up themolecule. We know that electrons carry anegative charge and like charges repel eachother. In molecules containing two or morebonds, the shape or molecule geometry ofthe molecule can be predicted using a modelknown as the valence shell electron pairrepulsion (VSEPR) model. As the name ofthe model implies, the shape of the moleculeis derived from the repulsion of electronpairs in the molecule.

A few examples will illustrate howVSEPR is used to predict molecular geom-etry. Beryllium chloride, BeCl2, has theLewis structure

Cl––Be––Cl

The central beryllium atom has two bond-ing pairs of electrons. These bonding pairsrepel each other and attempt to remain as far

Figure 7.7Arrangement of Sodium Atoms in MetallicSodium

Figure 7.8Sodium Ions in a Pool of DelocalizedValence Electrons

Chemical Bonding 81

apart as possible. The arrangement that sat-isfies this condition is linear, with each chlo-ride atom separated by an angle of 180°.

Now consider boron trifluoride, BF3 inwhich boron is covalently bonded to threefluorine atoms. For the electron pairs toremain as far apart as possible, the arrange-ment is triangular with a 120° angle betweenfluorines:

The Lewis dot structure for BF3 shows thecentral boron atom being surrounded byonly six electrons, which violates the octetrule. This illustrates that the octet rule, whileproviding general guidelines, has excep-tions.

For a final example, let’s considermethane, CH4 with four bonding electronssurrounding the central carbon atom.VSEPR predicts a tetrahedral arrangementwith bond angles of 109.5°.

In each of our examples, only bondingpairs of electrons surrounded the centralatom. Many times the central atom containsa single lone pair or lone pairs of electronsin addition to bonding pairs. The presenceof lone pairs, which occupy more space thanbonding pairs, affects the repulsive forcesbetween the valence electrons in molecules.

When lone pairs are present, several differ-ent repulsive forces exist: between bondingpairs, between bonding pair and a lone pair,and between lone pairs. The VSEPR modelis based on different repulsive forces for thethree situations. These are shown here:

The difference in repulsive forces betweenelectron pairs means that when lone pairsare present the geometries change. Let’sexamine two common substances to seehow the presence of lone pairs affects thegeometries of molecules. Ammonia, NH3,contains three bonding pairs and one lonepair surrounding the nitrogen atom:

The four pairs of electrons surroundingnitrogen are similar to the electron structurein methane. Methane, though, has fourbonding pairs, and ammonia has three bond-ing pairs and one lone pair. The repulsiveforce between the single lone pair and threebonding pairs is greater than the repulsiveforce between the bonding pairs. Thisresults in the hydrogens being “squeezed”together so that the H–N–H bond angles are 107° rather than the 109.5° found inmethane.

Water contains two bonding pairs andtwo lone pairs:

Greaterrepulsion

�

�lone pair and lone pairlone pair and bonding pairbonding pair and bonding pair

82 Chemical Bonding

Like ammonia, the structure is similar to thetetrahedral structure of methane. The twolone pairs repel each other in order to be asfar apart as possible. The squeezing of thehydrogens in water is even greater than thatin ammonia. The H–O–H bond angle inwater is 104.5°.

The VSEPR model lets us predictmolecular geometry based on the number ofbonding pairs and lone pairs on the centralatom. Table 7.4 summarizes several of thesimpler geometries based on the number ofbonding pairs and lone pairs possessed bythe central atom.

Modern Bonding TheoryThe use of Lewis dot structures and

VSEPR to explain bonding and predictmolecular geometry portrays electrons as dis-crete, fixed units of charge. It is important toremember that in the modern quantum viewof the atom electrons occupy a probabilisticregion of space surrounding the nucleus.Electrons are described in terms of wavesusing a probabilistic model. While quantumtheory is not required to understand the basicconcepts of bonding, the actual structure andproperties of many substances can beexplained only using quantum mechanical

Table 7.4Simple Molecular Geometries

Chemical Bonding 83

bonding theories. Two modern theories usedto describe bonding are the valence bond the-ory and the molecular orbital theory. Thesetwo models can be compared by applyingthem to the hydrogen molecule, H2.

Valence bond theory assumes that cova-lent bonds form by atoms sharing valenceelectrons in overlapping valence orbitals. Invalence bond theory, the individual atomspossess valence orbitals that take part inbonding. The overlap of valence orbitals isused to explain bonding. For the hydrogenmolecule, each hydrogen atom has onevalence electron in a 1s orbital. As twohydrogen atoms approach each other, eachelectron is attracted to the other atom’snucleus, while the electrons repel eachother. The overlapping 1s orbitals that resultwhen the hydrogen atoms combine is thecovalent bond:

In valence bond theory, the strength of thebond depends on the degree of overlap. Thegreater the overlap, the stronger the bond.One feature of the bond valence theory isthat the orbitals may combine to producehybrid orbitals. For example, an s and threep orbital may combine to form four sp3

hybrid orbitals.A second quantum mechanical bonding

theory is molecular orbital theory. This the-ory is based on a wave description of elec-trons. The molecular orbital theory assumesthat electrons are not associated with anindividual atom but are associated with theentire molecule. Delocalized molecularelectrons are not shared by two atoms as inthe traditional covalent bond. For the hydro-gen molecule, the molecular orbitals areformed by the addition of wave functions foreach 1s electron in each hydrogen atom. Theaddition leads to a bonding molecular

orbital when the waves reinforce each otherand to an antibonding orbital when thewaves cancel each other. A bonding molec-ular orbital diagram for the hydrogenmolecule is shown in Figure 7.9.

The two electrons in the hydrogenmolecule occupy the bonding orbital withthe lowest energy. The energy of this orbitalis lower than the energy of the electrons inindividual hydrogen atoms resulting in amore stable configuration.

One way of understanding molecularorbitals is to relate them to atomic orbitals.We know electrons fill atomic orbitals start-ing at the lowest energy levels. As moreelectrons are added, the atomic orbitals fill.Certain electron configurations are morestable than others, for example, the noblegas configuration. In molecular orbital the-ory, electrons fill molecular orbitals. Elec-trons in certain molecular orbitals are morestable than when they exist as electrons inindividual atoms. The stability of molecularorbitals is useful in predicting whether amolecule will form. The more stable the for-mation of molecular orbitals is, the morelikely a molecule will form. Molecularorbitals in which the arrangement of elec-trons produces a higher-energy, less-stableconfiguration predict nonbonding situations,that is, a molecule will not form. For exam-ple, the molecular orbital diagram for He2

results in two electrons in a bonding orbitaland two electrons in a nonbonding orbital.

Figure 7.9Bonding Molecular Orbital Diagram for H2

There is no decrease in energy in thisarrangement; therefore, we can concludethat He molecules do not form.

There are many aspects to valence bondtheory and molecular orbital theory. Eachtheory has its strengths and weaknesses, and

chemists must use both theories individuallyand collectively to explain chemical behav-ior. This section has provided just a briefglimpse into modern bonding theory. Moredetailed introductions can be found in manycollege texts.

84 Chemical Bonding

Introduction

The previous chapter dealt with chemicalbonding and the forces present between theatoms in molecules. Forces between atomswithin a molecule are termed intramolecu-lar forces and are responsible for chemicalbonding. The interaction of valence elec-trons between atoms creates intramolecularforces, and this interaction dictates thechemical behavior of substances. Forcesalso exist between the molecules them-selves, and these are collectively referred toas intermolecular forces. Intermolecularforces are mainly responsible for the physi-cal characteristics of substances. One of themost obvious physical characteristics relatedto intermolecular force is the phase or phys-ical state of matter. Solid, liquid, and gas arethe three common states of matter. In addi-tion to these three, two other states of matter exist—plasma and Bose-Einsteincondensate.

Plasma is a gas in which most of theatoms or molecules have been stripped ofsome or all of their electrons. The electronsremoved still stay within the vicinity of theatom or molecule. This state generallyoccurs when a gas is subjected to tremen-

dous temperatures, is electrically excited, oris bombarded by radiation. Plasma is theprincipal component of interstellar spaceand comprises over 99% of the universe.The plasma state is also present in gas-emitting light sources such as neon lights.The state of matter known as the Bose-Einstein condensate was theoretically pre-dicted by Satyendra Nath Bose (1894–1974)and Albert Einstein (1879–1955) in 1924. ABose-Einstein condensate forms whengaseous matter is cooled to temperaturesmillionths of a degree above absolute zero.At this temperature atoms, which normallyhave many different energy levels, coalesceinto the lowest energy level. At this level,atoms are indistinguishable from oneanother and form what has been termed a“super atom.” The first experimentally produced Bose-Einstein condensate wasreported 1995, and currently, a number ofresearch groups are investigating this newstate of matter.

In this chapter, we focus on solids andliquids. In the next chapter, we turn our atten-tion to gases. Except where noted, the dis-cussion focuses on the states of matter undernormal atmospheric conditions, that is, 1atmosphere pressure and room temperature.

8

Intermolecular Forces and the Solid and Liquid States

86 Intermolecular Forces and the Solid and Liquid States

Characteristics of Solids,Liquids, and Gases

Most of the naturally occurring ele-ments exist as solids under normal condi-tions (Figure 8.1). Eleven elements exist asgases, and only two exist as liquids.

Liquids and gases share some commoncharacteristics, such as the ability to flow andtogether are classified as fluids. Likewise, liq-uids also share some characteristics withsolids, and these two states are called con-densed states. Solids, liquids, and gases canbe differentiated by their properties. Solidsand liquids have relatively strong intermolec-ular attraction compared to gases, in whichthe intermolecular attraction can be consid-ered to be nonexistent under normal condi-tions. Solids and liquids contain moleculesand ions arranged close together with littleempty space between particles. Therefore,compared to gases, solids and liquids havehigh densities and are nearly incompressible.Gases are characterized by predominantlyvoid space through which the molecules

move. In solids, molecules and ions are heldrigidly in place preventing any appreciablemovement of particles with respect to oneanother. Solid have a fixed shape and volume.In liquids, the intermolecular forces areweaker than in solids, and so liquids have theability to flow and take the shape of the objectcontaining the liquid. Table 8.1 summarizessome of the characteristics of the solid, liquid,and gas states.

Intermolecular ForcesIntermolecular forces are responsible

for the condensed states of matter. The par-ticles making up solids and liquids are heldtogether by intermolecular forces, and theseforces affect a number of the physical prop-erties of matter in these two states. Inter-molecular forces are quite a bit weaker thanthe covalent and ionic bonds discussed inChapter 7. The latter requires several hun-dred to several thousand kilojoules per moleto break. The strength of intermolecularforces are a few to tens of kilojoules per

Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu

Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr

H He

Li Be B C N O F Ne

Na Mg Al Si P S Cl Ar

K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr

Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe

Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn

Fr Ra Ac Rh Db Sg Bh Hs Mt Uun Uuu Uub

Figure 8.1States of elements under normal conditions. Light gray indicates gas, dark gray indicatesliquids, all the rest are solids.

Intermolecular Forces and the Solid and Liquid States 87

mole. One of the strongest and most impor-tant types of intermolecular forces is thehydrogen bond. Its strength can be as highas 40 kilojoules per mole. The hydrogenbond was introduced in Chapter 7. Let’sexamine the hydrogen bond before lookingat several other types of intermolecularattractions.

Hydrogen bonds are formed betweenmolecules that contain hydrogen covalentlybonded to atoms of oxygen, fluorine, ornitrogen. Hydrogen bonded to any one ofthese elements produces a highly polarmolecule. The hydrogen takes on a partialpositive charge, and the electronegative ele-ment has a partial negative charge. Thehydrogen on one molecule has a high attrac-tion for the electronegative atom on an adja-cent molecule. It is important to rememberthat the hydrogen bond is an intermolecularforce.

Water is the most common substancedisplaying hydrogen bonding. The hydrogenon one molecule is attracted to an unsharedpair of electrons held by oxygen on an adja-cent water molecule (Figure 8.2). Because

each oxygen in water has two unshared pairsof electrons and there are also two hydrogenatoms in water, a network of water moleculesforms, held together by hydrogen bonds.Hydrogen bonding has a pronounced effecton the physical properties of water. Forexample, the boiling points of compounds

Table 8.1Characteristics of Solids, Liquids, and Gases

Figure 8.2Hydrogen bonding in water. Hydrogenbonds are represented by dashed lines. Onlyseveral water molecules are shown and notall bonds are displayed. The pattern for thecentral oxygen atom continues to form a net-work of hydrogen-bonded water molecules.


generally decrease up a group in the periodictable. The reason for this decrease involvesintermolecular forces and is explainedshortly. In the oxygen group, the boilingpoints of H2Te, H2Se, H2S, and H2O aregraphed in Figure 8.3.

The trend follows the expected trend forthe first three compounds with the boilingpoint decreasing from the largest compoundH2Te to H2S. According to the normal pat-tern, water should have the lowest boilingpoint. If the trend displayed for the first threecompounds continued, then water’s boilingpoint should be lower than that of H2S. Thereason that water has an abnormally highboiling point compared to similar com-pounds is that water molecules exhibithydrogen bonding. Because the watermolecules are hydrogen bonded, it takesmore energy to cause water to leave the liq-uid state and enter the gas state. It’s just asthough you were with a group of friends,and someone tried to pulled you away fromthe group. It would be a lot harder to pullyou, or any of your friends, away from thegroup if you were all holding on to eachother. In essence, hydrogen bonding is howwater molecules hold on to each other. Apattern similar to that shown in Figure 8.3exists for Groups 15 and 17 hydrogen com-pounds with ammonia, NH3 and hydrogenfluoride, HF, having the highest boilingpoints in these two groups, respectively. This

is not surprising because both ammonia and hydrogen fluoride exhibit hydrogenbonding.

Another interesting property of water isthat its maximum density is at 4°C. In gen-eral, almost all substances have a greaterdensity in their solid state than in their liq-uid state. Hydrogen bonding is also respon-sible for this unique property of water.Because each water molecule has two hydro-gen atoms and two lone pairs of electrons,the water molecules can form a three dimen-sional network of approximately tetrahe-drally bonded atoms. Each oxygen iscovalently bonded to two hydrogen atomsand also hydrogen bonded to two oxygenatoms (Figure 8.2). When water exists as ice,the molecules form a rigid three-dimensionalcrystal. As the temperature of ice increasesto the melting point of 0°C, hydrogen bond-ing provides enough attractive force to main-tain the approximate structure of ice. Theincrease in temperature above water’s melt-ing point causes thermal expansion, and thiswould normally lead to an increase in vol-ume and a corresponding decrease in den-sity. Remember, density is mass divided byvolume so a larger volume results in asmaller density. The reason water actuallybecomes more dense is that at the meltingpoint and up to 4°C there is enough energyfor some of the water molecules to overcomethe intermolecular attraction provided by thehydrogen bonding. These water moleculesoccupy void space in the still-presentapproximate crystalline structure. Becausemore molecules of water are occupying thesame volume, the density increases. Thisphenomenon continues until the maximumdensity is reached at 4°C. At this point, thetrapping of additional water molecules in thevoid space is not great enough to overcomethe thermal expansion effect that lowers thedensity; therefore, solid ice is less dense thanliquid water.

Figure 8.3Trend in Boiling of Several Group 16 Compounds


The fact that the maximum density ofwater is at 4°C has significant environmen-tal impacts. Consider the freezing of a fresh-water lake in winter. As the temperaturedecreases, the surface water becomes pro-gressively denser and sinks to the bottom.This process helps carry oxygen from thesurface to deeper water. The cycling of alake’s waters also helps to bring nutrients tothe surface. This process is sometimesreferred to as the fall overturn. Once thelake reaches a temperature of 4°C, the wateris at its maximum density. Further coolingresults in less dense water, which doesn’tsink, eventually forming a layer of ice iftemperatures are cold enough. The ice layereffectively insulates the rest of the lake.Because the water was replenished withoxygen during the fall overturn, the lakegenerally does not become anoxic. If waterbehaved like most substances, the entirelake would cool down to its freezing point,and then the entire lake would freeze solid.

Several other unique properties of watercan be attributed to hydrogen bonding. Wediscuss these after we have had a chance toexamine some of the other intermolecularforces. The hydrogen bond is actually anunusually strong form of what is known asa dipole-dipole force. Polar molecules giverise to dipole-dipole forces. A dipole-dipoleforce results from the electrostatic attractionbetween the partial positive and negativecharges present in polar molecules. Thestrength of the dipole-dipole force is directlyproportional to the strength of the dipolemoment of the molecule.

Just as two polar molecules, like oppo-site ends of a magnet, are attracted to eachother, a polar molecule may be attracted toan ion. This gives rise to an ion-dipoleforce. The negative ends of polar moleculesare attracted to cations and the positive endto anions. The charge on the ion and thestrength of the dipole moment determine the

strength of the ion-dipole force. When anionic compound such as salt, NaCl, dis-solves in water, ion-dipole forces come intoplay. The positive hydrogen end of water andnegative chloride ion are attracted to eachother, while the negative oxygen end andpositive sodium ion are attracted to eachother (Figure 8.4). Polar molecules tend to be soluble in water, while nonpolarmolecules are insoluble. This fact is illus-trated by a mixture of oil and water. Thenonpolar oil does not dissolve in water, andtwo separate layers result. A similar effectoccurs with the polar vinegar and nonpolaroil portions of a salad dressing.

Hydrogen bonding, dipole-dipole, andion-dipole forces all involve polar mole-cules, and all involve electrostatic attrac-tions between opposite charges. A forcethat is present in both polar and nonpolarmolecules is known as a dispersion orLondon force. The London force is namedfor Fritz London (1900–1954) whodescribed this force in 1931. Dispersionforces are a consequence of the quantumnature of the atom. Electrons move aroundatomic nuclei in a probabilistic fashion, butthey occupy specific energy levels. In non-polar molecules, on average, the movementof the electrons are such that the centers ofpositive and negative charge are the same.It is important to realize, though, that this isan average state and does not representwhat is occurring at any one instance.Because the electrons are in constantmotion, the electrons assume an asymmet-rical distribution producing an temporarydipole. An instant later, this situation

Figure 8.4Ion-Dipole Forces. M� represents a cationand X� an anion.


changes. Because electrons are constantlymoving and changing position, temporarydipoles are continually being formed anddissolved (Figure 8.5). Temporary dipolesinduce other temporary dipoles in adjacentatoms. That is, the negative and positiveends of the dipole will attract or repel electrons in adjacent atoms. This continual process leads to an ever-present attractionbetween nonpolar (as well as polar)molecules.

Dispersion forces depend on temporarydipoles inducing dipoles in adjacent atomsor molecules. The ease with which electronsin an atom or molecule can be distorted toform a temporary dipole is known as polar-izability. The larger the atom or moleculeis, the greater its polarizability. This isbecause larger atoms have more electrons,and these electrons are located farther fromtheir respective nuclei. Electrons fartherfrom the nucleus are held less tightly, andthis results in their greater polarizability.This is why larger molecules in a series tendto have higher boiling points. For example,the boiling points of methane (CH4),propane (C3H8), and butane (C4H10) are�161°C, �42°C, and 0°C, respectively.

An example of how dispersion forcesand polarizability affect physical propertiesis seen in the halogens. Moving down the

halogen group fluorine and chlorine aregases, bromine is a liquid, and iodine asolid. This is just what would be expected.Dispersion forces become stronger movingdown the halogen group as the atomsincrease in size and are more polarizable.Another illustration is seen in Table 8.2which lists the freezing point of the Nobelgases. The freezing point increases down thegroup. If the temperature of a mixture of thegases listed in the table was lowered, Xenonwould freeze first because of the presence ofgreater dispersion forces.

Dipole-dipole, ion-dipole, and disper-sion forces are collectively known as vander Waals forces. Johannes Diederick vander Waals (1837–1923) received the 1910Nobel Prize in physics for his work on flu-ids. We have seen how hydrogen bondingand van der Waals forces affect the physicalproperties of substances, and more is saidabout these forces as we examine the differ-ent states of matter.

Crystalline SolidsSolids are the most obvious state of mat-

ter for us; they are ubiquitous and come inall shapes and sizes. Solids are classified aseither crystalline solids or amorphoussolids. A crystalline solids displays a regu-lar, repeating pattern of its constituent parti-cles throughout the solid. At the microscopic

Figure 8.5On average, the electron distribution inhelium is symmetrical, and the centers ofpositive and negative charge are the same asshown on the left. At any instant, though, theelectrons may assume an asymmetrical dis-tribution as shown on the right. In this situ-ation, a temporary dipole is formed.

Table 8.2Freezing Points of Noble Gases


level, crystalline solids appear as crystals.An amorphous solid does not display a reg-ular, repeating pattern of its constituent par-ticles. Crystalline solids are composed ofatoms, molecules, or ions that occupy spe-cific positions in a repeating pattern. Theposition occupied by the particles arereferred to as lattice points. The most basicrepeating unit making up the crystallinesolid is known as the unit cell. The unit cellrepeats itself throughout the crystalline solid(Figure 8.6).

Crystalline solids can be classifiedaccording to the type of particles occupyingthe lattice points and the type of inter-molecular forces present in the solid. Ionsoccupy the lattice points in an ionic solid,and the crystal is held together by the elec-trostatic attraction between cations andanions. Common table salt is an example ofa crystalline solid. In molecular solids,molecules occupy the lattice points, and vander Waals forces and/or hydrogen bondingpredominate in this type of crystalline solid.Common table sugar, sucrose, is an exam-ple of a common molecular solid. Covalentcrystals are generally composed of atomslinked together by covalent bonds in a three-dimensional array. Carbon in the form ofgraphite or diamond is an example of acovalent solid. The final type of solid is ametallic solid in which metallic bonding is

characterized by a delocalized sea of elec-trons holding together metal cations.

Much of our knowledge on the structureof crystalline solids comes from x-ray crys-tallography or x-ray diffraction. Diffrac-tion is the scattering or bending of a waveas it passes an obstacle. X-rays exist as elec-tromagnetic waves. They are particularlyuseful for probing crystalline solids,because their wavelengths are roughly onthe same scale as that of atoms, 10–10m. Vis-ible light cannot be used for probing crys-talline solids, because the wavelength ofvisible light is far too large to produce animage. A simple analogy may help to putthis concept in perspective. Say you wantedto pick up a very tiny grain of rice with yourfingers. It would be impossible to use yourthumb and index finger in the usual mannerto do this; your fingers are just too large. Youcould pick up the grain with a pair of finetweezers, though. Just as your fingers are toolarge to perform certain tasks, visible lightis too large to probe at the atomic level. X-rays are the right size to “observe” crys-talline solids.

Max Theodor Felix von Laue (1879–1960) first predicted x-rays could bediffracted by a crystal in 1912. He receivedthe Nobel Prize in 1914 for his discovery ofx-ray crystallography. One year later, Wil-liam Henry Bragg (1862–1942) and his sonWilliam Lawrence Bragg (1890–1951)shared the Physics Nobel Prize for expand-ing on von Laue’s work. In x-ray crystallog-raphy, an x-ray beam is focused through acrystal. The particles making up the crystalscatter the x-rays resulting in a pattern ofconstructive interference and destructiveinterference. Constructive interference re-sults when waves are in phase and reinforceeach other; in destructive interference, thewaves are out of phase and cancel each other(Figure 8.7). Because the particles (atoms,molecules, or ions) making the crystalline

Figure 8.6A crystalline solid is composed of a repeat-ing unit known as the unit cell. Latticepoints are occupied by atoms, molecules, orions.


solid have an ordered arrangement, the x-rays scattered by the particles produce anordered pattern. The pattern can be capturedon a photographic plate or film with darkareas representing constructive interferenceand light areas showing destructive interfer-

ence (Figure 8.8). The diffraction patterncan then be used to help decipher the struc-ture of the crystal.

The x-ray diffraction pattern revealshow the particles are arranged in a crys-talline structure. A number of commonarrangements exist among crystalline solids.These can be classified according to thearrangement of particles making up the unitcell. Three common arrangements are sim-ple cubic, face-centered cubic, and body-centered cubic (Figure 8.9). In sodiumchloride, the Na� and Cl– ions are posi-tioned in a face-centered cubic arrangement.The alkali metals follow a body-centeredcubic arrangement. In addition to the ar-rangement of particles in the crystallinesolid, interpretation of the x-ray diffractionpattern allows chemists to determine thedistance between particles and how muchvoid space is present in the solid.

Amorphous SolidsThe term “amorphous” comes from the

Greek word for shape, “morph,” and meansdisordered or without shape (a � morph).Amorphous solids do not display a regularthree-dimensional arrangement of particles.The most common amorphous solid is glass.We normally use the term “glass” in associ-ation with silica-based materials, althoughthe term is sometimes used with any amor-phous solid including plastics and metals.The technical definition for glass is an opti-cally transparent fusion product of inorganic

Figure 8.7In constructive interference, waves reinforceeach other. In destructive interference,waves cancel each other.

Figure 8.8Picture of X-ray Diffraction Pattern (RaeDéjur)

Figure 8.9Three Simple Unit Cells for CrystallineSolids


material that has cooled to a rigid state with-out crystallizing. Our discussion in this sec-tion will primarily focus on silica-basedglass. Both crystalline quartz and the silicaglass are primarily composed of silica, SiO2.The difference between quartz and glass isthat quartz displays long-range order (Fig-ure 8.10). In quartz, the silicon is covalentlybonded to the oxygen in a tetrahedralarrangement. When glass is produced, SiO2

is heated to an elevated temperature andthen rapidly cooled. The rapid cooling doesnot allow the SiO2 to form a regular crys-talline structure. The result is a solid thatbehaves like a very viscous liquid when it isheated. Glass is sometimes called a solidsolution and does flow at a very slow rate.This flow can sometimes be seen in old win-dow glass where the bottom is slightlythicker than the top.

Quartz is a crystalline solid and has anorderly structure, but glass is an amorphoussolid that is disordered. The structures areactually arranged in a three-dimensionaltetrahedral pattern. They are shown in Fig-ure 8.10 as two-dimensional representations.

The first silica glass was producedaround 3500 B.C. in Mesopotamia (presentday Syria and Iraq), although there is evi-dence of early production in Egypt andPhoenicia (Lebanon). The melting point ofpure SiO2 is 1713°C, but the mixing of other

substances with the SiO2 lowers its meltingpoint. Egyptians added natron (sodium car-bonate) to SiO2. Some glasses are producedat temperatures as low as 600°C. As the art ofglass making developed, individuals discov-ered how to produce different glasses byadding various substances to the silica melt.The addition of calcium strengthened theglass. Other substances imparted color to theglass. Iron and sulfur gave glass a browncolor, copper produced a light blue, andcobalt a dark blue. Manganese was added toproduce a transparent glass, and antimony toclear the glass of bubbles. Most modernglass produced is soda-lime glass and con-sists of approximately 70% SiO2, 15% Na2O(soda), and 5% CaO (lime). Borosilicateglass is produced by adding about 13%B2O3. Borosilicate glass has a low coeffi-cient of thermal expansion, and therefore,is very heat resistant. It is used extensively inlaboratory glassware and in cooking whereit is sold under the brand name Pyrex®.

LiquidsLiquids share certain properties with

both solids and gases. Like solids, liquidsare almost incompressible and occupy a spe-cific volume; like gases liquids flow. Theintermolecular forces present in liquids areweaker than in solids. Liquid particles arenot held rigidly in place and are able to slidepast one another. The ability to flow variesgreatly among liquids. Anyone familiar withthe saying “slow as molasses” can appreci-ate the difference in the ability of molassesto flow compared to other liquids such aswater. Viscosity is a measure of a liquid’sresistance to flow. The ability of a liquid toflow is related to several factors includingintermolecular forces, size of particles, andstructure. The stronger the intermolecularforces in liquids are, the higher the viscosity.Small particles can slide past neighboring

Figure 8.10Both quartz and silica glass are primarilycomposed of silica, SiO2.


particles easier, and for this reason, sub-stances composed of smaller particles havelower viscosities. Another factor affectingviscosity is the size and shape of molecules.Structures composed of long chains havehigher molecular masses, and a greaterprobability of molecules entangling. Thedependence of size and shape on viscositycan be pictured by contrasting the differencein dumping out a box of marbles versusdumping out a box of rubber bands. Themarbles flow readily out of the box, whilethe rubber bands become entangled witheach other and do not flow smoothly.

Anyone who has heated honey or syrupis aware that viscosity decreases withincreasing temperature. Viscosity is a pri-mary concern in lubricants, and a practicalexample of viscosity deals with the labelingof motor oils. Motor oils have an arbitraryviscosity rating given by their SAE (Societyof Automotive Engineers) numbers. Thehigher the number is, the higher the viscos-ity. This means that oils with higher SAEnumbers are thicker. Most common motoroils are produced to have multiviscous char-acteristics to meet the needs of both coldstart-up conditions and the elevated temper-atures during operation. A typical SAEnumber would be 5W-30 or 10W-40. Thefirst number can be considered the basenumber appropriate for cold conditions (Wstands for winter). Cold conditions are con-sidered to be 0°F. The second number is theequivalent viscosity at an elevated tempera-ture, say 200°F. Multiviscous oils are usedto meet the typical driving cycle. When anengine is cold, it needs a lower viscosity(thinner) oil to lubricate the engine effec-tively. As the engine warms up, the oil insidealso warms and becomes even less viscous.To compensate for the lowering of the vis-cosity with increased temperature, multivis-cous oils contain carbon polymer additives.These polymers are coil shaped at low tem-

peratures and uncoil as the engine tempera-ture increases. The uncoiled polymersincrease the viscosity, thereby compensat-ing for the heating effect.

Another characteristic dependent on theintermolecular forces is the surface tensionof the liquid. Surface tension results fromthe unbalanced forces on molecules at thesurface of a liquid. Figure 8.11 shows howsurface tension results from these unbal-anced forces. Consider water as the liquid inFigure 8.11. A water molecule in the inte-rior of the liquid is surrounded on all sidesby other water molecules. Attractive inter-molecular forces pull the molecule equallyin all directions and these forces balanceout. A water molecule on the surface expe-riences an unbalanced force toward the inte-rior of the fluid. This unbalanced force pullson the surface of the water putting it undertension. This situation is similar to the tight-ening of the head of drum. The tensioncauses the surface of the water to act like athin film. If you carefully use tweezers toplace a clean needle on the surface of water,surface tension will allow the needle to floateven though the needle is denser than water.

Figure 8.11Surface tension results from the fact thatsurface molecules are pulled toward theinterior of the liquid as compared to interiormolecules where the forces are balanced.


Surface tension is directly related to themagnitude of intermolecular forces in a liq-uid. The greater the intermolecular force is,the greater the surface tension. For this rea-son, water has a high surface tension.Because of the surface tension, water dropswill assume a spherical shape. This shapeminimizes the surface area of the drop.Related to surface tension are the propertiesof cohesion and adhesion. Cohesion refersto the attraction between particles of thesame substance. That is, it is the inter-molecular force between molecules of theliquid. Adhesion is the attraction betweenparticles of different substances. Adhesionis seen in the ability of water to rise in thintubes by capillary action. The adhesionbetween water and glass is greater than thecohesion between water molecules. Theglass pulls water up the interior of the tube,and the cohesive force between watermolecules creates a concave meniscus (Fig-ure 8.12). The water rises until the cohesive

force balances the weight of the water pulledup on the side of the tube. Because smallerdiameter tubes contain less volume of water,water rises higher as the tube diameterdecreases.

Certain chemicals have the ability tolower the surface tension of water. Thisallows water to spread out over a surfacerather than bead up. Wetting agents decreasethe cohesive forces between water mole-cules, and this helps water to spread over thesurface of an object by adhesive force.

Material ScienceThe material in this chapter has focused

on the solid and liquid states. Up until recenttimes, there has been a clear distinctionbetween these states. Recent advances inmodern science have led to materials thatblur the distinction between states of matter.Material science is a relatively new inter-disciplinary area of study that appliesknowledge from chemistry, physics, andengineering. Material scientists study theproperties, structure, and behavior of mate-rials. Material science can further be dividedinto branches that deal in areas such asceramics, alloys, polymers, superconduc-tors, semi-conductors, corrosion, and sur-face films. Material scientists continue tomodify our traditional view of the differentstates of matter. A good example of a mate-rial that straddles two states of matter is liq-uid crystals. The term “liquid crystals”sounds contradictory, but it is an appropri-ate name for molecules classified as liquidcrystals. There are close to 100,000 differ-ent organic compounds identified as liquidcrystals. They are slender, rod-like mole-cules that have mobility like a liquid but dis-play order by having their axes line upparallel to each other. An analogy might bethat a liquid is like the people walking inone direction down a busy sidewalk, while a

Figure 8.12In capillary action, the adhesion between thewalls of the tube and the liquid is greaterthan the cohesion between liquid molecules.The liquid rises up the tube and forms ameniscus that is concave upward. Thesmaller the tube is, the greater the heightthe liquid rises.


liquid crystal is like a marching band. Onthe sidewalk, people flow in a definite direc-tion, but they are not aligned. People in themarching band are aligned in columns.There is a definite structure to the marchingband, but members of the band still flow.

Liquid crystals are all around us. Allyou have to do is look at a calculator, a dig-ital clock or watch, or a laptop computer andyou see liquid crystals. Devices with liquidcrystals are so commonplace that it is hardto believe that the first digital watch was notmarketed until 1973. If you examine a liq-uid crystal display closely, you will oftennotice a figure 8 arrangement consisting ofseven bars. The liquid crystals in these dis-plays are controlled by small electric fields;the crystals are very sensitive to these elec-tric fields and orient themselves accordingto the electric field. A liquid crystal displayworks by rotating polarized light so that iteither passes through or is blocked by a sec-ond polarizer. Think of unpolarized light aslight waves vibrating in all directions. In thiscondition, light is said to be incoherent, orin simple terms it is all scrabbled up. Apolarizer lines the light waves up in onedirection making the waves coherent. Inshort, it unscrabbles the light waves. Apolarizer is like an optical grating, allowinglight that has the right orientation to passthrough it. This is like a narrow passageallowing only those people who are turnedsideways to squeeze through. When liquidcrystals cause the light to be blocked by thesecond polarizer, the display registers a darkimage (Figure 8.13). These dark images aredisplayed as numbers or letters in the liquidcrystal display. By mixing other chemicalswith liquid crystal molecules, color displayscan be made.

Liquid crystals are just one of the manyadvances made in material science duringthe last thirty years. Many other new sub-stances with interesting properties have

been discovered. One that captured wide-spread attention starting in 1985 was the dis-covery of buckminsterfullerenes or “buckyballs.” This naturally occurring form of car-bon was named after the American architectF. Buckminster Fuller (1895–1983) whodesigned the geodesic dome. Buckminster-fullerene is said to be the most sphericalmolecule known. The first buckminster-fullerene identified was C-60 by RichardSmalley (1943–), Robert Curl (1933–), andHarold Kroto (1939–). These three sharedthe 1996 Nobel Prize in chemistry for theirpioneering discovery. C-60 has the shape ofa soccer ball with the sixty carbons making

Figure 8.13In a liquid crystal display, unpolarized lightpasses through a polarizer. The polarizedlight is either unaffected (1) or rotated (2) bythe liquid crystals depending on the voltageapplied to the liquid crystals. Light eitherpasses through the second polarizer and isreflected by the mirror producing a lightimage or is blocked by the second polarizerin which case a dark image appears.


up twelve pentagonal and twenty hexagonalfaces (Figure 8.14). Since the original dis-covery of C-60, numerous other fullerenesof various shapes have been produced rang-ing from C-28 to C-240. Research into theirpotential use is currently underway, but theirexpense, which is several times that of gold,currently limits widespread use.

Other advances in material sciencehave helped humans mimic nature in theproduction of certain materials. For exam-ple, in the last half of the twentieth centurywe have learned how to produce syntheticdiamonds. Diamonds were first producedcommercially in the 1950s by General Elec-tric by subjecting graphite to temperaturesof 2,500°C and pressures approaching100,000 atmospheres. Currently, well over ahundred companies produce synthetic dia-monds.

Other advances in material science arebeing made in a number of areas, includingmaterials collectively known as “smart” or“intelligent” materials. These materials havethe ability to adapt to their surroundings. Acommon example is lenses used in eyeglassesmade from material that darken in bright

light. Memory metals have the ability toreturn to a preformed shape at certain tem-perature. These metals are used in orthodon-tic devices and guidewires. A popularmemory metal is the nickel-titanium alloycalled Nitinol. Memory metals assume a“parent” shape above a specific transitiontemperature. Below the transition tempera-ture, the metal can be formed into variousshapes. When the memory metal is heatedabove the transition temperature, the metalgoes back to its parent shape as though it hasremembered it. Shape memory polymers thatwork similar to metals have recently beenintroduced. Advances in memory materialmay make it possible to repair bent objects bysimply heating the object. Another group ofmaterials is known as self-healing materials.To understand how self-healing materials aredesigned, think of how a cut heals. When youcut yourself, a number of signals muster achemical response to the wound. Over sev-eral days the wound heals. Self-healing plas-tics are currently being studied. Materialssuch as plastics develop numerous micro-cracks through everyday use. As time goeson, these micro-cracks propagate and grow,eventually leading to failure of the product.One concept is to strategically position liquidplastic-filled capsules in self-healing plastics.When the plastic object is dropped or han-dled in a certain way, the capsules wouldbreak open and flow into the micro-crackswhere it would harden. This self-healing process would lengthen the normal life of materials. Stealth weaponry is yet anotherexample of how materials can be modified tobehave in a totally new manner. Stealth air-planes are built with materials designed toabsorb radar waves rather than reflect a signalback to the receiver.

Research in material science continuesto modify the substances that make up ourworld. As the twenty-first century unfolds,new materials will improve our lives. Just

Figure 8.14Buckminsterfullerene


consider biomaterials used for bone replace-ment, drug delivery microchips, and skingrafts; composites used in cars, planes, andother transportation vehicles; building mate-rials that are lighter, provide insulation, andhave superior strength. All we have to do istake a good look around us to realize thatmany of the common materials presenttoday were not available even fifty yearsago, for example, teflon, fiberglass, and

semi-conductors. Material scientists con-tinue to improve old materials and createnew ones. In this manner, they are the mod-ern version of the ancient alchemists. Likealchemists, material scientists strive to per-fect the basic elements into something moreperfect, although without the mystical con-nection. As these new alchemists continueto perfect materials, chemistry will play acentral role in their quest.

IntroductionThis chapter continues to explore states ofmatter by focusing on the gas state. Table8.1 summarized the main features of gases.In the gaseous state, molecules are muchfarther apart than in either solid or liquids.Because of this distance molecules in thegaseous state have virtually no influence oneach other. The independent nature of gasmolecules means intermolecular forces inthis state are minimal. A gas expands to fillthe volume of its container. Most of the vol-ume occupied by the gas consists of emptyspace. This characteristic allows gases to becompressible, and gases have only about1/1,000 the density of solids and liquids.

Eleven elements (hydrogen, helium,oxygen, nitrogen, fluorine, chlorine, argon,neon, krypton, xenon, and radon) exist asgases under normal conditions. Addition-ally, many molecular compounds exist asgases. The most common gases are thosefound in our atmosphere. Nitrogen and oxy-gen comprise most of the atmosphere (Table9.1). The third most abundant gas in theatmosphere is argon. Only trace amounts ofcarbon dioxide, helium, neon, methane,

krypton, and hydrogen are present in theatmosphere.

PressureOne of the most important properties

characterizing a gas is its pressure. Pressureis defined as the force exerted per unit area.Atmospheric pressure is the force exerted bythe atmosphere on the Earth’s surface. Acommon device used to measure atmo-spheric pressure is the barometer. Thebarometer was invented in 1643 by Evange-lista Torricelli (1608–1647). Torricelli, whowas a student of Galileo, was presented withthe problem of why water could not raisemore than 32 feet with a suction pump.Galileo explained the rise of water up a pipeby invoking the statement “nature abhors avacuum.” When a pump created a vacuum,Galileo believed water rose to fill the voidcreated, but somehow there was a limit ofapproximately 32 feet. Torricelli reasonedcorrectly that the height to which watercould be pumped was due to atmosphericpressure. In his study of the pump problem,Torricelli took a sealed tube approximatelyfour feet in length and filled it with mercury.

9

Gases

100 Gases

He inverted the tube in a pool of mercury.Torricelli noticed that the mercury in thetube fell, creating a vacuum at the top of the tube. The final height of the mercury in the tube was roughly 28 inches. Becausemercury has 13.5 times the density of water,Torricelli proposed that water should riseabout 13.5 � 28 inches or about 32 feet.Torricelli reasoned that atmospheric pres-sure placed a limit on how high a liquidcould raise in the tube. His inverted mercury-filled tube was the first barometer(Figure 9.1).

Many units are used to express pressure.Because pressure is defined as force per unitarea, a common unit used in the UnitedStates is pounds per square inch. This unitis commonly used for tire inflation pressure.The atmospheric pressure at sea level isabout 14.7 pounds per square inch. In themetric system, the basic unit for force is thenewton, abbreviated N, and area is mea-

sured in squared meters, m2. Therefore, themetric unit for pressure is newtons persquared meter, N/m2. A N/m2 is also knownas a pascal, abbreviated Pa. As we haveseen, atmospheric pressure can also beexpressed by measuring the height of mer-cury in a barometer. This measurement isgiven in inches in the United States and ishow atmospheric pressure is reported indaily weather reports. In the metric system,the height of mercury is given in millime-ters of mercury. A common unit used bychemists to express pressure is atmospheres.One atmosphere is equal to normal atmo-spheric pressure at sea level. The variousunits for pressure and the value of standardatmospheric pressure are summarized inTable 9.2.

The values in column three of Table 9.2are atmospheric pressure at sea level. At thetop of Mt. Everest, the atmospheric pres-sure is only one-fourth of that at sea level.

Table 9.1Composition of Earth’s Atmosphere by Volume

Figure 9.1In a mercury thermometer, atmosphericpressure, P, forces mercury to rise in thetube to a height of approximately 30 inches

Gases 101

The pressure in Denver is about three-fourths of that at sea level. While changesin elevation have pronounced effects onatmospheric pressure, temperature andmoisture also affect atmospheric pressure.Changes in the moisture and temperature ofair masses produce high and low pressuresystems that move across the continent dic-tating our weather. Pressure increasesunderwater because the water pressure addsto the atmospheric pressure. Ten meters (33feet) of water is equivalent to 1 atmosphereof pressure.

We can understand how atmosphericpressure is just the weight of the atmospherepushing down on the Earth’s surface, buthow can we apply the basic definition ofpressure to a confined gas? To expand ourconcept of pressure and provide a basicframework for understanding the behaviorof gases, we use a simple model for a con-fined gas. This model is known as thekinetic molecular theory. The kineticmolecular theory states:

1. A gas consists of small particles, eitherindividual atoms or molecules, movingaround randomly.

2. The total volume of the gas particles is sosmall compared to the total volume the gasoccupies that we can consider the total par-ticle volume to be zero. This means that agas consists almost entirely of emptyspace.

3. The gas particles act independently of oneanother. A particle is not attracted to norrepelled from any other particle.

4. Collision between gas particles andbetween gas particles and the walls of thecontainer are elastic. This means that thetotal kinetic energy of the gas particles isconstant as long as the temperature is con-stant.

5. The average kinetic energy of the gas par-ticles is directly proportional to the abso-lute temperature of the gas.

The kinetic molecular theory is usedthroughout the discussion of gases andshould become clearer as examples are usedthat illustrate this theory.

First, the kinetic molecular theory isused to explain pressure. We can use asealed syringe as our container and assumeit contains a volume of air. To simplify ourdiscussion, we will also assume air consistsof 80% nitrogen and 20% oxygen (Figure9.2). The nitrogen and oxygen moleculesmove randomly in the barrel of the syringe.The molecules collide with the walls of thesyringe barrel and the face of the syringe’splunger (as well as with each other). Thecollisions exert a constant force on theinside surface area of the walls of the syr-inge barrel. The force exerted per unit areaof the syringe’s internal surface is the pres-sure of the gas. Because the molecules moverandomly throughout the syringe barrel, thepressure is the same throughout the syringe.If we assume the pressure outside thesyringe is 1 atmosphere and that the plunger

Table 9.2Pressure Units and Values for StandardAtmospheric Pressure

Figure 9.2Syringe filled with air that is assumed to be80% nitrogen and 20% oxygen

102 Gases

is not moving, then the pressure inside thesyringe must also be 1 atmosphere. So thepressure of a confined gas is nothing morethan the force caused by the constant bom-bardment of the gas particles on the sides ofthe container.

The Gas LawsContinuing to use a syringe as a con-

tainer, the basic gas laws can be explained.These laws apply to what is referred to as anideal or perfect gas. An ideal or perfect gascan be thought of as a gas that conforms tothe kinetic molecular theory. In reality, gasmolecules do have volume and exert forceson each other. Under normal conditions oftemperature and pressure, though, thekinetic molecular theory explains the behav-ior of gases quite well. It is only when a gasis at very low temperatures and/or underextremely high pressure that a gas no longerbehaves ideally.

The definition of pressure assumed thevolume of the air contained in the syringewas constant. What would happen if theplunger is pushed while making sure that theopening remained sealed (Figure 9.3)?Pushing in on the plunger obviouslydecreases the volume of the syringe’s barrel.Because the volume has decreased, theinside surface area has also decreased. Thismeans that the frequency of collisions per

unit area has increased, which translates intoan increase in pressure.

Remember, because pressure is forcedivided by area, if the area decreases thepressure increases. When we push in on theplunger, the pressure increases to somevalue above 1 atmosphere. The smaller thesyringe’s barrel volume becomes, the higherthe pressure exerted by the air in syringe.The relationship between volume and pres-sure is known as Boyle’s Law. Boyle’s Lawwas first mentioned in Chapter 3 and isnamed after Robert Boyle. Simply stated,Boyle’s Law says that the pressure and vol-ume of an ideal gas are inversely related, asone goes up, the other goes down. Boyle’sLaw can be stated mathematically as

Pressure � Volume � constantOr

PV � k

So if the gas in the syringe was originally at1 atmosphere and the volume was 50 mL,then PV would equal 50 atm-mL. If wepushed in on the plunger to decrease the vol-ume to 25 mL, then the pressure would haveto increase to 2 atmospheres for PV toremain constant. Remember, when we applyBoyle’s Law, the only two variables thatchange are pressure and volume. We areassuming the temperature of the gas and thenumber of molecules of gas in the syringeremain constant.

What happens when the pressure of agas remains constant and the temperatureand volume change? Before looking at therelationship between temperature and vol-ume, though, the concept of temperaturemust be understood. Temperature is one ofthose terms that is continually used, butrarely given much thought. As long as wecan remember, we have had our temperaturetaken, seen and heard the daily temperaturereported, and baked foods at various tem-peratures. Intuitively, we think of tempera-

Figure 9.3Pushing in on the plunger decreases the vol-ume, which causes an increase in pressuredue to the collision frequency increasing

Gases 103

ture in terms of hot and cold, but what doestemperature actually measure? Temperatureis a measure of the random motion of theparticles making up a substance. Specifi-cally, temperature is a measure of the kineticenergy of the particles in a substance.Kinetic energy is energy of motion. A bodypossesses kinetic energy when it moves. Theformula for kinetic energy is

Kinetic Energy � 1⁄2

� mass � velocity2

K.E. � 1⁄2

mv2

Kinetic energy depends on the mass andvelocity of an object. Each of the nitrogenand oxygen molecules in the syringe containa certain amount of kinetic energy by virtueof the fact that they have mass and are mov-ing. If we could somehow measure thespeed of particles in the syringe, we woulddiscover that they move at various veloci-ties. While most move at near some averagevelocity, some are moving much slower andothers much faster than average. A graph ofthe percent of molecules at a certain veloc-ity versus the velocity is called a Maxwell-Boltzmann distribution, named for LudwigBoltzmann (1844–1906) and James ClerkMaxwell (1831–1879) (Figure 9.4). Avelocity of 400 meters per second is roughly800 miles per hour and represents the typi-

cal velocity gas molecules would have atroom temperature. The shape of Figure 9.5is dependent on temperature. At a highertemperature, a greater percentage of themolecules would be moving with greatervelocities, and at a lower temperature, agreater percentage with slower velocities(Figure 9.5).

Higher and lower velocities translateinto higher and lower kinetic energy of themolecules, respectively. We can now con-sider what a change in temperature means atthe molecular level. An increase in temper-ature implies more energetic particles and adecrease in temperature less energetic par-ticles. When a mercury thermometer isplaced in a hot oven, it is exposed to morecollisions with greater kinetic energies. Thekinetic energy of the gas particles is trans-ferred to the kinetic energy of the mercury,which causes it to expand and register ahigher temperature. The expansion of themercury can be compared to a rack of bil-liard balls being struck by the cue ball. Atthe lower temperature, the column of mer-cury in the thermometer is like the rackedballs. An energetic cue ball is like the ener-getic gas molecules in the oven. The cue balltransfers its kinetic energy to the rackedballs causing them to spread out. Similarly,the energetic gas molecules in the ovenimpart their energy to the mercury atoms

Figure 9.4Distribution of Speeds in Gas—Maxwell-Boltzmann Distribution (Rae Déjur)

Figure 9.5Boltzmann Curves at Higher and LowerTemperatures (Rae Déjur)

104 Gases

causing them to expand up the thermometercolumn. Cooling would cause the mercury’svolume to decrease. If the mercury wascooled to its freezing point of –39°C, theformation of solid mercury might be con-sidered analogous to racking up the billiardballs.

The two most common scales used tomeasure temperature are the Fahrenheit andCelsius scales. The German Daniel Fahren-heit (1686–1736) proposed his temperaturescale in 1714 in Holland. Fahrenheitinvented the modern mercury thermometerand calibrated his thermometer using threedifferent temperature standards. He assigneda value of 0 to the lowest temperature hecould obtain using a mixture of ice, salt, andliquid water. A value of 30 was used for amixture of ice and freshwater. The third pointwas set at 96 based on the oral temperatureof a healthy person. Using his scale Fahren-heit determined the boiling point of waterwould be 212. He later changed the standardof the ice-freshwater mixture from 30 to 32in order to have an even 180 temperaturedivisions between the freezing and boilingpoints of water. The latter change resulted in98.6°F being the accepted normal body tem-perature of a healthy person.

Anders Celsius (1701–1744) from Swe-den devised his temperature scale in 1742.Celsius assigned a value of 0 to the boilingpoint of water and 100 to the temperature ofthawing ice. Instrument makers soonreversed the 0 and 100 to give us the mod-ern freezing and boiling points of water as0°C and 100°C, respectively. The relation-ship between the Fahrenheit and Celsiustemperature scales are given by the twoequations:

Degrees Celsius �5/9 � (Degrees Fahrenheit � 32)

Degrees Fahrenheit �(9/5 � Degrees Celsius) � 32

Both the Fahrenheit and Celsius tem-perature scales are based on the physicalcharacteristics of water. The use of water toestablish a zero point means that both 0°Fand 0°C are arbitrary and not true zeropoints. Because temperature is a measure ofthe random motion of the particles makingup a substance, true zero on either theFahrenheit or Celsius scale would imply thatthere is no motion at 0°F and 0°C. We knowthat substances at these temperatures parti-cles possess an ample amount of motion andkinetic energy. A temperature scale that hasa true zero is known as an absolute temper-ature scale. One problem that arises whenusing the Fahrenheit and Celsius tempera-ture scales is because they are not absolutescales, the mathematical comparisons do notmake sense. For example, if we compareequal quantities of water at 10°C and 20°C,we might expect the water at 20°C to betwice as hot or have twice as much energyas the water at 10°C. This is not the case. Asan analogy, consider a child who is 3 feet talland an adult who is 6 feet tall. Our absolutescale for measuring a person’s height is thevertical distance from the bottom of ourfeet, and using this absolute scale the adultis twice the height of the child. But what ifboth the child and adult stood on top of 5-foot ladders and we measured their heightsfrom the ground? The child would now mea-sure 8 feet and the adult 11. The adult onthis revised scale is no longer twice as tallas the child. The Fahrenheit and Celsiustemperature scales are like measuringheights of people standing on ladders. Itmight sound ridiculous to measure peoplestanding on ladders, but the heights wouldseem normal if that’s how we have alwaysmeasured height.

The most common absolute tempera-ture scale used by scientists is the Kelvintemperature scale. The Kelvin temperaturescale was proposed by William Thompson,

Gases 105

Lord Kelvin (1824–1907). The Kelvin tem-perature scale has an absolute zero. Truecomparisons can be made using the Kelvinscale. A substance at a temperature of 400Kelvins contains particles with twice asmuch kinetic energy as a substance at 200Kelvins. Absolute zero is the temperaturewhere the random motion of particles in asubstance stops. It is the absence of temp-erature. Absolute zero is equivalent to�273.16°C. How this value is determinedis discussed shortly after we discuss our nextgas law. The relationship between Kelvinand Celsius temperature is

Kelvins � Degrees Celsius � 273.16

For most work, it is generally sufficient toround off 273.16 to 273.

Now that the physical meaning of tem-perature has been explored, the relationshipbetween temperature and volume of an idealgas can be examined. Let’s consider asyringe containing a specific volume of airat a specific temperature. The syringe isoriginally at room temperature, about 20°C,and placed in a pot of boiling water at100°C. In the boiling water, energy is trans-ferred to the nitrogen and oxygen moleculesin the syringe barrel, and they move faster.The number of collisions with the interiorwalls of the syringe’s barrel increases. If thepressure within the syringe is to remain con-stant, the volume must increase. Anotherway of putting this is that an increase in tem-perature causes the pressure within thesyringe to increase. Because the pressureoutside the syringe is 1 atmosphere, the gaswill push out on the plunger until the pres-sure in the syringe returns to 1 atmosphere.There is a direct relationship between tem-perature and volume of an ideal gas. As tem-perature increases, volume increases. Thedirect relationship between temperature andvolume in an ideal gas is known as Charles’law. Jacques Alexandre Charles (1746–

1823) was an avid balloonist. Charles madethe first hydrogen-filled balloon flight in1783 and formulated the law that bears hisname in conjunction with his ballooningresearch.

Charles’ law can be stated mathemati-cally as

When applying Charles’ Law it is importantto remember that the absolute temperaturemust be used. Let’s use the syringe exampleto determine how much the volume wouldincrease if the temperature was raised from20°C to 100°C. Let’s assume the originalvolume at 20°C is 25 mL. First we’ll con-vert the temperatures to Kelvins, and thenapply Charles’ Law. The calculations are

Solving for V, we find that the volume at100°C is approximately 32 mL.

Louis Gay-Lussac continued the bal-looning exploits initiated by Charles, as-cending to over 20,000 feet in a hydrogenballoon in the early 1800s. Gay-Lussac’slaw defines the relationship between thepressure and temperature of an ideal gas. Ifthe temperature of the air in the syringeincreases while keeping the volume con-stant, the gas particles speed up and makemore collisions with the inside walls of thesyringe barrel. As we have seen, an increasefrequency in the number of collisions of the gas particles with a container’s walltranslates into an increase in pressure. Gay-Lussac’s law says that pressure is directly

Volume� Constant

Temperature

V� k

T

20�C � 273 � 293 Kelvins

100�C � 273 � 373 Kelvins

25 mL V�

293 Kelvins 373 Kelvins

106 Gases

proportional to temperature. Gay-Lussac’slaw expressed mathematically is

Again, the temperature must be expressed inKelvins. If the pressure and temperature ofthe air in the syringe are originally 1.0 atmo-sphere and 293 Kelvins and the syringe isplaced in boiling water, the pressure willincrease to approximately 1.3 atmospheres.The calculations are

The relationship between temperatureand pressure provides a method for deter-mining the value for absolute zero. By mea-suring the pressure of a gas sealed in aconstant volume container at different tem-peratures and extrapolating to a pressure of0 atmospheres gives the value for absolutezero (Figure 9.6).

Boyle’s, Charles’, and Gay-Lussac’slaws explain the relationships between pres-sure, volume, and temperature of an idealgas. In the examples thus far, the amount ofgas in the syringe was considered to be con-stant. The amount of gas is measured inmoles (the standard symbol for moles is n).In Chapter 4, we learned that Avogadro’shypothesis stated that equal volumes ofgases contain an equal number of mole-cules. This means that the volume of a gas isdirectly proportional to the number ofmolecules present, and therefore, the num-ber of moles of gas. This relationship isknown as Avogadro’s law. If temperatureand pressure are held constant and somehowmore molecules are added to the syringebarrel, the volume would increase. A betterexample to illustrate Avogadro’s Law is tothink about blowing up a balloon. As youblow up the balloon, you are adding moregas molecules (principally carbon dioxide)and the volume increases. It is assumed thepressure and temperature stay constant asthe balloon inflates, and that the increase involume is due to the increase in the numberof moles of gas entering the ballon. Becausethe relationship between pressure and molesis direct, the relationship is expressed math-ematically as

The four fundamental gas laws are sum-marized in Table 9.3.

The Ideal Gas LawThe gas laws discussed in the previous

section are limited, because they only allowus to examine the relationship between twovariables at a time. Fortunately, all four lawscan be combined into one general law called

Pressure� Constant

Temperature

P� k

T

1.0 atmosphere Pressure Final�

293 Kelvins 373 Kelvins

Pressure Final � 1.0 atmospheres373 Kelvins

� � 1.3 atmospheres293 Kelvins

Figure 9.6The value of absolute zero may bedetermined by extrapolating a plot ofpressure versus temperature to a value of 0atmosphere. This gives a value for absolutezero of �273.16°C

Volume� Constant

moles

V� k

n

Gases 107

the ideal gas law. The ideal gas law relatesthe four quantities pressure, volume, moles,and temperature. The ideal gas law is givenby the equation

Pressure � Volume � Moles � R � Temperature

PV � nRT

In this equation, R is called the ideal gas lawconstant. Its value depends on the unitsused, but assuming pressure is measured inatmospheres, volume in liters, and tempera-ture in Kelvins its value is 0.082 atm-L/mol-K. Other forms of the ideal gas law are

The four gas laws in the previous section areall special cases of the ideal gas law. We canuse the ideal gas law to calculate the volumeone mole of gas occupies at standard condi-tions. Standard conditions are 0°C (273 K)and 1 atm pressure. The volume at theseconditions is known as a standard molar vol-ume. Plugging the numbers into the idealgas law equation gives a value of 22.4 litersfor the standard molar volume.

Partial Pressure and VaporPressure

John Dalton was a contemporary ofCharles and Gay-Lussac. As a meteorolo-gist, he had a keen interest in the atmo-sphere and devoted much of his work to a

study of the behavior of gases. Dalton’sstudy of gases led him to formulate the lawof partial pressure and the concept of vaporpressure. Dalton’s Law of Partial Pressurestates that in a mixture of gases each gasexerts a pressure independent of the pres-sure exerted by other gases in the mixture.The pressure of the individual gases in themixture is the partial pressure of that gas.The sum of the partial pressures is the totalpressure. Figure 9.7 depicts the concept ofpartial pressure.

Partial pressure is directly related to themoles of gas present. The partial pressuresof gases in our atmosphere are approxi-mately 0.78 atm for nitrogen, 0.21 atm foroxygen, and 0.01 atm for all the other gasescombined.

At the same time that Dalton proposedhis ideas on partial pressure, he developedthe concept of vapor pressure. A vapor isthe gaseous form of a substance that nor-mally exists as a solid or liquid. A gas is asubstance that exists in the gaseous statesunder normal conditions of temperature andpressure. The vapor pressure of a liquid isthe partial pressure of the liquid’s vapor atequilibrium. Liquids with strong inter-molecular forces exert lower vapor pressuresthan those with weak intermolecular forces.In liquids with strong intermolecular forces,it is more difficult for the molecules to leavethe liquid state and enter the gaseous state.

Table 9.3Summary of Four Laws for an Ideal Gas

PV P V P V1 1 2 2� R or �

nT n T n T1 1 2 2

Figure 9.7Dalton’s Laws of Partial Pressure states thatin a mixture of gases each gas exerts a pres-sure independent of the other gases. Thetotal pressure is the sum of the partial pres-sures

108 Gases

Liquids with high vapor pressures areknown as volatile liquids.

To understand vapor pressure, let’s con-sider an empty jar that is partially filled withwater and then covered with a lid. We willassume the space above the water in the jarcontains only air when we screw on the jar’slid. After the lid is place on the jar, watermolecules leave the liquid and enter the airabove the liquid’s surface. This process isknown as vaporization. As time goes by,more water molecules fill the air spaceabove the liquid, but at the same time, somegaseous water molecules condense backinto the liquid state. Eventually, a point isreached where the amount of water vaporabove the liquid remains constant. At thispoint, the rates of vaporization and conden-sation are equal, and equilibrium is reached.The partial pressure exerted by the water atthis point is known as the equilibrium vaporpressure or just vapor pressure. Vapor pres-sure is directly related to the temperature,that is, the higher the temperature, thehigher the vapor pressure. Table 9.4 gives

the vapor pressure of water at several tem-peratures. The amount of water vapor in theair is known as humidity. Relative humid-ity is the ratio of the actual vapor pressurecompared to the equilibrium vapor pressureat a particular temperature. If the vaporpressure was 14.0 at a temperature of 20°C,then the relative humidity would be14.0/17.5 � 80%.

The vapor pressure of a liquid dictateswhen a substance will boil. In fact, the boil-ing point of a substance is defined as thetemperature at which the vapor pressureequals the external pressure. Typically, theexternal pressure is equal to atmosphericpressure, and we define the normal boilingpoint as the temperature when the vaporpressure equals 1 atmosphere. If we con-sider water heated on a stove, the bubblesthat develop in the liquid contain watervapor that exerts a pressure at the specificvapor pressure of water at that temperature.For example, when water reaches 60°C, anybubbles that form will contain vapor at 149mm Hg (see Table 9.4). At this pressure, andany other pressure below 760 mm Hg (1atmosphere), the external pressure of 1atmosphere causes the bubbles to immedi-ately collapse. As the temperature of thewater rises, the vapor pressure continuallyincreases. At 100°C, the vapor pressureinside the bubbles finally reaches 760 mmHg. The vapor pressure is now sufficient toallow the bubbles to rise to the surface with-out collapsing. At higher elevations wherethe external pressure is lower, liquids boil ata lower temperature. At the top of a 15,000-foot peak, water boils at approximately85°C rather than 100°C. This increases thecooking time for items, as noted in the di-rections of many packaged food. If theexternal pressure is increased, the boilingtemperature also increases. This is the con-cept behind a pressure cooker. The sealedcooker allows pressure to build up inside it

Table 9.4Vapor Pressure of Water at DifferentTemperatures

Gases 109

and increases the boiling point allowing asubstantial increase in cooking temperatureand decrease in cooking time.

Applications of the GasLaws

The pressure cooker is just one practi-cal example of the behavior of gases ineveryday life. In this section, we look at afew more relevant examples involvinggases. A process analogous to the pressurecooker which we all can relate to is poppingcorn. Each kernel of corn contains about15% water. The hard kernel acts as a con-stant volume container. When the kernel isheated, the temperature increases, andaccording to Gay-Lussac’s law, the pressureincreases. A small amount of the waterinside the kernel vaporizes into superheatedsteam, but most of the water remains in liq-uid form because of the increased pressure.At some point when the pressure increasesto several atmospheres, the kernel pops asthe steam transforms the starchy kernel intogelatinous globules of popped corn.

One national news story in the year2000 involved the defect in certain brandsof tires and how tread separation causedvehicle accidents leading to personal injuryand in some cases deaths. One prominentaspect of the tire story involved whether thetires were properly inflated to the correctpressure. The air inside a tire can be consid-ered an ideal gas, and we can apply the gaslaws to the air inside. The inflation pressurerecommendations stamped on the sides oftires call for tires to be inflated under coldconditions, or before the vehicle is driven.Because the most important aspect of tireperformance is correct inflation pressure, itis important to adhere to the recommendedtire pressure. Many individuals disregardthis recommendation. It is not uncommonfor someone to pull into a gas station and fill

a warm tire to the recommended pressure.How much can this affect the pressure in thetire? In this example, the important variablesare pressure and temperature. We willassume that the amount of air in the tire andvolume of the tire are constant, and will seehow much of an error can be introduced byignoring the inflation temperature. Let’sassume the recommended inflation pressureis 30 pounds per square inch (psi), and wefill the tire on a hot day after pulling off thefreeway and the tire temperature is 90°F(32°C). We will assume the cold pressureapplies to a temperature of 70°F (21°C).Gay-Lussac’s law can be applied to deter-mine the actual cold inflation pressure.

We can write Gay-Lussac’s Law as

This is a relatively small change, but con-sider the change if you inflate the tire in aheated garage at 70°F and then drive in win-ter when the temperature is 10°F. In thiscase, the pressure drop would be approxi-mately 5 psi. A pressure change of thismuch can lead to as much as a 25% reduc-tion in fuel efficiency. Additionally, under-inflated tires result in overloading andpossible tire failure. Reduced tire pressuresof as little as 4 psi were cited in some of thenational stories involving tire failure andvehicle accidents.

Another area in which the gas laws playa key role is in scuba diving. At the surface,we breath air at a pressure of approximately1 atmosphere. The partial pressures ofnitrogen and oxygen are 0.78 and 0.20atmosphere, respectively. A scuba diverbreaths compressed air that is delivered ata pressure that corresponds to the pressureat the depth of the diver. Because 33 feet of

P Pcold hot�

T Tcold hot

294 KP � 30 psi � � 28.9 psicold 305 K

110 Gases

water corresponds to 1 atmosphere of pres-sure, a diver at 33 feet would breath air at atotal pressure of 2 atmospheres; one atmo-sphere of this pressure is due to the air pres-sure at sea level and 1 atmosphere is due tothe pressure of the water. The partial pres-sures of the nitrogen and oxygen would betwice as great as that at the surface. Mostserious problems in diving arise whendivers descend to depths in excess of 100feet. The total pressure at depths greaterthan 100 feet can cause problems such asnitrogen narcosis and oxygen toxicity.Nitrogen narcosis, also known as rapture ofthe deep, is due to breathing nitrogen underpressure. At a depth of 100 feet, the totalpressure of air delivered to the diver isabout 4 atmospheres. This results from awater pressure of slightly greater than 3atmospheres plus the surface air pressure of1 atmosphere. The partial pressure of nitro-gen at this depth would be about 3.2 atmo-spheres. Nitrogen at this pressure producesan anesthetizing effect similar to nitrousoxide (laughing gas). The result is a state ofnitrogen intoxication with loss of judgment,a state of euphoria, drowsiness, andimpaired judgment. If nitrogen narcosisdevelops, it can be quickly reversed bymoving to a shallower depth. To preventnitrogen narcosis, deep divers use a mixtureof helium, oxygen, and nitrogen. Helium ismuch less soluble in body tissues and likenitrogen is inert and does not play a role inmetabolism.

Another example of a diving problemthat is a direct consequence of Dalton’s Lawof partial pressure concerns oxygen toxicity.The deeper a diver descends, the greater thepartial pressure of oxygen. At a depth of 130feet, the total pressure is close to 5 atmo-spheres and the partial pressure of oxygenwill be close to 1 atmosphere (21% � 5atmospheres). What this means is thatbreathing compressed air at 130 feet is like

breathing pure oxygen at the surface.Breathing pure oxygen or breathing com-pressed air at 130 feet is only a problem ifdone for several hours. Oxygen toxicity candamage the lungs, affect the nervous sys-tem, and in extreme cases, result in seizures.Recreational divers do not stay down longenough at great depths for this to become aproblem, but it can be a problem for com-mercial divers. When the partial pressure ofoxygen approaches 2 atmospheres, it maytake only several minutes for symptoms ofoxygen toxicity to appear. To prevent oxy-gen toxicity, divers use a lower percentageof oxygen for deep dives. For example, adiver at 200 feet might breath from a tankthat contains only 4% oxygen.

Both nitrogen narcosis and oxygen tox-icity involve deep diving. Other problemsresult when the diver ascends too quickly.We know from Boyle’s Law that as pressuredecreases, volume increases. Decompres-sion sickness develops when a diver ascendstoo fast and nitrogen bubbles develop in thetissue. (This is also related to Henry’s lawand the solubility of gases in the blood; thistopic is covered in Chapter 11.) Symptomsmay include dizziness, paralysis, shock, andjoint and limb pain. The term “the bends”comes from the pain that appears in bodyjoints such as the elbows, knees, ankles, andshoulders. During a normal ascent frommoderately deep depths (greater than 60feet), a diver must be sure to take scheduledstops and allow gases such as nitrogen tocome out of the body tissue slowly. Anotherserious problem with surfacing too quicklyis an air embolism. An air embolism devel-ops when air expands in the lungs and bub-bles are forced into the blood vessels,possibly cutting off blood flow to the brainand resulting in a loss of consciousness.

Both decompression sickness and airembolism can be treated by placing the vic-tim in a hyperbaric chamber. “Hyper”

Gases 111

means greater or above normal and “baric”means weight or pressure. In a hyperbaricchamber, the victim is placed in an envi-ronment of increased pressure and enrichedoxygen. By decreasing the pressure slowly,the gases can come out slowly. Theenriched oxygen in the hyperbaric chamberhelps to purge and dilute the nitrogen in thetissues.

A final example of a practical applica-tion of the ideal gas law involves snowmaking. Many ski resorts and Hollywoodoften employ snow guns to produce theirown snow (Figure 9.8). A snow gun essen-tially uses compressed air to atomize waterinto droplets. The droplets are sprayed intothe air where they condense into ice crys-tals. As the compressed air and waterdroplets leave the snow gun, the decrease inpressure causes the air to expand resultingin a drop in temperature. The cooling of airas it expands is also demonstrated whenmoist air masses from the Pacific moveinland over the West Coast. The air masseswill ride up the western edge of coastalmountain ranges such as the Cascades. Asthis moisture-laden air moves to higher ele-vations, the pressure decreases causing theair to expand and cool. When the relativehumidity reaches 100%, precipitationresults in the form of rain or snow. On theeastern side of the mountains, dry airmasses descend and become compressed asthey are subjected to increased pressure atlower elevations. This results in dry, desert-

like conditions on the eastern side of thecoastal mountain ranges.

The few examples illustrated abovedemonstrate how gases and the gas lawsaffect our lives on a daily basis. We are sur-rounded by gas and our lives are dependenton a steady supply of gas. Humans can sur-vive a few weeks without food, a few dayswithout water, but only a few minutes with-out a steady supply of air. In this chapter, wehave focused on some of the unique proper-ties of gases and the gas laws. We continueour exploration of gases when we examinephase changes in Chapter 10 and solubilityin Chapter 11.

Figure 9.8Snow Gun (Rae Déjur)

IntroductionChapters 8 and 9 explored the three com-mon states (phases) of matter—solid, liquid,and gas. Substances are generally associatedwith one of these three phases under normalconditions of 1 atmosphere pressure androom temperature; yet phase changes con-stantly occur all around us. Take an ice cubeout of the freezer, and in a few seconds itwill start to melt changing from the solid tothe liquid phase. The familiar phase changesof water can be used to explore changes ofstate and also introduce the topic of ther-mochemistry. Thermochemistry is thestudy of the energy changes associated withchemical processes. The term “chemicalprocess” will be used throughout this chap-ter, and it includes both physical processesand chemical reactions. After looking at thethermochemistry of phase changes, chemi-cal processes in general will be examined.

Changes of StateThe phase changes of water result in the

transitions between the solid, liquid, and gasstates:

H2O(s) H2O(l) H2O(g) (1)

The double arrows indicate that two pro-cesses are possible at each stage. The termsfreezing and melting indicate the generaldirection of the phase change. The formerfrom liquid to solid, and the latter fromsolid to liquid. Similarly, condensation andvaporization characterize a phase changefrom gas to liquid and from liquid to gas,respectively. Equation 1 is somewhat mis-leading because it implies the liquid phaseis always intermediate between the solidand gas phases. A solid can pass directlyinto the gas phase without passing throughthe liquid phase. This process is known assublimation. Solid carbon dioxide, alsoknown as dry ice, is a substance that sub-limes under normal conditions. The changefrom a gas directly to a solid is called depo-sition. Table 10.1 summarizes the differentphase changes.

Energy and Phase ChangesBefore looking at the energy associated

with phase changes, some basic termsrelated to thermochemistry need to be

10

Phase Changes and Thermochemistry

114 Phase Changes and Thermochemistry

defined. Terms, such as “work,” “heat,” and“energy” are used every day, but they havevery specific meaning in the context ofchemistry. Whenever a chemical processsuch as a phase change or chemical reactionis examined, it is important to specifyexactly what is being examined. This state-ment might seem obvious, but it is espe-cially important when considering theenergy aspects of a chemical process. Thechemical system or simply system is thatpart of the universe a person is interested inand wants to examine. The concept of a sys-tem is a somewhat abstract concept. Attimes, it is very easy to define a system,while in other cases it might not be obviouswhat the system includes. For example,when you take an ice cube out of the freezerand observe how it changes, you can defineyour system as the ice cube. If you are look-ing at the reaction between vinegar and bak-ing soda, the system might be a beakercontaining these two substances. A systemmight be microscopic, or it might be as largeas the planet. For example, a biochemistlooking at the effect of toxins on humansmight define a system as the human cell;whereas an atmospheric chemist looking athow the chemical composition of the atmo-sphere affects global temperatures mightdefine the system as the entire planet and its

atmosphere. Everything outside the systemis considered the surroundings. This meansthe entire universe excluding the definedsystem comprises the surroundings. A use-ful conceptual formula is

System � Surroundings � Universe

Energy is defined as the ability to dowork or produce heat. In Chapter 9, kineticenergy was defined as the energy of motion.The motion of molecules and the electronswithin the atoms make up the kinetic energyof a chemical system. Potential energy canbe considered stored energy. Potentialenergy results from the position or configu-ration of a system. Just as the gravitationalpotential energy of an object is due to itsposition with respect to the Earth’s surface,chemical potential energy results from theposition of electrons with respect to nuclei.Chemical potential energy is stored in thebonds of compounds. During a chemicalreaction, atoms are rearranged, and it is thisreconfiguration of the atoms that causes theenergy changes accompanying the reaction.The sum total of the kinetic and potentialenergy of all particles in the chemical sys-tem is termed the internal energy.

Inherent in the definition of energy arethe terms “work” and “heat.” Work isdefined as force times distance. In physics,work is generally calculated by multiplyingthe resultant force applied to an object byhow far the object moves. The physical definition of work can be modified into a chemical definition for our purposes.Chemical work can be considered expan-sion work. To understand chemical work,consider a chemical reaction that takesplace in a container equipped with a fric-tionless piston (Figure 10.1). Let’s assumea gas is generated during the chemical reac-tion, for example alka-seltzer plus water.The gas generated in the reaction expandsin the container and forces the plunger out.

Table 10.1Phase Changes

Phase Changes and Thermochemistry 115

The external pressure outside the containercan be assumed to be 1 atmosphere. It canbe shown that the work done during thechemical reaction is equal to the externalpressure times change in volume. This workis called PV work, and it is what is meantby chemical work.

Energy may also produce heat. Heat isdefined as a transfer of energy from a bodyof higher temperature to a body of lowertemperature. Energy and heat are often con-sidered as synonymous terms, but heatinvolves the flow of energy. When examin-ing thermochemistry, the transfer of energybetween the chemical system and its sur-roundings is of primary concern.

Work, heat, and energy can be ex-pressed using various units. The metric unitfor energy is the joule, abbreviated J. Thejoule was named in honor of the Englishphysicist James Prescott Joule (1818–1889).

The joule is a rather small unit of energy. Ittakes approximately 10,000 joules to raisethe temperature of a coffee cup of water by10°C. Often kilojoules (kJ) are used with 1 kJ equal to 1,000 J. Another popular unitis the calorie, abbreviated with a c. A calo-rie is the amount of heat necessary to raiseone gram of water by 1°C. One calorie isequal to 4.18 joules. The physical unit of acalorie should not be confused with a foodcalorie. A food calorie, symbolized with acapital C, is equivalent to 1,000 calories.When you eat a piece of cake with 400Calories, you have actually consumed400,000 calories.

Heating CurvesPhase changes are most often the result

of a heat added or released from a sub-stance. Phase changes are also pressurerelated, but the thermal aspects of phasechanges are the primary focus in this sec-tion. The overall transition between phasesof a substance can be graphically displayedusing a heating curve. A heating curved isa graph showing the temperature of the sub-stance versus the amount of heat absorbedover time. Note that it could have beencalled a cooling curve and the transitionsexamined in terms of heat released overtime. Whether called a heating curve orcooling curve, the graph would be exactlythe same.

Figure 10.2 shows the heating curve forwater. To understand Figure 10.2, the pathof an ice cube on the curve is followed asheat is supplied to it. To make our examplerealistic, assume the ice cube has a mass of10 grams and is taken out of a freezer whereits temperature is �10°C. Consider whathappens when the cube is placed in a pan ona stove and heat applied. Initially, the icecube is at �10°C when it is removed fromthe freezer. The temperature of the ice cube

Figure 10.1A reaction generates a gas, which causes theplunger to expand. The volume increases byan amount equal to �V. The externalpressure is P. Because pressure � force(f)/area(A), then the force is equal to PA. Thedistance the plunger moves is d. Work �force � distance gives work � PAd. The fig-ure shows Ad is equal to �V. This meanswork is equal to P�V (Rae Déjur)


immediately starts to raise, but it does notstart to melt until its temperature reaches0°C. The amount of heat absorbed between�10°C and the melting point of 0°C can bedetermined using the specific heat capac-ity, or just specific heat, of ice. The specificheat of a substance is a measure of theamount of heat necessary to raise the tem-perature of one gram of a substance by 1°C.The specific heats of several common sub-stances are listed in Table 10.2.

Table 10.2 demonstrates that the spe-cific heat of a substance depends on itsphase. The specific heat of liquid water isapproximately twice that of ice and steam.Water has one of the highest specific heatscompared to other liquids. The high specificheat of liquid water is directly related to thepresence of hydrogen bonds. Because spe-cific heat is a measure of the heat needed tochange the temperature of a substance, andtemperature is a measure of the randomkinetic energy of the particles, substanceswith hydrogen bonds have a greater ten-dency to resist temperature change. Thehigh specific heat of water explains whycoastal environments have more moderateweather than areas at similar latitudeslocated inland. Water’s high heat capacity

means coastal regions will not heat up orcool down as much as areas not adjacent toa body of water.

The relationship between heat, specificheat, and temperature change is given by theequation Q � mc�T. In this equation, Q isthe amount of heat absorbed or released injoules, m is the mass in grams, c is the spe-cific heat of the substance, and �T is the

Figure 10.2The Heating Curve for Water

Table 10.2Specific Heats of Some Common Substances


change in temperature equal to the finaltemperature minus the initial temperature.For our ice cube, the amount of heat neededto increase the temperature from –10°C to0°C is equal to (10 g)(2.1 J/g°C)(10°C) �210 J. If the temperature of a substancedecreases, heat must be released from thesubstance. In this case, the change in tem-perature is negative resulting in a negativevalue for Q. A negative sign for Q alwaysindicates that heat is being removed fromthe substance rather than being added to it.

Once the ice cube reaches 0°C, thecube starts to melt. The heating curveplateaus at the melting point even thoughheat is still being added to the ice-liquidmixture. It may seem strange that eventhough heat is added the temperature doesnot increase, but this is because the addedheat is used to convert ice to liquid water.The heat supplied at the melting point is notbeing used to increase the random kineticenergy of the water molecules, but to over-come intermolecular attractions. Theenergy supplied at the melting point goesinto breaking the water molecules free fromthe crystalline structure resulting in theless-structured liquid state. As long as iceis present, any heat added goes into thephase change. The heat necessary to meltthe ice is called the heat of fusion forwater, and its value is 6.0 kJ per mole ofwater. Our 10-gram ice cube is equal toapproximately 0.55 moles of water. Thismeans it will require about 3.3 kJ or 3,300joules to melt the ice cube. The heat offusion of a substance is a measure of howmuch energy is required to convert a solidinto a liquid. It would be expected thatsolids, which are tightly held together,would have high heats of fusion. Sodiumchloride with its strong ionic bonds has aheat of fusion of 30.0 kJ per mole; alu-minum has a heat of fusion 10.7 kJ permole. Substances held together by weak

dispersion forces have low heats of fusion.The heat of fusion of oxygen is 0.45 kJ permole.

Once all the ice has melted, the heatadded to the liquid water can now go toincreasing the kinetic energy of the watermolecules and the temperature begins to riseagain. The rate at which the temperaturerises is governed by the specific heat of liq-uid water, 4.2 J/g°C. The specific heat of liq-uid water is twice that of ice; therefore, therate at which the temperature increases forliquid water is only half of what it was forice. The heat necessary to increase liquidwater from 0°C to 100°C can be found bymultiplying 10 g by 4.2 J/g°C by 100°C.This is equal to 4,200 joules. The heatingcurve reaches a second plateau at water’sboiling point of 100°C. At this temperature,any heat added to the water goes into break-ing the hydrogen bonds as liquid water isconverted to steam. Just as with the firstplateau at the melting point, the temperatureremains constant as long as the liquid andgas phases coexist. The wide plateau at theboiling point indicates that much moreenergy is needed for the vaporization pro-cess than for the melting process. The heatneeded to convert liquid water to steam iscalled the heat of vaporization of water.The heat of vaporization for water is 41 kJper mole. This value is roughly seven timesthe heat of fusion and indicates it takes only1/7 of the energy to melt ice as compared tovaporizing water. Vaporization should takesignificantly more energy than meltingbecause to convert liquid water to steamrequires completely breaking the hydrogenbonds. In the gas phase, the water moleculescan be thought of as independent moleculeswith minimal intermolecular attraction. Theheat of vaporization of a substance is gener-ally several times that of its heat of fusion.This is because the intermolecular forcespresent in the condensed phases must be


overcome before a substance can be con-verted to gas.

In closing the discussion of the heatingcurve for water, a couple of points should bemade. The last section of the heating curverepresents the situation when all the liquidwater has been converted to steam. At thispoint as the temperature of the steam beginsto rise, superheated steam would be ob-tained. In our ice cube example, we wouldnot observe this temperature rise becausethe water we heated in our pot (our system)would escape to the surroundings. It shouldalso be remembered that phase changes canbe considered in terms of a cooling curve.In this case, we would follow the curve fromright to left and 41 kJ per mole of heatwould have to be released to condense steamto liquid water and 6.0 kJ per mole wouldhave to be released by liquid water for it tofreeze (Figure 10.2).

Many everyday examples illustrate theconcepts embodied in the heating curve.Perspiration is a physiological response forcooling the body. For sweat to evaporatefrom our skin, heat must be supplied. Thisheat comes from our body. If we considerour body as a chemical system, sweating issimply a means of transferring heat fromour body to the surroundings. The sameevaporative cooling effect occurs when step-ping out of a shower or getting out of waterafter a swim. While the vaporizationrequires energy, a heating effect occurs dur-ing condensation and freezing. One practiceused in agriculture to protect plants fromfrost damage is to spray the plants withwater. Typically, growers using this methodspray their crops with water the eveningbefore a hard freeze is forecasted. Whentemperatures fall during the night and earlymorning, the water freezes releasing energyto the surrounding plants. This energy maybe sufficient to keep the plants themselvesfrom suffering frost damage. Constantly

occurring phase changes in the atmospherefigure prominently in our weather. Clouds,rain, snow, hail, and humidity all demon-strate phase changes on a grand scale. Asnoted in Chapter 9, ski resorts and the enter-tainment industry try to duplicate the snow-making process by manufacturing snow.Snow guns use a source of water and com-pressed air to spray water droplets into theair. The expansion of the air as it leaves thesnow gun creates a drop in the temperatureof the surrounding air helping the dropletscool and fall as snow. Depending on theambient temperature, the snow guns mayhave to be elevated and/or the water cooledto get water to undergo the phase changefrom liquid to solid.

CalorimetryIn the previous section, the thermo-

chemistry of water’s phase changes wasexamined. The heat of fusion and heat ofvaporization were used to determine theenergy flow that accompanies phase changes.All chemical processes involve energy. Tocharacterize the energy changes involved inthese processes, chemists use the term “heatof ” followed by the process. For example,heat of vaporization characterizes the energychanges involved in the vaporization process.Table 10.3 defines a number of “heat of ”processes chemists use.

A basic method for determining theenergy change involved with many chemi-cal processes is calorimetry. A calorimeteris an insulated container used to carry out achemical process. A thermometer is used tomeasure temperature changes that takeplace during the process. A simple constant-pressure calorimeter is shown in Figure10.3. This type of calorimeter derives itsname from the fact that it is open to theatmosphere and the pressure remains con-stant during the process. Constant-pressure


calorimeters are used to determine all theprocesses listed in Table 10.3 except com-bustion. The general procedure in using aconstant-pressure calorimeter is to carry outthe process in the calorimeter, observe thetemperature change that takes place, andthen use the equation Q � mc�T to deter-

mine the heat involved in the process. Theprocesses involved using constant-pressurecalorimeters generally involve solutions oremploy water. A description of a simpleexperiment that can be used to determinethe heat of fusion of water can illustrate theprocedure. To determine the heat of fusionof water, a measured volume of water isplaced in the calorimeter. The temperatureof the water is measured before placing ameasured mass of ice into the calorimeter.After the ice is placed in the calorimeter, themixture is gently stirred until all the ice hasmelted. At this point, the final temperatureof the water in the calorimeter is measured.The basic principle used to determine theheat of fusion assumes that heat required tomelt the ice comes from the water in thecalorimeter. The amount of heat that flowsfrom the water to the ice is found by usingthe change in temperature of the water, thespecific heat of water, and the mass of water.This heat is divided by the moles of ice(mass of ice/18 g per mole) to give an esti-mate of water’s heat of fusion.

Table 10.3Energy Changes of Some Chemical Processes

Figure 10.3Simple Calorimeter (Rae Déjur)


A second type of calorimeter is a con-stant-volume or bomb calorimeter (Figure10.4). This type of calorimeter is used todetermine the energy content of materials. Asthe name implies, the process takes place in aheavy walled metal “bomb” of constant vol-ume. The sample is placed in the bomb withample oxygen to sustain the combustion. Thebomb is surrounded by water. The sample isignited electrically, and the temperaturechange of the water is monitored. The energycontent of the sample is calculated in a man-ner similar to that used in constant-pressurecalorimetry. Bomb calorimetry is used exten-sively to determine the energy content of fuelsand foods. The energy of food is expressed infood calories. The relationship between foodcalories and joules is 1 food calorie equals4,180 joules. The conversions mean that acandy bar containing 200 food calories has anenergy equivalent of 837,000 joules.

Heat of ReactionThus far our discussion on thermo-

chemistry has focused on physical processes

such as phase changes. Bomb calorimetryinvolves a chemical reaction. In a chemicalreaction, substances are rearranged, bondsare broken, and bonds are formed. Thereconfiguration of atoms during a chemicalreaction involves changes in the potentialenergy of the electrons. A lower chemicalpotential energy and a release of energyresults when bonds are formed. Conversely,an energy input is required to break bonds.It is not surprising that practically all chem-ical reactions produce or absorb energy.

The combustion of natural gas, meth-ane, to produce heat can be summarized inthe following reaction:

CH4(g) � 2O2(g) CO2(g) � 2H2O(l) � 890 kJ

The energy released in this reaction is usedto heat homes, generate electricity, and cookfood. Figure 10.5 is a schematic diagram ofthe process. When heat is released during achemical reaction, the reaction is said to beexothermic. Another way of saying this isthat there is a heat transfer from the systemto the surroundings. Examples of exother-mic processes include combustion reac-tions, condensation, and freezing. Reactionsin which the potential energy of the products

Figure 10.4Bomb Calorimeter (Rae Déjur)

Figure 10.5Energy involved in the combustion ofmethane, CH4. In the combustion process,the chemical potential energy of thereactants is higher than the products. Thedifference in potential energy is the heatinvolved in the reaction


is greater than the potential energy of thereactants are called endothermic reactions.In an endothermic reaction, the heat flowsfrom the surroundings into the system.Melting and vaporization are endothermicprocesses.

Another means to describe the transferof energy into or out of a chemical systemis in terms of internal energy. It was previ-ously stated that internal energy was equalto the sum total of the potential and kineticenergy of all particles making up a chemi-cal system. It is impossible to measuredirectly the internal energy of a system, butthe first law of thermodynamics states thatthe change in internal energy of a system isequal to the heat plus work:

�E � q � w (3)

The first law of thermodynamics simplysays that energy cannot be created ordestroyed. With respect to a chemical sys-tem, the internal energy changes if energyflows into or out of the system as heat isapplied and/or if work is done on or by thesystem. The work referred to in this case isthe PV work defined earlier, and it simplymeans that the system expands or contracts.The first law of thermodynamics can bemodified for processes that take place underconstant pressure conditions. Because reac-tions are generally carried out in open sys-tems in which the pressure is constant, theseconditions are of greater interest than con-stant volume processes. Under constantpressure conditions Equation 3 can berewritten as

�E � qp – P�V (4)

Written as Equation 4, qp is the heatexchange at constant pressure and P�V isthe work done. It is important to understandthe significance of the mathematical signs inEquation 4, and how they relate to thechange in internal energy. When qp is posi-

tive, heat is transferred from the surround-ings into the system causing the internalenergy to increase. Pressure is always a pos-itive value, but �V (� means “change in,”so �V means change in volume) can be pos-itive or negative. When �V is positive, workis being done by the system on its sur-roundings. Think of a syringe in which areaction results in the production of gas thatexpands and forces the plunger out. In thiscase, �V is positive and this causes theinternal energy of the system to decrease.When the gas expands, some of the internalenergy of the system is being used to dowork on its surroundings. There are actualfour cases that describe how heat and workcan affect the internal energy of the system.These are summarized in Table 10.4.

Reactions under constant pressure con-ditions are so prevalent in chemistry that thequantity qp is defined as the change inenthalpy of a chemical system and is sym-bolized as �H. Substituting �H for qp inEquation 4 and rearranging gives �H � �E� P�V. The change in enthalpy is typicallyused to characterize a reaction as exother-mic or endothermic. When qp is positive,heat flows into the system increasing thepotential energy of the system, and thisresults in an endothermic reaction. A nega-tive qp signifies a heat flow out of the system decreasing the potential energy andgiving an exothermic reaction.

When writing a chemical reaction, �Hshould be included with the reactants andproducts; it is assumed that the reactiontakes place under the standard conditions of25°C and 1 atmosphere. Several simplereactions with their �Hs are

2H2(g) � O2(g) 2H2O(g) �H � –484 kJ

N2(g) � 3H2(g) 2NH3(g) �H � –93 kJ

6CO2(g) � 6H2O(l) C6H12O6 �6O2(g) �H � �2,816 kJ

The first two reactions given are for the for-mation of water and ammonia, respectively.


The �H values for these two reactions arenegative, indicating that these reactions areexothermic. In both of these reactions, theproducts are formed from their basic ele-ments so the �Hs represent heats of forma-tion. The third reaction is the familiarreaction for photosynthesis. The energyneeded for this endothermic reaction is sup-plied by the sun.

Applications of Heat ofReactions

We have seen that energy changes arean integral part of chemical processes. Weclose this chapter with a brief look at someof the applications of energy changes ineveryday life. All life and our society aredependent on a multitude of exothermicreactions. The metabolic breakdown of foodin organisms supplies the energy theyrequire to exist. For humans, the exothermiccombustion of fossil fuels powers our mod-ern industrial society. It is interesting to lookat the energy content of different fuels. Table10.5 lists the energy content of several fuels.

The figures in Table 10.5 indicate thatthe biomass fuel ethanol, proposed as analternative to fossil fuels, yields signifi-cantly less energy per gram than octane.This means that driving a vehicle withethanol will decrease the fuel efficiency of

Table 10.4Four Cases Describing How Heat and Work Affect Internal Energy of a System

Table 10.5Energy Content of Selected Fuels


a car. It should be noted, though, thatethanol is sometimes used as an additive todecrease the amount of carbon monoxideemissions. The table also shows that theenergy content of fat is over twice that ofcarbohydrates, and also almost twice that ofprotein. The higher energy content of fatsand the fact that they are more readily storedin the body explain why hibernating organ-isms build up a high fat content during theyear.

One common application of thermo-chemistry in everyday life is the use of hotand cold packs. A variety of these are usedto treat injuries and provide warmth. Onetype of pack used to provide heat utilizes theexothermic oxidation of iron to produceheat. The reaction can be represented as

4Fe(s) � 3O2(g) 2Fe2O3 �H � –1650 kJ

A hot pack is activated when someone ini-tiates the reaction by some physical actionsuch as shaking or breaking a seal on thepack. The actual hot pack reaction involvesseveral other chemicals, but the primaryreaction is the oxidation of iron. The ironin the hot pack is a fine powder to increasethe efficiency of the reaction. As expected,cold packs depend on some type ofendothermic process. One common coldpack is based on the mixing of ammonianitrate and water. In this case, the heat of

solution is positive, indicating an endother-mic reaction:

NH4NO3(s) � H2O(l) NH4�

(aq) �NO3

–(aq) �H � �26kJ

When the reaction is initiated by breaking abarrier separating the ammonia nitrate andwater contained in the pack, heat isabsorbed from the surroundings. The imme-diate surroundings in this case happen to bethe body part to which the cold pack isapplied.

In this chapter, we have examined theenergy aspects of chemical processes. Whileour focus on chemistry until this chapter hasbeen on matter, we must realize that a com-plete study of chemical processes must alsoinclude energy. The principles of thermo-chemistry can help us answer fundamentalquestions such as how organisms maintainthemselves or what is the ultimate fate of theuniverse. On a more practical level, phasechanges and thermochemistry form thebasis of many common devices. Car radia-tors, refrigerators, power plants, air condi-tioners, and heat pumps all employ phasechanges to the benefit of modern society.The combustion of fuels powers modernsociety. Now that we have laid the ground-work for including energy considerations inour examination of chemical principles, weinclude these principles in coming chapters.

Chapter 11

IntroductionMuch of the world around us is composedof solutions. Air, oceans, steel, gasoline,and soda pop are just a few examples ofcommon substances that exist as solutions.A solution is nothing more than a homoge-neous mixture in which the particles rangein size from 0.2 to 2.0 nm. There are otherhomogeneous mixtures with smaller parti-cle sizes. These are known as colloids andsuspension, but our main focus in this chap-ter is on solutions. When considering solu-tions, it helps to think in terms of onesubstance dissolved in another substance.The simplest solutions consist of two com-ponents. The component in the greateramount is referred to as the solvent, and thecomponent in the lesser amount is called thesolute. Therefore, a solution can be thoughtof as a homogeneous mixture containing asolute dissolved in a solvent:

solute � solvent � solution

By far, the most familiar solutions con-sist of solid solutes dissolved in water. Theoceans are giant solutions containing hun-dreds of solutes; the major ones exist in

ionic form and include sodium, chloride,magnesium, iodide, calcium, and carbonate(Table 11.1). Although the most familiarsolution consists of a solid dissolved in aliquid, solutions can consist of numerouscombinations of phases. Table 11.2 summa-rizes different types of solutions based onthe state of the solute and solvent.

11

Solutions

Table 11.1Major Ions Present in Seawater

126 Solutions

Table 11.2 shows that many differenttypes of solutions exist. In this chapter, weare primarily concerned with solutions inwhich water is the solvent. These are knownas aqueous solutions.

The Solution ProcessWhy is it that some substances readily

mix to form solutions while others do not?Whether one substance dissolves in anothersubstance is largely dependent on the inter-molecular forces present in the substances.For a solution to form, the solute particlesmust become dispersed throughout the sol-vent. This process requires the solute andsolvent to initially separate and then mix.Another way of thinking of this is that thesolute particles must separate from eachother and disperse throughout the solvent.The solvent may separate to make room forthe solute particles or the solute particlesmay occupy the space between the solventparticles. Determining whether one sub-stance dissolves in another requires exam-ining three different intermolecular forcespresent in the substances—between the

solute particles, between the solvent parti-cles, and between the solute and solvent par-ticles. The formation of a solution can occurwhen the magnitude of all three of theseforces are similar. Conversely, if the inter-molecular forces for solute particles aremuch stronger than their attraction for sol-vent particles, then mixing is not favored.The tendency of substances with similarintermolecular forces to mix leads to thegeneral rule of thumb: “like dissolves like.”Substances with “like” intermolecularforces tend to form solutions. As a simpleanalogy for the solution process, assumetwo situations take place in a stadium in atown with two rival teams. In the first situ-ation, consider a football game between thetwo cross-town rivals. The fans from oneside of town will occupy one side of the sta-dium, while the fans from the opposite sideof town will sit together on the other side.There is no attraction of either side for theother, and this corresponds to the situationwhen two substances do not mix to form asolution. On the other hand, consider a sec-ond situation in which a rock concert is heldin the stadium. In this case, people from the

Table 11.2Examples of Solutions

Solutions 127

opposite sides of towns will mix throughoutthe stadium. Fans from both sides of townare equally attracted to the music, and thereis an equal desire to get the best seats. Thissecond situation is analogous to what hap-pens when two substances mix to form asolution.

The “like dissolves like” rule providesa general guideline for determining whichtype of substances will form solutions. Tosee how this rule applies, let’s look at thesolubility of some of the major types ofcompounds. Ionic compounds tend to dis-solve in polar solvents. Because water is themost common polar solvent, ionic com-pounds tend to form aqueous solutions. Forexample, NaCl dissolves in water. Considerwhat happens when a few grains of salt aresprinkled into water. Each salt grain is acrystalline lattice structure containing mil-lions of unit cells of NaCl. Although theionic bond holding the NaCl unit cellstogether is strong, the polarity of watercauses the positive hydrogen end of water tobe attracted to the chloride ion (Cl–) and thenegative oxygen end to be attracted to the

sodium ion (Na�). Water molecules initiallyinteract with the NaCl crystals on the outersurface of the grains. Water molecules actcollectively to pull individual ions out of thecrystalline structure (Figure 11.1). The gen-eral process by which a solute becomes dis-persed in a solution is known as solvation.When water is the solvent, the process isknown as hydration.

While many ionic compounds are solu-ble in water, many are not. The term “solu-bility” is somewhat subjective. There areactually degrees of solubility. A substance isconsidered soluble if 0.1 moles of it can dis-solve in 1 liter of water. If less than 0.001mole of the substance dissolves in water, asubstance is considered insoluble. Partiallysoluble substances fall between these twoextremes. Table 11.3 summarizes the solu-bility of some major groups of ionic com-pounds in water.

Whether an ionic compound dissolvesin water depends on the strength of theionic bond holding the compound together.Water must have sufficient strength tobreak the ionic bond. The strength of the

Figure 11.1Diagram of Hydration of NaCl (Rae Déjur)

128 Solutions

bond depends on the size and charge of theions in the compound. Stronger ionic bondsoccur when the ions are smaller and theions carry multiple charges. Because ofthis, many ionic compounds may not be sol-uble in water, as is seen by noting theexceptions listed in Table 11.3.

Polar covalent molecules may or maynot dissolve in water depending on whetherthey have the ability to form hydrogenbonds. For example, many alcohols will dis-solve in water because the OH group char-acteristic of alcohols gives them the abilityto hydrogen bond with water. The attractionbetween alcohol and water is demonstratedwhen equal volumes of the two are mixed togive a total less than the sum of their indi-vidual volumes, for example, 100 mL ofalcohol � 100 mL of water produces lessthan 200 mL of solution. Nonpolar sub-stances tend not to dissolve in water, butthey do dissolve in nonpolar solvents. Fats,oil, grease, and gasoline, for example, donot dissolve in water, but they form a layeron top of water. When substances do notmix but form distinct layers, they arereferred to as immiscible. Nonpolar sub-

stances can be dissolved in a nonpolar sol-vent such as benzene or carbon tetrachlo-ride.

ElectrolytesA useful characteristic in classifying

solutions is their ability to conduct electric-ity. Solutions that are good conductors ofelectricity are said to contain strong elec-trolytes. Strong electrolytes have the abilityto produce ions when dissolved in water; thegreater the degree of ionization of a sub-stance is, the stronger the electrolyte. Strongelectrolytes include soluble salts, strongacids, and strong bases. Many substances,such as weak acids and weak bases, only par-tially ionize when dissolved in water. Thesesubstances are referred to as weak elec-trolytes. Still other substances do not ionizeat all, but they dissolve as whole moleculesand are referred to as nonelectrolytes. Sugaris an example of a nonelectrolyte. Whensugar dissolves, the covalent bonds holdingthe molecule together are sufficiently strongto keep the molecule together. Table sugar,C12H22O11, contains a number of oxygen

Table 11.3Solubility of Ionic Compounds in Water at 25°C

Solutions 129

atoms making it a polar molecule. Its highnumber of oxygen atoms allows it to hydro-gen bond to many water molecules, andtherefore, sugar is quite soluble.

Soluble ionic compounds tend to bestrong electrolytes, while alcohols andorganic compounds are nonelectrolytes.Remember that classification as a strongelectrolyte, weak electrolyte, or nonelec-trolyte is somewhat subjective. Freshwatercan be either a weak electrolyte or a non-electrolyte depending on its purity. Theimportant consideration in classifying a sub-stance is to what extent an aqueous solutionof the substance will conduct electricity.

Concentration of SolutionsThe amount of solute and solvent in a

solution can be quantitatively expressedusing numerous concentration units. Thechoice of a particular concentration unitdepends largely on practice and convenience.We have probably all made solutions usingrecipes or directions that tell us to add somuch water to a substance. In the field ofchemistry, the most common concentrationunits are molarity, molality, percent by mass,and “parts per.” Each of these is defined here:

Molarity is the moles of solute per liter of solu-tion. It is abbreviated with a large “M.” A 1 Msolution contains 1 mole of the substance dis-solved in 1 liter of solution. It is important torealize molarity is the amount of solute per literof solution. If a 1 M solution of sodium hydrox-ide were to be prepared, 40 grams of NaOHwould be diluted to 1 liter. Special flasks of var-ious volumes are made for just this purpose.Adding 1 liter of water to 40 grams would notproduce a 1 M solution, but a solution of slightlyless than 1 M because adding 1 liter of water to40 grams of NaOH makes the final volume ofthe solution greater than 1 liter.

Molality is the number of moles of solutedissolved in 1 kilogram of solvent. It is abbrevi-ated by “m.” A 1 m solution of NaOH contains

40 grams of NaOH dissolved in 1 kilogram ofwater. Molality is especially important in work-ing with colligative properties such as boilingpoint elevation and freezing point depression.

Percent by mass is the mass of the solutedivided by the mass of solution multiplied by100.

Parts per a particular quantity is used exten-sively in environmental work and is useful whenexpressing the concentration of dilute aqueoussolutions. Common concentration expressionsusing this unit include parts per thousand, partsper million, and parts per billion. Ocean salinityis expressed in parts per thousand (ppt). Typicalocean salinity is 35 ppt and can be translated tomean that for every 1,000 grams of ocean waterthere are 35 grams of dissolved material. Partsper million (ppm) is used frequently in environ-mental work and to express the solubility ofgases in water. For example, the solubility ofoxygen in a lake may be given as 10 ppm.Because the density of water is approximately 1mg/L, a ppm is the same as mg/L for dilute aque-ous solutions. A part per billion (ppb) is also frequently used for water standards. The recom-mended drinking water standard for lead is 15ppb. A part per billion is the same as 1 g/L.Modern instruments have allowed chemists tomeasure progressively smaller concentrations ofsubstances in the environment, some at parts pertrillion concentrations. Even though it is possibleto detect toxic substances at very low concen-trations, this should not be interpreted as mean-ing these substances present a threat to humanhealth. A low concentration may be statisticallyequivalent to zero, presenting no problem what-soever.

While concentration units indicate howmuch solute and solvent make up a solutionat any given time, often it is desirable toknow the concentration when the solution issaturated. A solution is saturated when themaximum amount of solute has dissolved inthe solvent at a particular temperature. Whenthe solvent holds less than its maximumamount of solute, the solution is said to beunsaturated; when the solution holds more

130 Solutions

than the maximum amount, it is said to besupersaturated. The amount of solute thatcan dissolve in a solvent can be expressedusing different measures of solubility. Molarsolubility is the number of moles of solute in1 liter of saturated solution. Handbooks gen-erally report the solubility of aqueous solu-tions by giving the amount of solute that willdissolve in 100 grams of water at a specifictemperature. The solubility of NaCl is 35.7grams per 100 grams of water at 0°C and37.3 grams per 100 grams of water at 60°C.Once a solution becomes saturated, anysolute added will settle out or precipitate.The additional solute that settles out is calledthe precipitate. For example, using the solu-bility of NaCl at 60°C, the solution would beundersaturated until 37.3 g NaCl have beenadded to 100 grams of water. At this point,the solution becomes saturated and any addi-tional salt added will precipitate out as solid.

In any discussion of solubility, it isimportant to remember solubility is temper-ature dependent. Generally, the solubility ofsolids in liquids increases with temperature.The variation of solubility in water variesgreatly for different solutes. Figure 11.2demonstrates that the solubility may increase

significantly with temperature for somesolutes, may be relatively constant (NaCl),and may even decrease with temperature.The variation in solubility with temperatureshould not be confused with the rate at whicha solute dissolves. Increasing the tempera-ture causes the solute to dissolve faster asdoes stirring and increasing the surface areaof the solute, but just because a substancedissolves faster does not mean more of it dis-solves.

In some cases a solution may becomesupersaturated. A good example of this issodium acetate, CH3COONa. When a solu-tion containing sufficient sodium acetate isheated until the solid dissolves and allowedto cool slowly with minimal disturbance, thesodium acetate does not crystallize out asthe solution cools. For the sodium acetate tosettle out, a small seed crystal must beadded to the solution to initiate the crystal-lization process.

Solubility of GasesThe focus in this chapter has primarily

been on solid solutes in liquid solvents. Thesolubility of gases in water and other liquids

Figure 11.2The solubility of several ionic compounds varies according to temperature.

Solutions 131

(such as in blood) varies in several respectsfrom that of solids. One of the most obviousdifferences is the degree of solubility. Gasesare not very soluble in water. The saturatedvalue for oxygen in freshwater is 14.6 mg/Lat 0°C and 9.2 mg/L at 20°C. The saturationvalues for oxygen illustrate that the solubil-ity of gases is inversely related to tempera-ture. The fact that gas solubility decreaseswith increased temperature is the reason forthermal pollution associated with powerplants. Power plants use vast quantities ofwater to condense steam in the generationof electricity. As seen in the last chapter,condensation is an exothermic process. Theheat given up by the steam raises the tem-perature of the cooling water. When thewarmed cooling water is released into theenvironment, for example into a river orlake, the elevated temperatures cause a dropin the oxygen solubility of the water. Thelowered oxygen levels may place a signifi-cant stress on aquatic organisms living inthe vicinity of the power plant. This stress isnot only due to the decreased oxygen con-tent, but also due to the impact of elevatedtemperatures on the metabolism of cold-blooded organisms.

Another factor that differentiates thesolubility of gases from solids and liquids isthe effect of pressure. The effect of pressureon gas solubility was studied extensively bya contemporary and close associate of JohnDalton named William Henry (1775–1836).Henry’s Law states that the solubility of agas is directly proportional to the partialpressure of that gas over the solution. Statedmathematically, Henry’s Law is c � kP,where c is the concentration of the dissolvedgas in moles per liter, k is Henry’s law con-stant for the solution, and P is the partialpressure of the gas above the solution.Henry’s Law is demonstrated every time acarbonated beverage is opened. During thecarbonation process, carbon dioxide is dis-

solved in the beverage under increased pres-sure. When the beverage container is sealed,a small space above the beverage fills withcarbon dioxide along with other gases (air,water vapor). The partial pressure of the car-bon dioxide above the beverage in theclosed container is many times higher thanthat found in the atmosphere, where it isonly about 0.0004 atmosphere. When thecontainer is opened, the partial pressure ofthe CO2 will immediately drop. Accordingto Henry’s Law, a drop in the partial pres-sure of CO2 above the solution results in adecreased solubility of carbon dioxide in thebeverage. This is seen whenever a carbon-ated beverage is opened and the CO2 isobserved escaping as a steady stream ofbubbles. Henry’s Law also explains why thebeverage goes flat if it is left open to theatmosphere for any length of time.

Colligative PropertiesAnyone who has ever made ice cream

knows that the addition of rock salt to icecauses it to melt and produce a liquid-icesolution below 0°C. This is just one exampleof how the physical properties of a solutiondiffer from those of a pure solvent. Proper-ties that depend on the amount of solute present in a solution are termed colligativeproperties. Colligative means collectiveproperties. These properties are termed col-ligative, because the properties depend onthe collective number of particles present insolution rather than the types of particles.The major colligative properties and howthey affect solutions compared to their puresolvents are summarized in Table 11.4.

To examine the four properties listed inTable 11.4, we can use the simple case of anonelectrolyte, nonvolatile solute. The low-ering of the vapor pressure is a conse-quence of nonvolatile solute particlesoccupying positions at the surface of the

132 Solutions

solution. Figure 11.3 shows two identicalcontainers, one containing an aqueous solu-tion, for example a sugar solution, and theother water. The surface of the solution isoccupied by both solute and solvent parti-cles. Because the solute is nonvolatile,solute particles do not vaporize but stay insolution and reduce the surface area occu-pied by the solvent water molecules. Thereduction in the surface area for the solu-tion means that fewer molecules are able toenter the vapor phase as compared to thepure solvent. The solute particles can bethought of as a barrier that impedes the sol-vent from entering the vapor phase, andtherefore, reducing the vapor pressure. Thehigher the concentration of the solution, themore the vapor pressure is lowered.

The boiling point is defined as the tem-perature at which the vapor pressure of aliquid is equal to the external pressure, typ-ically 1 atmosphere. If adding a solute low-ers the vapor pressure, then a highertemperature is required to reach the pointwhere the vapor pressure of the liquid isequal to the external pressure. Therefore,the boiling point is higher for a solution ascompared to its pure solvent. The change inboiling point due to addition of a solute isknown as boiling point elevation. The boil-ing point elevation can be calculated usingthe equation �Tb � Kbm, where �Tb is theincrease in boiling point for the solution, Kb

is the molal boiling point elevation con-stant, and m is the molality of the solution.Molal boiling point elevation constants aregiven in reference books for different sol-vents. Water’s Kb is equal to 0.51 °C/m.Using this value, a 2.0 molal sugar solutionboils at a temperature of 101.02°C ratherthan 100°C.

A solution also exhibits a depression inits freezing point. The freezing point depres-sion is the decrease in the temperature of thefreezing point due to the addition of asolute. It is calculated using the equations�Tf � Kfm, where �Tf is the decrease infreezing point for the solution, Kf is themolal freezing point depression constant,and m is the molality of the solution. Water’sKf value is 1.86°C/m.

Table 11.4Colligative Properties

Figure 11.3In a solution, solute particles at the surfacereduce the amount of solvent that can enterthe vapor phase resulting in a lower vaporpressure.

Solutions 133

A relevant application of freezing pointdepression and boiling point is the additionof antifreeze to a car’s radiator. Ethyleneglycol, CH2(OH)CH2(OH), is a popularantifreeze. Producers of antifreeze recom-mend that a 50% antifreeze and a 50% watermixture be used for protection. An averagecapacity for the cooling system of a car isabout 14 quarts. Seven quarts of water hasa mass of approximately 6.6 kg. The densityof ethylene glycol is about 1.1 g/mL, and itsmolar mass is 62 g. These values are used todetermine that there are about 117 molescontained in 7 quarts of ethylene glycol. Themolality of the water-antifreeze mixturewould be 117 moles/6.6 kg � 18 m. Usingthe values for Kb and Kf, gives the freezingpoint of the water-antifreeze solution as–33°C and the boiling point as 109°C. Thesetemperatures would be the same for any50% antifreeze solution independent of thecooling system capacity.

The last of the four colligative proper-ties in Table 11.4 is osmotic pressure.Osmosis occurs when solvent moleculesmove through a semipermeable membranefrom the pure solvent or diluted solution toa more concentrated solution. A semi-permeable membrane acts like a sieve or fil-ter. It allows movement of solvent particlesto pass through, but it blocks solute particles(Figure 11.4). Figure 11.4 shows only thenet movement of solvent particles from thesolvent to the solution. It is important torealize that solvent particles pass throughthe membrane in both directions, but therate of movement from the solvent to soluteis greater than in the reverse direction. Thepressure required to stop the net movementof solvent particles across the semiperme-able membrane is called the osmotic pres-sure of the solution. The osmotic pressure is given by the formula � � MRT, where �is the osmotic pressure in atmospheres, M isthe molarity of the solution, R is the ideal

gas law constant, and T is the absolute tem-perature.

Osmosis is extremely important inphysiology. Cell membranes, capillaries,and the lining of digestive tracts act assemipermeable membranes. Osmosis con-stantly occurs through cell membranes.Cells generally have the same concentra-tion as the fluids surrounding them. Ininstances where the surrounding fluid hasa higher or lower solvent concentrationcompared to the cell, water flows into orout of the cell, respectively. The formercondition is known as plasmolysis and thelatter as crenation.

A process similar to osmosis is dialysis.In dialysis, a dialyzing membrane allowsboth solvent and minute solute particles of acertain size to pass. The passage of theseparticles, like osmosis, is from a region ofhigh concentration to a region of low con-centration. Our kidneys are a dialysis sys-tem responsible for the removal of toxicwastes from the blood. In artificial dialysis,blood is circulated through dialysis tubesmade of membranes, which are immersed ina “clean” solution that lacks the wastes inthe blood. The impurities are filtered out of the blood as they move across the wallsof the dialyzing tubes.

Figure 11.4In osmosis, there is a net movement ofsolvent from a region of higherconcentration to a region of lowerconcentration.

134 Solutions

Reverse osmosis is a process in whichfreshwater is obtained from saltwater byforcing water from a region of low fresh-water concentration to a region of highfreshwater concentration. This is opposite ofthe natural process, and hence the namereverse osmosis. In a typical reverse osmo-sis process, a pressure of approximately 30atmospheres is required to force freshwaterto move from seawater across a semiperme-able membrane (Figure 11.5).

In the discussion of colligative proper-ties, it was assumed that the solutions werenonelectrolytes. When electrolyte solutionsare considered, remember that colligativeproperties are dependent on the number ofparticles in solution. While a 1 m concen-tration of a nonelectrolyte, such as sugar,yields 1 mole of particles per kg of solvent,an electrolyte such as NaCl theoretically dis-sociates to give 1 mole of sodium ions and1 mole of chloride ions. In reality, a 1 mNaCl solution produces slightly less than 1mole of sodium and chloride ions. This isbecause water is not perfectly efficient inhydrating all ions, and some Na� and Cl– areattracted to each other in solution and act asa single particle. This single particle is

known as an ion pair. The greater the con-centration of the solution is, the greater thechance of ion pairs forming. For dilute con-centrations of NaCl, it is assumed there areabout 1.9 moles of particles for a 1 m solu-tion of NaCl. The 1.9 should be used as acorrection factor in the boiling point ele-vation, freezing point, and osmotic pressureequations. To see how this works, let’s deter-mine the osmotic pressure of seawater,which is 30 atm. The equation is � � MRT.If it is assumed that the salinity of seawateris 35 ppt, this means there are 35 grams ofsalt per 1,000 grams of seawater. Assumethat all the salt is in the form of NaCl, andthe density of seawater is approximately 1g/mL; this gives the number of moles of saltas 35g/58 g/mole � 0.60 mole. The molar-ity of seawater is found by taking this valueand dividing by the number of liters of solu-tion: molarity � .60 mole/ 1.0 L � 0.60 M.If we use the equation for osmotic pressure,we find � by multiplying (0.60m)(0.082atm-L/m)(298K). This value is slightly lessthan 15 atmospheres. The calculated valuefor seawater does not take into account thatwe get almost two ions for every NaCl par-ticle. Therefore, our calculated value mustbe doubled to give 30 atmospheres for theosmotic pressure of seawater. To determinethe boiling point elevation or freezing pointdepression of a salt solution, the factor 1.9would be used in the appropriate equations.Correction factors, called van’t Hoff fac-tors, for a number of different electrolytescan be found in chemical reference books.

Reactions of SolutionsMost reactions in chemistry involve

solutions. Reactions between solutions con-tinually take place in the atmosphere, ocean,and natural environment. A host of chemi-cal processes such as the refining ofpetroleum, production of steel, purification

Figure 11.5In reverse osmosis, a pressure equal to theosmotic pressure of seawater is applied toobtain freshwater from seawater.

Solutions 135

of water supplies, and refining of mineralsinvolve solutions. In fact, solution reactionsnot only occur all around us, but also withinus. Our survival depends on numerous solu-tion reactions taking place within our bod-ies. The digestion of food, the transport ofoxygen to cells, and the removal of wasteproducts are just a few body functionsinvolving solution reactions. We concludethis chapter on solutions by taking a brieflook at precipitation reactions. A precipita-tion reaction takes place when two solutionsare mixed, and a solid, called a precipitate,separates from the solution.

To understand how a precipitation reac-tion works, let’s examine the reaction be-tween two aqueous ionic solutions, silvernitrate (AgNO3) and potassium chloride(KCl). According to the solubility rules pre-sented in Table 11.3, both these ionic com-pounds are soluble in water. When thesesolutions are mixed, a white precipitateimmediately forms. Referring again to Table11.3, rule 3 indicates chlorides are soluble,but that a notable exception is for chloridecompounds containing Ag�. Based on thisinformation, the white precipitate must besilver chloride, AgCl. The mixing of the twosolutions is presented in Figure 11.6. Thetwo original solutions contain the potas-sium, chloride, silver, and nitrate ions inaqueous solution. When these are mixed, thesilver and chloride react to produce a pre-cipitate. The potassium and chloride ionsremain in solution and the process can berepresented with the molecular equation:

KCl(aq) � AgNO3(aq) AgCl(s) � KNO3(aq)

Rather than write the chemical equationusing whole units, a more accurate pictureof the reaction is represented by writing thecomplete ionic equation:

K�(aq) � Cl–

(aq) � Ag�(aq) � NO3

–(aq)

AgCl(s) � K�(aq) � NO3

–(aq)

Writing the reaction using ions shows thatthe potassium and nitrate ions are present onboth sides of the equation. These ions do nottake part in the overall reaction and arecalled spectator ions. The general ionicequation can be simplified by dropping thespectator ions and writing the net ionicequation:

Ag�(aq) � Cl–

(aq) AgCl(s)

The net ionic equation shows only the ions that are actually taking part in the reac-tion.

Precipitation reactions follow the samegeneral pattern as potassium chloride andsilver nitrate. The solubility rules in Table11.3 can be used to predict whether a pre-cipitate will form. In many cases, solutionsdo not form a precipitate when mixed. Inthese cases, it is assumed all ions stay insolution. Precipitation reactions can be usedto our advantage. In water quality work, onemethod of removing pollutants is to precip-itate the pollutant out of the water in a treat-ment plant. The pollutant can then berecycled or disposed of rather than beingreleased into the environment. For example,phosphates causing eutrophication (rapidaging) in freshwater lakes can be removedby adding FeCl3 and alum to the water.Another use of precipitation reaction is to

Figure 11.6When two solutions are mixed, a solidprecipitate may form when the ions form aninsoluble compound.

136 Solutions

identify the presence of specific ions inwater. For instance, if a water sample wassuspected of containing lead, Pb2�, the sam-ple could be mixed with a solution contain-ing iodide (or chloride, bromide) to see if aprecipitate formed. The method of qualita-tive analysis involves conducting a system-atic series of reactions using differentsolutions to determine the presence orabsence of particular ions. By noting whichsolutions produce precipitates and which donot, many unknowns can be identified.

A number of precipitation reactions aredetrimental and cause problems. Scaling inpipes due to the precipitation of calcium isa problem in regions of hard water. Hardwater is defined as water having a high, typ-ically greater than 100 ppm, concentrationof calcium and/or magnesium. Water heatedin hot water tanks and hot water heating sys-tems results in the precipitation of calciumcarbonate according to the following reac-tion:

Ca2�(aq) � 2HCO3

–(aq) CaCO3(s)

� CO2(aq) � H2O(l)

This reaction shows calcium ion reactingwith bicarbonate to yield calcium carbonate,carbon dioxide, and water.

Another example of an undesirable pre-cipitation reaction involves the formation ofkidney stones in the human body. The majortype of kidney stones consists of calcium incombination with oxalate (C2O4

2�). Thereaction can be represented as

CaCl2(aq) � H2C2O4(aq) CaC2O4(s) � 2HCl(aq)

This reaction shows calcium chloride react-ing with oxalic acid to produce calciumoxalate and hydrochloric acid. The calciumoxalate is the kidney stone. Calcium oxalateprecipitates out of the urine in the kidneysof all individuals. Normally, the solidsformed are grain size and do not cause prob-

lems. In certain individuals, though, the pre-cipitates can grow to appreciable size, caus-ing considerable discomfort and healthproblems. One recommendation for thosewith kidney stones is to increase their waterintake and reduce the consumption of foodsrich in oxalic acid such as chocolate,peanuts, and tea. Other precipitates that canform in the kidney and lead to stone forma-tion are calcium phosphate, Ca3(PO4)2, andcalcium carbonate, CaCO3.

Colloids and SuspensionsThe two other types of mixtures men-

tioned at the beginning of this chapter werecolloids and suspensions. These mixturescan be differentiated from solutions basedon particle size. Colloid particles are largerthan particles in solution. They range indiameter from 2.0 to 1,000 nm. Colloid particles cannot be filtered out, remain insuspension, and scatter light. Commonexamples of colloids are fog and milk. Sus-pensions consist of particles greater than1,000 nanometers in diameter. They aregenerally opaque to light, separate on stand-ing, and can be filtered. Blood is an exampleof a suspension.

Colloids were initially differentiatedfrom solutions by Thomas Graham (1805–1869) as a result of his work concerning thediffusion of particles through membranes.Graham observed that certain substancessuch as starches, glues, and gelatins did notdiffuse through membranes like ionic solu-tions. Graham used the term “colloids,”which means glue in Greek, to describethese substances. Graham’s work on the dif-fusion of materials through membranes alsoled him to the discovery that membranesallowed small molecules and ions to pass,but blocked the movement of colloids. Gra-ham coined the word dialysis to describethis process.

Solutions 137

A colloid is also called a colloidal dis-persion. A colloid dispersion consists of twocomponents similar to a solution. The parti-cles themselves are the dispersed phaseand are analogous to the solute in a solution.The dispersing medium is similar to thesolvent. Some examples of different typesof colloids are summarized in Table 11.5.

Colloid particles tend to acquire similarcharges on their surface. This results in theparticles repelling each other, and therefore,they stay dispersed and tend not to settle.Another characteristic of colloids is theirability to scatter light. This phenomenon isknown as the Tyndall effect, named afterthe English physicist John Tyndall(1820–1893). This effect is observed whencolloidal dust and water particles in the airscatter light from a movie projector or spot-light making the light beam visible.

As the examples in Table 11.5 indicate,many familiar products are colloids. Thecleaning ability of soap is due in part to dirtparticles emulsifying in the soap. Food andother nutrients in our blood are transportedto cells as colloidal particles. The process ofjelly making involves producing a sol thatsets up as a semisolid. A similar processoccurs when making gelatin. Sols that set uplike jellies are technically called gels.

Just as solutions precipitate, there areprocesses that cause colloidal particles togroup together and settle. The process inwhich colloidal particles group together iscalled coagulation. Coagulation can be ini-tiated by heating or adding an acid to a col-loid. The coagulation of colloids is observedevery time an egg is fried. The colloids inthe albumen of the egg coagulate producinga solid white mass when heated.

Table 11.5Types of Colloids

IntroductionTwo important aspects of chemical reactionsare how fast they occur and to what extentthe reaction takes place. The area of chem-istry that deals with the speed of chemicalreactions is known as chemical kinetics. Inmany reactions, reactants are only partiallyconverted into products. In these reactions astate is reached in which the concentrationsof reactants and products remain constant. Atthis point, the reaction is said to have reachedequilibrium. In this chapter, we exploreeach of these important areas of chemistry.

Reaction RatesIt is well known that reactions take place

at different rates. Some reactions occur in aninstant, while others can be measured in sec-onds, minutes, and hours. The explosivecombustion of gasoline in the cylinder of acar occurs in a fraction of a second. Reac-tions used to develop film take place overseconds to minutes. Chemical reactionsresponsible for changing the color of fallleaves take days to weeks, and the oxidationof metals takes place over years. Wheneverreactions occur, chemical kinetics play a

vital role. Kinetics dictate how fast productscan be produced in the chemical industry,how long it takes for medicines and drugs toact in our bodies, how quickly the depletionof the ozone layer occurs, and how long ittakes food to spoil. Knowledge of kineticsgives us insights into not only how fast theyoccur, but also how they occur.

Because kinetics deals with the rates ofchemical reactions, it is important to estab-lish a definition for reaction rate. Becausemost reactions occur in or between solu-tions, it is logical to define reaction rates interms of changes in concentration. Reactionrate is typically defined as the change inconcentration of a reactant or product overtime. For example, hydrogen peroxidedecomposes slowly over time. Its decompo-sition is accelerated by heat and exposure tolight, and therefore, it is stored in opaquecontainers. The reaction is represented bythe equation:

2H2O2(l) 2H2O(l) � O2(g)

Hydrogen peroxide sold in stores is typicallya 3% solution. This solution is equivalent toapproximately 0.90 M. The rate of thedecomposition of H2O2 can be expressed

12

Kinetics and Equilibrium

140 Kinetics and Equilibrium

by the change in hydrogen peroxide con-centration over time:

If after one year the concentration was 1.5%or 0.45 M, then the reaction rate would beequivalent to –0.45 M/year. Rather thanlooking at the disappearance of hydrogenperoxide, the reaction rate could beexpressed in terms of the production ofwater or oxygen. Expressing the reactionrate in terms of the amount of oxygenformed, the stoichiometry of the reactioncould be used to determine the rate of for-mation of oxygen. The reaction shows 1mole of molecular oxygen is formed forevery 2 moles of hydrogen peroxide thatdecompose. The amount of oxygen formedafter one year would be half of 0.45 M orabout �0.22 M/yr. Notice that the plus signindicates a species is forming and a negativesign indicates a species is disappearing.

The Collision TheoryAn understanding of reaction rates can

be explained by adopting a collision modelfor chemical reactions. The collision theoryassumes chemical reactions are a result ofmolecules colliding, and the rate of the reac-tion is dictated by several characteristics ofthese collisions. An important factor thataffects the reaction rate is the frequency ofcollisions. The reaction rate is directlydependent on the number of collisions thattake place, but several other important fac-tors also dictate the speed of a chemicalreaction.

Consider a piece of coal, which for ourpurposes can be considered pure carbon, sit-ting on the table. The coal is exposed to bil-lions upon billions of collisions with oxygenmolecules, O2 each second. Yet, the reactionC � O2 CO2 does not occur. The coal

could sit on the desk for years and experi-ence countless collisions with oxygen with-out reacting. Obviously, merely the presenceof collisions does not mean a reactionoccurs. It’s similar to saying we bruise whenwe come into contact with an object. Whilecontact with an object may be a necessarycondition to bruising, we do not get a bruiseevery time we touch something.

To understand why most collisions donot cause a reaction, the energetics of areaction must be considered. As noted sev-eral times, reactions involve the breakingand formation of bonds. Energy is requiredto break bonds, and energy is released whenbonds form. According to the collision the-ory, the energy needed to break bondscomes from the kinetic energy of the col-liding molecules. During a collisionbetween molecules, kinetic energy is con-verted to potential energy. More often thannot, there is simply not enough kineticenergy to break bonds and initiate a reac-tion. The situation can be compared todropping a rubber ball from a certain heightonto the floor. As the ball falls, gravitationalpotential energy is converted to kineticenergy. The ball’s kinetic energy reaches amaximum value just before it contacts thefloor. When it strikes the floor, the ball isdeformed as kinetic energy is converted topotential energy. This potential energy isreferred to as elastic potential energy and issimilar to the energy stored when a spring is compressed. When the ball has reachedits maximum deformation, it momentarilystops and has no kinetic energy. At thispoint, the elastic potential energy stored inthe rubber ball is at a maximum. Thisenergy is converted back into kinetic energyas the ball bounces back up. If the ball werethrown with sufficient energy, it wouldbreak apart when it struck the floor. This isanalogous to molecules colliding. Most ofthe time they just bounce off each other, but

� ConcentrationRate �

� time

Kinetics and Equilibrium 141

if they have enough kinetic energy, they canbreak apart.

During chemical reactions, the kineticenergy of the molecules are converted intopotential energy during collisions. The col-lisions distort and stretch the bonds; thebonds can be viewed as acting like tinysprings or rubber bands. If the energy of thecollision is sufficient, bonds break, similarto the breaking of a rubber band if it isstretched too far. In a chemical reaction, theminimum energy needed to break bonds andinitiate a reaction is known as the activationenergy. No matter how many collisionsoccur between reactants, a reaction does notoccur unless the energy of the collision isequal to the activation energy. Therefore,while the carbon in the lump of coal experi-ences an endless number of collisions withoxygen, these collisions are not sufficient toprovide the activation energy for the com-bustion of coal. The energy in the collisionsis insufficient to break the double bond ofthe oxygen molecule.

The activation energy is often depictedas an energy hill or barrier that reactantsmust overcome for a reaction to take place.

Two situations are shown in Figure 12.1 cor-responding to exothermic and endothermicreactions. When the reaction is exothermic,the total potential energy of the products isless than the potential energy of the reac-tants. The difference in potential energies ofthe reactants and products equals the changein internal energy of the system. Becauseenergy flows from the system to the sur-roundings in an exothermic reaction, theinternal energy decreases and �E is nega-tive. In an endothermic reaction, the prod-ucts have a higher potential energy than thereactants, energy flows from the surround-ings into the system, and �E is positive.

At the top of the energy hill a transitionstate exists called the activated complex.The activated complex represents the stagein the reaction where reactants may be trans-formed into products. The idea of an acti-vated complex was first proposed by SvanteAugust Arrhenius (1859–1927) around1890. Arrhenius received the Nobel Prize inchemistry in 1903. Reactants may reach theactivated complex stage and still not be con-verted into products, but “slide” down theactivation energy hill back to reactants. On

Figure 12.1The activation energy, Ea, acts as an energy barrier. For a reaction to occur, the amount ofkinetic energy converted to potential energy during a collision must be greater than the activa-tion energy.


the other hand, when a reaction does occur,then the activated complex stage is wherebonds are broken and formed as reactantsare changed into products.

Even though molecules may haveenough kinetic energy to supply the activa-tion energy, a reaction may still not takeplace. A collision between molecules mustnot only possess sufficient energy, but themolecules must also have the proper orien-tation during the collision. To appreciate theimportance of orientation, think of the dam-age caused when two cars collide. A head-on collision results in appreciable damagecompared to a sideswipe or indirect colli-sion. Similarly, the energy transferred whenmolecules collide can vary significantlydepending on how they collide. Experi-ments demonstrate that the rate of a chemi-cal reaction is often less than what isexpected at a particular temperature. Onemajor reason for this disparity is that manyof the collisions do not have the correct ori-entation to break the bonds of reactants.

Factors Affecting ReactionRates

The collision theory can be used toexamine those factors that affect reactionrates. The most obvious factor is tempera-ture. There is a direct relationship betweentemperature and kinetic energy. Kineticenergy is equal to 1⁄

2mv2 and is also a mea-

sure of temperature. By increasing the tem-perature at which a reaction occurs, thekinetic energy of the molecules alsoincreases. Recall from Chapter 9 the distri-bution of velocities at two different tem-peratures. Figure 12.2 shows that as thetemperature increases, there is a greaterproportion of molecules with higher veloc-ities, and therefore, more molecules withsufficient kinetic energy to overcome theactivation energy barrier. The higher the

temperature is, the greater the proportion ofparticles with higher kinetic energies.

It’s often taken for granted that increas-ing the temperature increases the reactionrate. Hot water is used for cleaning, elevatedtemperatures to cook and bake, and sparksand flames to ignite fuels. One general ruleof thumb regarding temperature and reactionrate is that for every increase of 10°C thereaction rate doubles. This rule is sometimesreferred to as the van’t Hoff rule. JacobusHendricus van’t Hoff (1852–1911) was oneof the foremost chemists of his time. Van’tHoff’s did work on thermochemistry, solu-tions, and organic chemistry. He received thefirst Nobel Prize awarded in chemistry in1901. Van’t Hoff’s rule can also be stated inreverse by saying that the reaction rate ishalved for every decrease of 10°C.

Every time an item is placed in therefrigerator we depend on lower tempera-tures to slow reaction rates to prevent foodspoilage. The effect of temperature on reac-tion rates is also illustrated by its impact onhuman survival. Normal body temperatureis 37°C. An increase of body temperature ofjust a few degrees to produce a fever condi-tion increases the metabolic rate, while low-ering the body temperature slows downmetabolic processes. The slowing of human

Figure 12.2As the temperature increases, a greater pro-portion of particles reach higher velocitiesand therefore higher kinetic energies. (RaeDéjur)


metabolic processes explains why individu-als who have fallen through ice and beensubmerged under cold water for long peri-ods can sometimes be revived. In somecases, people have been under cold water for20 to 30 minutes and been revived with lit-tle or no permanent damage. Under normalconditions the brain requires a steady sup-ply of oxygen. If this supply is cut off for afew minutes, permanent brain damageresults. In cold icy waters, the need for oxy-gen is greatly reduced; therefore, someonemay be able to survive a lack of oxygen foran extended period. This phenomenon isalso applied in some surgical procedures,such as heart surgery, where the body tem-perature is lowered during an operation tomitigate reduced oxygen flow to the brain.

Another factor that affects the rate of achemical reaction is the concentration ofreactants. As noted, most reactions takeplace in solutions. It is expected that as theconcentration of reactants increases morecollisions occur. Therefore, increasing theconcentrations of one or more reactants gen-erally leads to an increase in reaction rate.The dependence of reaction rate on concen-tration of a reactant is determined experi-mentally. A series of experiments is usuallyconducted in which the concentration of onereactant is changed while the other reactantis held constant. By noting how fast thereaction takes place with different concen-trations of a reactant, it is often possible toderive an expression relating reaction rate toconcentration. This expression is known asthe rate law for the reaction.

The decomposition of hydrogen perox-ide illustrates the dependence of reactionrate on concentration. One way to determinehow fast hydrogen decomposes is to mea-sure the amount of oxygen generated. Byperforming a series of trials using differentconcentrations of hydrogen peroxide, it isfound that the reaction rate increases in

direct proportion to the original concentra-tion of H2O2. This is expressed mathemati-cally as

rate � k�H2O2

This equation is the rate law for the decom-position of hydrogen peroxide. The constantk is referred to as the rate constant and thebrackets indicate that the concentration ofhydrogen peroxide is expressed in molarunits. When the concentration is directlyproportional to the concentration of a reac-tant, the reaction is first order with respectto that reactant. Therefore, the decomposi-tion of H2O2 is first order with respect toH2O2. The reaction order tells us how thereaction rate is related to the concentration.As another example, consider the reactionof hydrogen and iodide to produce hydrogeniodide:

H2(g) � l2(g) 2Hl(g)

The rate law for this reaction has been deter-mined to be

rate � k�H2 �l2

Therefore, for this reaction, the rate is first order with respect to hydrogen and firstorder with respect to iodide. The over-all order of a reaction is found by summingthe individual orders. Therefore, the overallorder for the above reaction is 2.

The reaction order is equal to the expo-nent of the concentration in the rate law. Ageneral expression for the rate law can bewritten using the standard reaction: A � B

products, where A and B represent reac-tants. The general rate law for this expres-sion can be written as

rate � k�A x�B y

In this expression, A and B represent molarconcentrations of the reactants and x and y


are the orders with respect to each reactant.The overall order of the reaction would beequal to x � y. Reaction orders typicallyhave values like 0, 1, and 2, although it isnot unusual to have fractional reactionorders. A reaction order of 0 for a reactantmeans that the rate does not depend on theconcentration of that reactant. This makessense if it’s remembered that any value tothe 0 power is one. For example, if the reac-tion did not depend on the concentration ofreactant A, then �A 0 � 1 and the rate lawbecomes rate � k�B y.

As stated previously, the rate law for achemical reaction is determined experimen-tally. Because rates are not constant withtime but change during the course of thereaction, rates are measured during the ini-tial stages of the reaction. Once the rate lawfor a reaction is determined, an expressioncan be derived using calculus that expressesthe change in concentration of reactant withtime. One of the most widely used forms ofthese expresses the change in concentrationusing half-life. The half-life is simply thetime it takes for the original concentrationof a reactant to drop to half its originalvalue. A shorter half-life means that thereaction is faster. Half-life also is usedextensively in areas such as nuclear reac-tions and toxicology. Half-life is coveredmore extensively in Chapter 17.

A third important factor in determininghow fast a reaction occurs involves the sur-face area of the reactants. Because collisionstake place on the surface of reactants,increasing the surface area increases the rateof chemical reactions. The more a substanceis divided, the greater the amount ofexposed surface area (Figure 12.3).

Dividing a reactant into small piecesmakes a significance difference in the reac-tion rate. A familiar example is the differ-ence between igniting a log versus ignitingwood chips. One situation that illustrates

how the reaction rate depends on surfacearea is the occasional violent explosion ofgrain dust. Finely powdered grain is asexplosive as dynamite due to its tremendoussurface area. In grain elevators and storagebins, ventilation and other methods of dustcontrol are used to reduce the chances of aviolent explosion.

Reaction Mechanisms andCatalysis

Another method of increasing the reac-tion rate is by employing a catalyst. A cata-lyst is a substance that speeds up thereaction but is not consumed in the reaction.A substance that slows down or stops a reac-tion in known as an inhibitor. To under-stand how catalysts work and their role inreaction kinetics requires knowledge ofreaction mechanisms. A reaction mecha-nism is the series of reactions or stepsinvolved in the conversion of reactant to

Figure 12.3Dividing a substance into smaller piecesincreases the amount of exposed surfacearea. The volume of the large cube and thetotal volume of the eight small cubes are thesame, but the small cubes have twice thetotal surface area.


products. When a chemical reaction is writ-ten, the equation often represents the sum ofa series of reactions. For example, the deple-tion of ozone in the atmosphere can be rep-resented with the reaction

O3 � O 2O2

According to this reaction, ozone, O3, com-bines with atomic oxygen to produce molec-ular oxygen. One mechanism to explain thisreaction is

The reaction mechanism shows that whilethe sum of reactions 1 and 2 results in theoriginal reaction given for ozone depletion,the ozone does not react directly withatomic oxygen but with chlorine. Steps 1and 2 are called the elementary steps in themechanism. Summing the elementary stepsin a reaction mechanism gives the overall ornet reaction. Chemical reactions are gener-ally presented as the net reaction, and theelementary steps are typically omitted. It isimportant to remember that while the netreaction gives the reactants and products,the mechanisms show how the reactantsbecame products. We can think of the reac-tants and products as the start and destina-tion of a trip. We may start a trip in NewYork City and end in Los Angeles, but thereare numerous paths we could use to makeour trip. We might go directly from NewYork to L.A., but we just as easily couldhave stopped in St. Louis and Denver on theway. Similarly, reactants may go directly toproducts, but there may also be intermedi-ate reactions along the way.

The reaction mechanism shown forozone depletion includes chorine. Chlorinein this reaction acts as a catalyst. A principalsource of this chlorine is from the ultravio-let breakdown of CFC (chlorofluorocar-

bons) in the upper atmosphere. A catalystworks by altering the reaction mechanism sothat the activation energy of the reaction islowered. If the reaction mechanism is con-sidered as a path, a catalyst can be viewedas providing a shortcut around the activationenergy barrier. As an analogy, consider agroup of bicycle riders who need to climb asteep mountain. One path might be a roadstraight to the top, while another might be atunnel bored through the mountain. Only theheartiest cyclists may be able to make it overthe mountain by going up the road, butmany more will be able to get to the otherside by using the tunnel. To see how a cata-lyst works in a chemical reaction, reconsiderthe decomposition of hydrogen peroxide. Aspointed out at the start of this chapter,hydrogen peroxide slowly decomposes intowater and oxygen. Iodide serves as a cata-lyst for this reaction according to the fol-lowing mechanism:

In the decomposition of hydrogen peroxide,the iodide catalyst reacts in the first ele-mentary step, and it is regenerated in thesecond step.

In a reaction mechanism with two ormore steps, the slowest step will control therate of the net reaction. This step is referredto as the rate determining step. The ratedetermining step in a reaction mechanismcan be compared to the slowest step in aseries of activities. For example, say wewere mailing out letters and set up anassembly line of several people thatincluded the following tasks: 1) take enve-lope out of box, 2) place stamp on envelope,3) put letter in envelope, 4) address enve-lope, and 5) seal envelope. All the stepsexcept step 4 could be done in a matter ofseconds. It might take a minute or two to

Cl � O3 r ClO � O2 1� ClO � O r Cl � O2 2

O3 � O r 2O2

H2O2(aq) � r H2O(l) �� l Ol(aq) (aq)

� H2O2(aq) � r H2O1 � O2(q) �� Ol l(aq) (aq)

2H2O2(aq) r 2H2O(l) � O2(g)


address the envelope, and this step essen-tially controls how fast the letters can beprepared. In a similar fashion, the slowestreaction in a series of reactions making upthe mechanism controls the overall rate ofthe reaction.

The importance of catalysts in chemicalreactions cannot be overestimated. In thedestruction of ozone previously mentioned,chlorine serves as a catalyst. Because of itsdetrimental effect to the environment, CFCsand other chlorine compounds have beenbanned internationally. Nearly every indus-trial chemical process is associated withnumerous catalysts. These catalysts makethe reactions commercially feasible, andchemists are continually searching for newcatalysts. Some examples of important cat-alysts include iron, potassium oxide, andaluminum oxide in the Haber process tomanufacture ammonia; platinum andrhodium in the Ostwald synthesis of nitric

acid; and nickel when vegetable oil is hydro-genated to produce saturated fats such asmargarine. Catalytic converters in automo-biles rely on platinum, rhodium, and palla-dium as catalysts. The catalytic convertersaid in the complete combustion of emissionsfrom engines by converting carbon monox-ide to carbon dioxide, nitric oxides intonitrogen and oxygen, and hydrocarbons intocarbon dioxide and water. Because catalyticconverters must perform both oxidation andreduction (see Chapter 14 on electrochem-istry), catalytic converters generally work intwo stages. The hot exhaust first passesthrough a bed of rhodium catalyst to reducenitric oxides and then through a platinumcatalyst to oxidize CO and hydrocarbons(Figure 12.4).

Enzymes are catalysts that acceleratemetabolic processes in organisms. The word“enzyme” comes originally from the Greekterm “enzymos” for leavening. Enzymes are

Figure 12.4Catalytic Converter (Rae Déjur)


a critical ingredient in bread, wine, yogurt,and beer. Without enzymes, the biochemi-cal processes would be much too slow.Enzymes have the ability to speed chemicalreactions tremendously. It is not unusual forenzymes to increase the reaction rate by afactor of a million. Like any other catalyst,enzymes work by lowering the activationenergy. Enzymes, though, are very specificwith respect to the compound they interactwith and the reactions they catalyze. Thecompound associated with a particularenzyme is known as the substrate. A gen-eral model to explain how enzymes work isthe lock-and-key model (Figure 12.5). Inthis model, the enzyme, which is typically alarge protein molecule, has a “keyhole” orregion known as an active site. The activesite has the specific shape and chemicalcharacteristics to act on a particular sub-strate. The substrate enters the active sitejust as a key would enter a keyhole. Theenzyme-substrate forms a complex wherethe chemical reaction takes place. Theenzyme catalyzes the substrate into theproduct and then the product separates from

the enzyme. The human body contains thou-sands of enzymes involved in all biochemi-cal processes such as digestion, respiration,and reproduction.

Enzymes in humans work best at tem-peratures of 37°C. When temperaturesclimb too high, the efficiency of an enzymecan be greatly reduced. This is one examplewhere an increase in temperature can retardthe reaction rate. Another problem is thatcertain substances can disrupt enzymes byblocking active sites and preventing the sub-strate from bonding with the enzyme. Sub-stances that disrupt enzymes are known asinhibitors. Many poisons and drugs fit inthis category.

Chemical EquilibriumIn the discussion of reactions in Chapter

5, all reactions were written as completereactions. Complete reactions are writtenwith a single arrow pointing to the right (r),indicating reactants are converted into prod-ucts. For complete reactions, reactants areconverted into products until one of thereactants disappears. Many reactions areactually reversible reactions. Reversiblereactions are written with a double arrow(p or ). Reversible reactions actuallyconsist of two reactions called the forwardreaction and the reverse reaction. The for-ward reaction represents the conversion ofreactants into products, while the reversereaction represents the conversion of prod-ucts back to reactants. The reaction ofhydrogen and nitrogen to form ammonia isa reversible reaction:

N2(g) � 3H2(g) 2NH3(g)

If it’s assumed that there is no ammonia whenthe reaction starts, then the reactants initiallycombine according to the forward reaction to produce ammonia. As the ammonia is

Figure 12.5Lock-and-Key Model of Enzyme (RaeDéjur)


formed, some of this product decomposesback into hydrogen and nitrogen as indicatedby the reverse reaction. Eventually, a point isreached where the rate of the formation ofammonia is equal to the rate of decomposi-tion of ammonia. It is also true that bothnitrogen and hydrogen are forming and dis-appearing at a constant rate. At this point, theconcentrations of reactants and products donot change and the reaction has reached equi-librium. Figure 12.6 shows how the concen-tration of nitrogen, hydrogen, and ammoniachange with time. According to the reaction,1 mole of nitrogen reacts with 3 moles ofhydrogen to produce 2 moles of ammonia.Therefore, hydrogen disappears three timesas fast as nitrogen, while ammonia is formedat twice the rate that nitrogen disappears.

Whenever equilibrium is considered,certain assumptions concerning the reactionare made. One assumption is that the reac-tion takes place in a closed system, that is, itis assumed reactants are mixed in a closedcontainer. It is also assumed that the reactiontakes place under constant temperature andpressure conditions. When conditions underwhich the reaction takes place change, theequilbium is affected. This topic is addressedshortly, but first the equilbrium state isexamined.

It is important to realize that once equi-librium is reached the reaction does not

stop. At equilibrium, reactants continue tochange into product as products are con-verted back into reactants. As long as thereaction is at equilibrium, the rates of theforward and reverse reactions are equal andthe concentrations remain constant. Equil-brium can be compared to a departmentstore during the day of a big sale. We canconsider the store to be empty or containonly a few people before it opens. Similarly,before a reaction starts we can assume thereare little or no products. When the storeopens, people rush in and start to fill thestore; very few, if any, people exit the storeat this point. This corresponds to the initialstage of a reaction when the concentrationof reactants is at its highest. As the day pro-gresses, people continue to enter, but now ata slower rate than when the store opened.The number of people exiting also increasesas the day progresses. Eventually, the rate atwhich people enter and exit is equal. At thispoint, the number of people in the storeremains constant and we can consider thisto be equilibrium. This equilibrium condi-tion exists until closing time approaches.Now the rate of people exiting the storeincreases and the rate of people enteringdecreases. Finally, at closing time the storeis nearly empty again. Equilibrium wasestablished in the store once the number ofpeople entering equaled the number of peo-ple exiting. Equilibrium exists in the storeuntil some condition or event changes theequilibrium. In our example, we can con-sider the store to be initially at equilibriumwith only a few salepeople (product) andlots of people outside the store (reactants).The original condition exists until the storeopens, and then a new equilibrium point isestablished that is maintained throughoutmost of the day. Equilibrium again changesas the closing hour approaches. Other con-ditions during the day may cause the equi-librium in the store to shift. For example, if

Figure 12.6Change in H2, N2, NH3 (Rae Déjur)


a fire alarm sounded, the store wouldquickly empty or there may be an increasein shoppers around lunch-time. Chemicalreactions are similar to our store with theforward and reverse reactions representingshoppers entering and exiting. Equilibriummust always be considered with respect tothe existing conditions. If these conditionschange, then the equilibrium shifts until anew equilibrium position is reached.

Le Châtelier’s PrincipleThe equilibrium position of a chemical

reaction changes when the conditions underwhich the reaction takes place change.These conditions include adding or remov-ing reactants or products, changing the tem-perature, and changing the pressure. Achange in any one of these or several ofthese simultaneously shifts the equilibrium.A change in conditions can favor the for-ward reaction or reverse reaction. In the for-mer case, more reactants form into products,while the latter causes more products tochange to reactants. When the equilibriumfavors the forward reaction, the reactionshifts to the “right,” and when the reversereaction is favored, the equilibrium shifts tothe “left.” The terms “right” and “left” areoften used to designate how the equilibriumshifts.

How a chemical system at equilibriumchanges when conditions change was firststated by Henri Louis Le Châtelier(1850–1936) in 1884. Le Châtelier was aprofessor at a mining school in France whoworked on both the theoretical and practicalaspects of chemistry. His research on thechemistry of cements led him to formulate aprinciple to predict how changing the pres-sure affected a chemical system. In the pub-lication Annals of Mines in 1888, LeChâtelier stated the principle that bears hisname: “Every change of one of the factors

of an equilibrium occasions a rearrangementof the system in such a direction that the fac-tor in question experiences a change insense opposite to the original change.”According to Le Châtelier’s Principle,changing the conditions of a chemical sys-tem at equilibrium places a stress on the sys-tem. A stress is just a change in conditionssuch as temperature or pressure. The equi-librium shifts to the right or left to relievethe stress placed on the system. To see howLe Châtelier’s Principle works, consider thereaction used to prepare hydrogen gas:

C(s) � H2O(g) CO(g) � H2(g) �H � �131 kJ

Let’s assume this reaction is at equilib-rium and use Le Châtelier’s Principle to pre-dict what happens when we change theconcentration, pressure, and temperature ofthe system. Le Châtelier’s Principle says thatwhen a stress is placed on the system, theequilibrium shifts to relieve the stress. Oneway to stress the system is by adding orremoving reactants or products, that is,changing the concentration of reactants orproducts. In the hydrogen gas reaction ifmore steam (H2O) was added, the reactionwould shift to the right producing morehydrogen and carbon monoxide. The reac-tion shifts to cause more hydrogen to react.This shift counteracts the stress of addingmore hydrogen to the system. Likewise, byremoving hydrogen and carbon monoxide,the equilibrium shifts to the right, or if moreof these products are added, the reactionshifts to the left.

Another way to stress a chemical sys-tem is by changing the pressure. What if thepressure of the system increased in theabove reaction? How might this stress berelieved? The key to this lies in the stoi-chiometry of the reaction. The reactionshows 1 mole of solid carbon reacting with1 mole of steam to yield 1 mole of carbon


monoxide and 1 mole of hydrogen. Tounderstand the effect of pressure on equi-librium, remember that pressure is the forceper unit area, and this force results from thecollisions of gas particles with the containerwalls. The reaction shows that there is 1mole of gas on the left side of the equationand two moles of gas on the right side.Therefore, when the pressure is increased,the reaction shifts toward the side with thefewer number of moles of gas. This meansthe reaction shifts to the left. Each time acarbon monoxide molecule combines witha hydrogen molecule to produce a watermolecule and solid carbon, there is a loss ofone gas molecule. A decrease in the numberof gas molecules means there are fewer col-lisions, and fewer collisions translates intoa pressure decrease. Using the same reason-ing, decreasing the pressure on the systemwhen it is at equilibrium shifts the reactionto the right.

The effect of pressure on equilibriumdepends on the number of moles of gas par-ticles on the right and left sides of the bal-anced equation. In reactions where thenumbers of moles of gas are equal, a changein pressure has no effect on the equilibrium.Changes in pressure are directly related tochanges in volume. This is because adecrease in volume gives the same result asincreasing the pressure, while an increase involume is the same as a decrease in pres-sure. This situation exists when equilibriumoccurs in a cylinder with a movable piston,and the volume is changed by moving thepiston.

Pure solids and liquids do not need tobe considered when using Le Châtelier’sPrinciple, because the concentrations ofpure solids and liquids are constant duringthe course of the reaction. For example, con-sider the solid carbon in the hydrogen gasreaction. Figure 12.7 shows two situationswith different amounts of carbon. Assume

the system is at equilibrium with a smallamount, say 12 g, of carbon. The concentra-tion of carbon can be calculated using thisinformation. The molar concentration of asubstance is the number of moles of thatsubstance divided by the volume it occupies.Twelve grams of carbon are equivalent to 1mole of carbon. The volume occupied by thecarbon can be found by dividing the mass,12 g, by its density, 2.6 g/cm3, and convert-ing to liters. This gives a volume of 0.0046L. The molar concentration of carbon isfound by dividing 1 mole by 0.0046 L,which equals 217 M. If more carbon isadded so that the amount of carbon in thesystem is 120 g, then there are now 10moles of carbon in the container. The vol-ume of carbon is also ten times greater.Therefore, the molar concentration remainsconstant at 217 M. Whenever pure solidsand liquids are present in the system, theydo not need to be considered in the equilib-rium. Changing the amount of hydrogen inour reaction shifts the reaction because theconcentration of hydrogen changes. Addingor removing carbon has no effect on theequilibrium because the concentration ofcarbon does not change.

Figure 12.7When more solid carbon is added to the sys-tem at equilibrium, the equilibrium is notaffected. The concentration of carbonremains constant.


The final stress to be considered is achange in temperature. To apply Le Châte-lier’s Principle with temperature changes,the sign of �H for the reaction needs to beknown. The �H in our example is � �131kilojoules. This indicates that the forwardreaction is endothermic and the reversereaction is exothermic. When the tempera-ture of a system at equilibrium is increased,the equilibrium will favor the endothermicreaction. One way to think of the effect oftemperature is to think of energy as a reac-tant or product. This is seen when the for-ward and reverse reactions are written astwo separate reactions:

Forward Reaction131 kJ � C(s) � H2O(g) CO(g) � H2(g)

Endothermic

Reverse ReactionCO(g) � H2(g) C(s) � H2O(g) � 131 kJ

Exothermic

In an endothermic reaction, energy can beconsidered a reactant, and in an exothermicreaction, energy can be considered a prod-uct. When the temperature of a system atequilibrium is increased, energy flows intothe system. To relieve this stress, energy isused in the endothermic reaction to estab-lish a new equilibrium. The effect of addingenergy is no different from the effect ofadding another reactant, for example, steam.Decreasing the temperature of the systemremoves energy from the system; to coun-teract this stress, the exothermic reaction isfavored.

The Equilibrium ConstantTwo Norwegians named Cato Maximil-

ian Guldberg (1836–1902) and Peter Waage(1833–1900) carried out a comprehensivestudy of equilibrium between 1864 and1879. Their work led to what is known as

the law of mass action. For a generalreversible reaction given by the equation: aA� bB cC � dD, the law of mass actionstates an equilibrium constant, Keq, can bedefined at a particular temperature:

The small letters are the coefficients in thebalanced equation and the brackets indicatethe molar concentrations (moles per liter) ofthe reactants and products. Applying the lawof mass action to the hydrogen gas equationgives

Notice that solid carbon does not appear inthe reaction. Again, this is because the con-centration of solid carbon is constant and isassumed to be part of Keq. The value for Keq

is approximately 2 at 1,000°C.Values for equilibrium constants have

been tabulated for numerous reactionsoccurring at various temperatures. Becausemany of these constants are associated withindustrial processes, the temperature atwhich a Keq is reported varies considerably.That is, it might seem odd that a Keq valuefor a reaction is given at what seems like anarbitrary temperature, but it should beremembered that these constants have prob-ably been determined with regard to a spe-cific industrial process, for example,1,000°C for the production of hydrogen gasfrom steam and coke (carbon).

Values for equilibrium constants varywidely. For example, consider the followingtwo reactions:

N2(g) � O2(g) 2NO(g)

Keq � 1.0 � 10�31 @ 25°C

H2(g) � Cl2(g) 2HCl(g)

Keq � 1.8 � 1033 @ 25°C

c d[C] [D]K �eq a b[A] [B]

[CO][H ]2K �eq [H O]2


The value of an equilibrium constant indi-cates whether reactants or products domi-nate at equilibrium. The small Keq value forthe first reaction above shows that, for allpractical purposes, nitrogen and oxygen donot react to form nitric oxide, NO, at 25°C.Conversely, hydrogen and chloride are quitereactive and react almost completely at25°C. Remember that the equilibrium con-stant is reported at a specific temperature. Itis known that nitrogen and oxygen do notreact at room temperature from the fact thatthe reaction is not occurring while you sithere and read these words. Nitric oxide,though, is a significant air pollutant that is aby-product of automobile exhaust andpower plant combustion. The higher tem-peratures associated with fuel combustionshifts the equilibrium to the right andincreases the Keq value for the formation ofnitric oxide. The value of Keq for this reac-tion is approximately 1.7 � 10�3 at2,000°C.

Many different types of reversible reac-tions exist in chemistry, and for each ofthese an equilibrium constant can bedefined. The basic principles of this chapterapply to all equilibrium constants. The dif-ferent types of equilibrium are generallydenoted using an appropriate subscript. Theequilibrium constant for general solutionreactions is signified as Keq or Kc, where thec indicates equilibrium concentrations areused in the law of mass action. When reac-tions involve gases, partial pressures areoften used instead of concentrations, and theequilibrium constant is reported as Kp (pindicates that the constant is based on par-tial pressures). Ka and Kb are used for equi-libria associated with acids and bases,respectively. The equilibrium of water withthe hydrogen and hydroxide ions isexpressed as Kw. The equilibrium constantused with the solubility of ionic compoundsis Ksp. Several of these different K expres-

sions will be used in chapters to come, espe-cially in the next chapter on acids and bases.

Ammonia and the HaberProcess

Similar to chemical kinetics, chemicalequilibrium plays a critical role in all areasof chemistry. The chemical industry con-stantly uses equilibrium principles toincrease product yield, design efficientchemical processes, and ultimately increaseprofits. One example that illustrates howequilibrium and Le Châtelier’s Principleapply to the chemical industry is in the pro-duction of ammonia. Ammonia is widelyused in fertilizers, refrigerants, explosivesand cleaning agents and to produce otherchemicals, such as nitric acid. Approxi-mately 18 million tons of ammonia are man-ufactured annually in the United States, andeach year it ranks as one of the top tenchemicals produced. During ancient times,ammonia was produced by the distillation ofanimal dung, hooves, and horn. In fact, thename is derived from the ancient Egyptiandiety Amun who was known to the Greeksas Ammon. Greeks observed the preparationof ammonia from dung in Egypt from salammoniac (ammonium chloride) near thetemple of Amun. Sal ammoniac or “salt ofAmmon” is an ammonium-bearing mineral(ammonium chloride) of volcanic origin.

The importance of nitrogen fertilizers inagriculture was established during the mid-1800s, and this coupled with the growth ofthe chemical industry, provided incentive tofind a method for fixing nitrogen. Nitrogenfixation is a general term to describe theconversion of atmospheric nitrogen, N2, intoa form that can be used by plants. Onemethod to fix nitrogen mimicked the natu-ral fixation of nitrogen by lightning. Theprocess involved subjecting air to a highvoltage electric arc to produce nitric oxide:


N2(g) � O2(g) 2NO(g)

The nitric oxide was then converted intonitric acid and combined with limestone toproduce calcium nitrate. The problem withthis process was that the procedure wasenergy intensive, and it was only economi-cal where there was a steady and cheap sup-ply of electricity.

Another way to fix nitrogen is by thesynthesis of ammonia. The conversion ofnitrogen and hydrogen into ammonia hadbeen studied since the mid-1800s, but seri-ous work on the subject did not occur untilthe turn of the century. In 1901, Le Châtelierattempted to produce ammonia by subject-ing a mixture of nitrogen and hydrogen to apressure of 200 atmospheres and a temper-ature of 600°C in a heavy steel bomb. Con-tamination with oxygen led to a violentexplosion, and Le Châtelier abandoned hisattempt to synthesize ammonia. Toward theend of his life, Le Châtelier was quoted assaying “I let the discovery of the ammoniasynthesis slip through my hands. It was thegreatest blunder of my scientific career.”

By 1900, the increasing use of nitratefertilizers to boost crop production providedample motivation for finding a solution tothe nitrogen fixation problem. During thelast half of the nineteenth century and theearly twentieth century, the world’s nitratesupply came almost exclusively from theAtacama Desert region of northern Chile.This area contained rich deposits of salt-peter, NaNO3, and other minerals. The Warof the Pacific (1879–1883) between Chileand Bolivia and its ally Peru was foughtlargely over control of the nitrate-producingregion. Chile’s victory in the war (some-times referred to as the “nitrate war”) forcedBolivia to cede the nitrate region to Chile,and as a result, Bolivia not only lost thenitrate-rich lands but also became land-locked.

After Le Châtelier’s failed attempt tosynthesize ammonia, the problem was tack-led by the German Fritz Haber (1868–1934). As early as 1905, Haber’s progress onthe problem indicated it might be possibleto make the process commercially feasible.Haber continued to work throughout thenext decade and by 1913, on the eve ofWorld War I, had solved the problem. Theimplementation of Haber’s work into a com-mercial process was carried out by KarlBosch (1874–1940) and became known as the Haber or Haber-Bosch process.Haber’s work was timely because Germanyacquired a source of fixed nitrogen to pro-duce nitrate for fertilizer and explosives tosustain its war effort. Without this source, ablockade of Chilean nitrate exports wouldhave hampered German’s war effortseverely. Haber received the Nobel Prize inchemistry in 1918 for his discovery of theprocess that bears his name. Karl Boschreceived the Nobel Prize in 1931 primarilyin conjunction with his work on the produc-tion of gasoline.

The Haber process for the synthesisillustrates several concepts presented in thischapter. The reaction is represented as fol-lows:

N2(g) � 3H2(g) 2NH3(g) �H � –92.4 kJ

According to Le Châtelier’s Principle,the production of ammonia is favored by ahigh pressure and a low temperature. TheHaber process is typically carried out atpressures between 200 and 400 atmospheresand temperatures of 500°C. While LeChâtelier’s Principle makes it clear why ahigh pressure would be favorable in theHaber process, it is unclear why a high tem-perature would be desirable because thereaction is exothermic. An increase in tem-perature shifts an exothermic reaction to theleft. Even though the equilibrium shifts to


the left at high temperatures, the reaction istoo slow at lower temperatures. This demon-strates the trade-off between equilibriumand kinetics in this reaction. It is fruitless tokeep the temperature low to produce ahigher yield if it takes too long. For instance,it would make more sense to accept a 10%yield if you could get this in 30 minutes asopposed to a 60% yield that took 12 hours.In the commercial production of ammonia,NH3 is continually removed as it is pro-duced. Removing the product causes more

nitrogen and hydrogen to combine accord-ing to Le Chatelier’s Principle. Unreacted,nitrogen and hydrogen are separated fromthe product and reused.

An important aspect of the Haber pro-cess is the use of several catalysts in the reac-tion. A major problem in Haber’s attempt tosynthesis ammonia was identification of suit-able catalysts. Haber discovered that ironoxides worked as the primary catalysts incombination with smaller amounts of oxidesof potassium and aluminum.

Introduction and HistoryTwo of the most important classes of chem-ical compounds are acids and bases. A smallsampling of acids and bases found aroundthe home demonstrates their importance indaily life. A few of these include fruit juice,aspirin, milk, ammonia, baking soda, vine-gar, and soap. Beyond their presence innumerous household items, acids and basesare key ingredients in the chemical processindustry. More sulfuric acid is producedthan any other chemical in the United Stateswith an annual production of 40 milliontons. While the commercial applications ofacids and bases illustrate their importancein everyday life, on a more fundamentallevel each one of us inherited our character-istics and genetic make-up through the acidDNA, deoxyribonucleic acid.

Human use of acids and bases datesback thousands of years. Probably the firstacid to be produced in large quantities wasacetic acid, HC2H3O2. Vinegar is a dilutedaqueous solution of acetic acid. This acid isan organic acid that forms when naturallyoccurring bacteria called acetobacter aceticonvert alcohol to acetic acid. AncientSumerians used wine to produce vinegar for

use in medicines and as a preservative. Asignificant advance in chemistry occurredaround the year 1200 when alchemists dis-covered how to prepare strong mineralacids. These acids include sulfuric, nitric,and hydrochloric acid. Mineral acids wereprepared by distilling various vitriols (sul-fate compounds) or virtriol mixtures. Dis-tilling green vitriol (FeSO4) and dissolvingthe vapor in water produced sulfuric acid,formerly called oil of vitriol. The distillationof vitriol and saltpeter (KNO3) resulted innitric acid. Aqua regia or royal water con-sists of a mixture of one part nitric acid andthree parts hydrochloric acid. The nameaqua regia denotes the ability of this mix-ture to dissolve precious metals such asgold. The word “acid” comes from the Latinword “acere,” which means sour.

The use of bases or alkalines also datesback thousands of years. Bases were nodoubt created as prehistoric humans carriedout their daily activities. Bases are a keyingredient of soap; some of the first soaprecipes date back to 2800 B.C. from theBabylonian period. The Egyptians combinedlime (calcium oxide) and soda ash (sodiumcarbonate) and evaporated the product to

13

Acids and Bases

156 Acids and Bases

produce caustic soda (NaOH). They usedthis base to produce cleansers and dyes formaterials and in the preparation of papyrus.Bases were used in ancient China to preparepaper. The use of lye (NaOH) to producesoap dates back to at least Roman times, andit is still used today to prepare soap. Alkalineis derived from the Arabic “al qualy,” whichmeans to roast. Evidently the term referredto roasting or calcinating plant material andleaching the ash residue to prepare a basiccarbonate solution.

Acids were of critical concern toLavoisier as he gathered evidence disputingthe phlogiston theory. In fact, Lavoisier mis-takenly thought that oxygen was a commonchemical in all acids and adopted the term“oxygen,” which means “acid former,” for thiselement. Lavosier’s oxygen theory of acidspersisted for two decades. Chemical analysesconducted by Humphrey Davy and severalother prominent chemists around 1810demonstrated that oxygen was not present inmany acids, for example, hydrochloric acid(HCl). This led Davy to abandon the oxygentheory of acids and suggest it was hydrogenthat characterized acids. Most chemists had

abandoned Lavoisier’s oxygen acid theory andaccepted Davy’s hydrogen theory by 1820. In1838, Justus Liebig (1803–1873) advancedDavy’s concept of an acid and defined an acidas a compound that contains hydrogen thatcan be replaced by a metal. While chemiststhroughout the 1800s continued to improvetheir description of an acid, they were unableto furnish an adequate definition of a base. Forthe most part, a base was considered a sub-stance that neutralized an acid.

It was not until the last decade of thenineteenth century that chemists had an ade-quate theoretical description of acid andbases. Until then, most acids and bases wereclassified according to their general proper-ties. Chemists knew acids and bases dis-played contrary properties, and these wereadequate for identifying many acids andbases. Some of these properties are listed inTable 13.1.

Chemical Definitions ofAcids and Bases

A general chemical definition for acidsand bases was proposed by Svante Arrhe-

Table 13.1General Properties of Acids and Bases

Acids and Bases 157

nius in 1887. Arrhenius defined an acid as asubstance that dissociates in water to givehydrogen ions (H�) and defined a base as asubstance that dissociates in water to givehydroxide ions (OH�). A hydrogen ion issimply a hydrogen atom minus its electron.Because a hydrogen atom consists of a sin-gle proton and a single electron, removingan electron leaves just the proton. Therefore,a hydrogen ion is equivalent to a proton, andboth can be symbolized as H�. If the gen-eral formula of an acid is represented as HAand a base as BOH, Arrhenius’ definitionsfor an acid and base can be represented bythe following general reactions:

Acid HA � H2O H�(aq) � A�

(aq)

Base BOH � H2O OH�(aq) � B�

(aq)

When an acid dissociates to produce hydro-gen ions in water, the hydrogen ions do notremain as individual ions but are attractedto the polar water molecules represented bythe following reaction:

H� � H2O H3O�

The H3O� ion is called the hydronium ion.

To be technically correct, Arrhenius’ defi-nition for an acid should state an acid is asubstance that produces hydronium ions insolution. Several water molecules may actu-ally be associated with a single hydrogen ionto produce ions such as H5O2

� or H7O3� in

acids. The important thing to remember isthat hydrogen ions do not exist as individualions in water, but become hydrated. Whileyou will often see the symbol H� in reac-tions involving acids, the H� should beinterpreted as a hydronium ion rather than ahydrogen ion. Throughout this chapterhydrogen ions and hydronium ions shouldbe considered synonymous.

The Arrhenius definition implies thatacids contain hydrogen ions and bases con-

tain the hydroxide ion. This is illustrated inthe chemical formulas for some commonacids and bases displayed in Table 13.2.

Table 13.2 illustrates the presence ofhydrogen in acids. It is also apparent thatbases contain hydroxide ions, but the weakbase ammonia seems to be an exception.Ammonia illustrates one of the short-comings of the Arrhenius definition of acidsand bases; specifically, bases do not have tocontain the hydroxide ion to producehydroxide in aqueous solution. Whenammonia dissolves in water, the reaction isrepresented by:

NH3(g) � H2O(l) NH4�

(aq) � OH�(aq)

Table 13.2Some Common Acids and Bases

158 Acids and Bases

This reaction shows that the hydroxide ionscome from ammonia pulling a hydrogenaway from water resulting in the formationof OH�. Therefore, a compound does nothave to contain hydroxide to be a base. Inaddition to this limitation, the Arrheniusdefinition limits acids and bases to aqueoussolutions.

Recognizing these limitations, the Dan-ish chemist Johannes Brønsted (1879–1947)and the English chemist Thomas Lowry(1874–1936) independently proposed a newdefinition for acids and bases in 1923. Brønsted and Lowry noted that in certainreactions substances acted like acids eventhough hydrogen ions were not present.Brønsted and Lowry viewed acids and basesnot as isolated substances but as substancesthat interacted together as a pair. Accordingto the definition, an acid is a proton donorand a base is a proton acceptor. An anal-ogy helps to explain the idea behind theBrønsted-Lowry theory. Acids and basescan be compared to throwing a football. Aquarterback throws a football to a receiver.In this example, the quarterback representsthe acid, the football the proton, and thereceiver the base. Applying the Brønsted-Lowry definition to nitric acid, it is seen thatHNO3 donates a proton to water, Water inthis case acts as a base.

When ammonia dissolves in water, wateracts as an acid when it donates a proton tothe base ammonia:

Just as a quarterback and receiver act as apair, Brønsted-Lowry acid and bases act aspairs (Figure 13.1). Acids and bases are

defined in terms of reactions and the inter-action of two substances, rather than as indi-vidual substances. In the two examples it isseen that water, which is not commonly con-sidered an acid or base, is an acid when itreacts with ammonia and a base when itreacts with nitric acid.

To complete our discussion of the Brønsted-Lowry theory, we see that when aproton is transferred, two new species areformed. In the nitric acid reaction, water istransformed into the hydronium ion andnitric acid into the nitrate ion, NO3

�. Thehydronium and nitrate ions formed arethemselves an acid and base, respectfully.Again, using our quarterback and receiverexample, once the receiver has the football,he or she can act as an acid and pass it on tosomeone else. Likewise, the quarterbackcan act as a base by receiving the football.The HNO3–NO3

� and H2O–H3O� are

referred to as conjugate acid-base pairs.Notice that the only difference between thesubstances making up a conjugate acid-basepair is the presence of a proton. In theammonia reaction, the conjugate acid basepairs are NH3(base)–NH4

� (acid) and H2O(acid)–OH� (base).

In the same year that Brønsted andLowry proposed their definition for acidsand bases, the American G. N. Lewis pro-posed an alternative definition based on the

Acid Base� �HNO � H O r H O � NO3(aq) 2 (l) 3 (aq) 3

| H� H

Base Acid� �NH � H O r NH � OH3(g) 2 (l) 4 (aq) (aq)

H� |H

Figure 13.1Football Analogy (Rae Déjur)

Acids and Bases 159

valence electron structure of substances.Brønsted-Lowry bases must contain at leastone unshared electron pair to accept a pro-ton. We can see this if we look at the Lewisdot structure of several bases:

Lewis defined a base as an electron pairdonor and an acid as an electron pair accep-tor. Lewis’ electron pair donor was the sameas Brønsted-Lowry’s proton acceptor, andtherefore, was an equivalent way of defininga base. Lewis’ acids were defined as a sub-stance with an empty valence shell thatcould accommodate a pair of electrons. Thisdefinition broadened the Brønsted-Lowrydefinition of an acid. The three definitionsof acids and bases are summarized in Table13.3.

Each of the three definitions expandsour concept of acids and bases. Arrhenius’basic definition is adequate for understand-ing many of the properties of acids andbases. It is important to recognize, though,that acids and bases are not fixed labels thatcan be applied to a substance. Brønsted-Lowry and Lewis showed that acid-basecharacteristics are dependent on the reac-tions that take place between substances. A

substance can act as an acid or base depend-ing on the conditions. When a substance hasthe ability to act as an acid or a base, it iscalled amphoteric. We have already seenthat water is an amphoteric substance. Fur-thermore, substances not typically classifiedas acids or bases, such as water, act as acidsand bases in many instances. Throughoutthe remainder of this chapter, the Arrhenius’definition will be used. This simple defi-nition, though it has limitations, is useful forunderstanding most of the basic concepts ofacids and bases.

Strengths of Acids andBases

Many of us start our day by consumingacid in the form of fruit juice and washingourselves with a base-like soap. We are alsoaware that contact with some acids andbases causes severe burns. Our personalexperiences with different acids and basespoint out the wide differences in thestrengths of these compounds. The strengthof acids and bases is a function of their abil-ity to ionize. Strong acids and bases can beconsidered 100% ionized or completely dis-sociated. Weak acids and bases only partlyionize in water. We can apply our conceptsof equilibrium from Chapter 12 to differen-tiate between strong and weak acids orbases. For example, nitric acid is a strongacid so HNO3 dissociates completely intoH� and NO3

�. When we represent the dis-sociation of HNO3 in aqueous solution withthe equation HNO3(aq) � H2O(l) H3O

�(aq)

� NO3�

(aq), the equation is written with asingle arrow pointing to the right to indicate100% ionization (complete dissociation).On the other hand, a weak acid such asacetic acid dissociates very little. Only about0.4% of a 1 M acetic acid solution dissoci-ates into ions, that is, most of the acetic acidremains in the combined form of HC2H3O2

Table 13.3Definitions of Acids and Bases

160 Acids and Bases

in solution. Similarly, in an aqueous solutionof the weak base NH3 the concentrations ofNH4

� and OH� ions are less than 1% of thatof the concentration of NH3. Equations rep-resenting the dissociation of weak acids orbases in solution are written with a doublearrow to indicate that equilibrium existsbetween the acid or base and its ions. Theequilibrium condition for weak acids orbases is such that the equilibrium favors thecombined states of acids or bases rather thantheir dissociated ion state:

weak acid or weak base partially dissociated ions

strong acid or base total dissociation

A quantitative measure of the degree ofdissociation is given by the equilibrium con-stant for the acid or base. The higher theequilibrium constant is, the greater the per-cent dissociation of the acid or base. There-fore, a higher equilibrium constant means astronger acid or base. Equilibrium con-stants, Ka and Kb, are listed for several com-mon weak acids and bases in Table 13.4.

A final point concerning our discussionof acid/base strengths deals with thestrength of conjugate acid-base pairs. Whenan acid donates a proton, what remains is abase, and when a base accepts a proton, theproduct is an acid. Whenever we considerBrønsted-Lowry acid-base pairs, a generalrule is that a strong acid or base produces acorresponding weak base or acid. Thismeans that strong acids produce weak bases,weak bases produce strong acids, and so on.To illustrate this point, let’s consider a strongacid (HCl), a weak acid (HC2H3O2), and aweak base (NH3). Equations for their solu-tion in water are

HCl(aq) � H2O(l) H3O�

(aq) � Cl�

HC2H3O2(aq) � H2O(l) H3O�

(aq) � C2H3O2�

NH3(aq) � H2O(l) NH4�

(aq) � OH�(aq)

In the first reaction, HCl donates a proton towater and the conjugate base of HCl is Cl�.Because HCl is a strong acid (indicated bythe one-way arrow), Cl� is a weak base. Inthe next reaction, acetic acid donates a pro-ton, creating the acetate ion, C2H3O2

�. Theacetate ion is the conjugate base of aceticacid and is a strong base, because it isformed from a weak acid. Finally, in theammonia reaction, the base NH3 accepts aproton to become the ammonium ion,NH4

�. Ammonium is the conjugate acid ofammonia. The ammonium ion is a strongacid, because it is produced from the weakbase ammonia. There is an inverse relation-ship between the strength of an acid or baseand the conjugate formed when a proton istransferred.

Acid ConcentrationIt is important not to confuse acid/base

strength, which depends on the degree ofionization, with concentration. Concentra-tion measures how much acid or base is in

Table 13.4Ka and Kb for Weak Acids and Bases at 25°C

Acids and Bases 161

solution, and it is typically measured inmoles per liter or molarity. For instance,acetic may have a high concentration, butbecause it is a weak acid most of it will existin solution as HC2H3O2. Thus, a 10 M aceticacid solution is highly concentrated, butonly a small fraction of the acid ionizes toproduce hydrogen ions; the concentration ofhydrogen ions is perhaps only 0.1 moles perliter. Conversely, a 10 M solution ofhydrochloric acid (HCl) ionizes 100% toproduce 10 moles each of hydrogen andchloride ions per liter.

Another concentration unit often asso-ciated with acids and bases is normality.Normality is similar to molarity, but ratherthan a measure of the number of moles perliter it measures equivalents per liter (N).For acids and bases, an equivalent can beconsidered the mass that produces 1 mole ofhydrogen or hydroxide ions, respectively. Tounderstand normality, we have to consideracids and bases that yield more than 1 moleof H� or OH� when they dissociate. Sulfu-ric acid (H2SO4), for example, will yield 2moles of hydrogen ions for every mole ofsulfuric acid that dissociates, and each moleof calcium hydroxide, Ca(OH)2, gives 2moles of hydroxides when it dissociates. Anequivalent is found by dividing the molarmass by the number of moles of hydrogenor hydroxide ions that the molecule pro-duces. For H2SO4, an equivalent would beequal to 49 grams because the molar massof sulfuric acid is 98 grams and each moleof sulfuric acid produces 2 moles of hydro-gen ions in solution. Thus, a 1 normal solu-tion of H2SO4 contains 49 grams of H2SO4

in 1 liter of solution. It can easily be shownthat a 1 M H2SO4 solution is the same as a 2N H2SO4 solution:

1 M H2SO4 contains 1 mole or 98 g H2SO4

per liter98 grams is the same as 2 equivalents of

H2SO4

so there are 2 equivalents of H2SO4 per literwhich � 2 N

Many acids such as HCl and NaOH yield amole of ions per every mole of acid or base.In these cases, the normality and molarityare equal.

Commercial acids and bases are sold instandard high concentrations and diluted tothe desired concentration. The standard con-centrated forms are referred to as stock solu-tions. Concentrations of stock solution aregiven in Table 13.5.

pH and pOHThe basic criterion for distinguishing

acids and bases in aqueous solutions is theconcentration of hydrogen and hydroxideions. When an acid dissolves in water, thewater acts as a base and accepts protons toform the hydronium ion. Water acts as anacid to produce hydroxide ions when itdonates protons to a base dissolved in water.While we can considered water to be ampho-teric when it interacts with other acids andbases, we can also consider pure water to beboth an acid and base. The reason for this isthat water itself actual dissociates to a very

Table 13.5Stock Solutions of Acids and Bases

162 Acids and Bases

small extent to produce hydrogen andhydroxide ions:

H2O(l) H�(aq) � OH�

(aq)

The above reaction depicts water as anArrhenius acid and base. Treating water interms of the Brønsted-Lowry theory, a moreappropriate reaction would be

H2O(l) � H2O(l) H3O�

(aq) � OH�(aq)

At 25°C the concentrations of hydroniumand hydroxide ions in water are both equalto 1.0 � 10�7 M. To put this concentrationinto perspective, consider that about two outof every billion water molecules dissociatesinto ions. The equilibrium constant for theabove reaction is known as the ion productconstant of water, symbolized by Kw, and isequal to the product of the H� and OH�

molar concentrations:

Kw � �H� �OH� �(1.0 � 10�7) (1.0 � 10�7) � 1.0 � 10�14

The equation for Kw applies to both purewater and aqueous solutions.

Because the ion concentrations aresmall and the negative exponents makethem tedious to work with, Soren Peer Lau-ritz Sorenson (1868–1939), a Danish bio-chemist, devised the pH concept in 1909 toexpress the hydrogen ion concentration. Theabbreviation pH comes from the Frenchpouvoir hydrogène meaning power ofhydrogen. The pH of a solution is given bythe equation:

pH � –log�H�

Remember that brackets are used to signifythe molar concentration of a substance. Thisequation states that pH is equal to the nega-tive log of the hydrogen ion molar concen-tration, but it should be remembered it isreally the hydronium ion concentration in

solution that is being measured. In a similarfashion, the pOH of a solution can bedefined as:

pOH � –log�OH�

Using these definitions, the pH andpOH of a neutral solution at 25°C are bothequal to 7. We can see from the expressionfor Kw that �H� and �OH� are inverselyrelated, and consequentially pH and pOHare inversely related. We can picture pH andpOH as sitting on opposite sides of a see-saw, as one goes up, the other always goesdown. The product of the hydrogen andhydroxide concentrations will be equal to1.0 � 10�14, while the sum of the pH andpOH will be equal to 14. In an acidic solu-tion, the hydrogen ion concentrationincreases above 1.0 � 10�7, the hydroxideconcentration decreases, and the pH valuegets smaller. The relationship between thetype of solution, pH, pOH, and ion concen-trations is shown in Table 13.6. The pHs ofa number of common substances are pre-sented in Table 13.7.

Because the pH and pOH scales arebased on logarithms, a change in 1 pH orpOH unit represents a change in ion con-centration of a factor of ten. Coffee with apH of 5 has approximately 100 times thehydronium ion concentration as tap waterwith a pH of 7.

The pH of a solution can be measuredusing several methods. One popular method

Table 13.6Solution, pH, pOH, and Ion Concentrations

Acids and Bases 163

is to use an indicator. An indicator is a sub-stance that “indicates” the condition ofanother substance to which the indicator hasbeen added. Indicators used to determinethe pH of a substance are known as acid-base indicators. Many common acid-baseindicators come from plant pigments. Lit-mus consists of blue and red dyes extractedfrom lichens originally found in the Nether-lands. While litmus paper only gives a gen-eral measure of pH by identifying asubstance as an acid or a base, there are hun-dreds of other indicators that can be used tomore accurately measure pH. Simple indi-cators can be obtained by boiling plantssuch as beets, cabbage, and spinach. Severalindicators are often mixed to produce a uni-versal indicator. Universal indicators showvarious colors and shades across a wide pHrange and can measure pH in the range from2–12. Numerous synthetic indicators havealso been produced. One of the most com-mon is phenolphthalein, which is colorlessin acidic solutions and turns pink in basicconditions.

The way indicators work can be under-stood if we consider indicators to be weakacids. If we let Hln represent the general for-mula of an indicator, then we can write thefollowing equilibrium expression:

Hln(aq) H�(aq) � ln�

(aq)

color different color

Indicators have different colors in the com-bined and dissociated forms. The equilib-rium of the indicator in solution shiftsaccording to Le Châtelier’s principle as anacid or base is added to the solution. Theshift in equilibrium to the right or left causesthe color to change accordingly.

Accurate and precise measures of pHcan be made with a pH meter. Typical pHmeters usually contain a glass electrode and reference electrode arranged similar toan electrochemical cell. We discuss electro-chemical cells in Chapter 14. For now,though, consider a pH meter as essentially amodified voltmeter in which the voltagemeasured is directly proportional to thehydrogen ion concentration of the solution.Simple pH meters are capable of measuringpH within � 0.1 pH units, while moresophisticated instruments are precise towithin 0.001 pH units.

Neutralization ReactionsWhen acidic and basic solutions con-

taining hydroxide are mixed, a neutraliza-tion reaction occurs in which the acid and

Table 13.7Approximate pH Values of Common Substances

164 Acids and Bases

base lose their characteristic properties.Acid and base solutions combine duringneutralization to form a salt and water. Asan example, consider the reaction betweenhydrochloric acid and sodium hydroxidesolutions:

HCl(aq) � NaOH(aq) H2O(l) � NaCl(aq)

For complete neutralization to take place,the proper amounts of acid and base must bepresent. The salt formed in the above reac-tion is NaCl. If the water were evaporatedafter completing the reaction, we would beleft with common table salt. Sodium chlo-ride is just one of hundreds of salts that formduring neutralization reactions. While wecommonly think of salt, NaCl, as a season-ing for food, in chemistry a salt is any ioniccompound containing a metal cation and anonmetal anion (excluding hydroxide andoxygen). Some examples of salts that resultfrom neutralization reactions include potas-sium chloride (KCl), calcium fluoride(CaF2), ammonium nitrate (NH4NO3), andsodium acetate (NaC2H3O2).

Neutralization reactions are very impor-tant in both nature and industry. Because pHhas a profound effect on many chemicalprocesses, neutralization can be used to con-trol the pH. A common application of neu-tralization reactions is in the use of antacids.As noted in Table 13.7, the pH of stomachacid is approximately 1.5. The hydrochloricacid present in gastric juices helps us digestfood and activates specific enzymes in thedigestive process. Our stomach and diges-tive tract are protected from the hydrochlo-ric acid secreted during digestion by aprotective mucous lining. Cells lining ourdigestive system constantly regenerate thisprotective layer. Stomach ulcers developwhen the protective mucous lining is weak-ened or excess stomach acid is generateddue to health problems. Antacids are a com-mon method used to neutralize excess stom-

ach acid. An antacid is nothing more than abase or salt that produces a basic solutionwhen ingested. Some popular antacids andtheir associated bases are shown in Table13.8.

The reactions in Table 13.8 show thatcarbon dioxide is a common product inmany neutralization reactions. This isclearly displayed when a drop of vinegar(acetic acid) is added to baking soda(sodium bicarbonate). Some aspirinincludes an antacid in their formulation toneutralize some of the acidity imparted bythe aspirin (acetylsalicylic acid). These arecommonly referred to as buffered aspirins.

A detrimental form of neutralizationoccurs when acid precipitation reacts withbuilding materials to accelerate weathering.Natural precipitation (rainfall not affectedby pollution) is acidic with a pH of about 6.The acidic pH is due to the presence of CO2

in the atmosphere and the production of car-bonic acid when carbon dioxide dissolves inwater. Acid precipitation or acid rain is gen-erally classified as rain in which the pH islower than its natural level due to humaneffects. In extreme cases, acid rain may havea pH as low as 2 or 3. The primary cause ofacid rain is the burning of fossil fuels. Thecombustion of sulfur-bearing coals in powerplants results in the production of sulfurdioxide, SO2. This sulfur dioxide is thenconverted into sulfurous acid (H2SO3) andsulfuric acid (H2SO4) in the atmosphere.Natural gas combustion and automobilesproduce nitrogen oxides. These can be con-verted into nitrous and nitric acids in theatmosphere. Reactions summarizing theformation of acids in rain are displayed inFigure 13.2. Acid rain causes millions ofdollars of damage each year to buildingsand structures made of marble and lime-stone. In areas impacted by acid precipita-tion, such as the northeast United States andnorthern Europe. Limestone and marblestructures suffer significant damage due to

Acids and Bases 165

its effects. Limestone and marble are cal-cium carbonate; a typical reaction repre-senting the destruction of limestone by acidrain would be

CaCO3(s) � H2SO4(aq) Ca2�(aq)

� SO42�

(aq) � H2O(l) � CO2(g)

In this reaction, calcium carbonate reactswith sulfuric acid resulting in the dissolu-tion of calcium carbonate.

In addition to the damage acid raincauses to structures, acid rain also affectsnatural environments. Significant loss ofspruce forests due to the burning of spruceneedles by acid rain has occurred in Scan-dinavia. Acid rain also extracts a heavy tollon aquatic systems and associated organ-isms. Most adult fish cannot tolerate pHsmuch lower than 5.0, and even the most tol-erant species will not survive below a pH of4.0. Fish larvae are even more susceptible tolow pH levels. Insects and their larvae alsoperish when pH approaches 4.0 in aquaticsystems. Numerous lakes in upstate New

Table 13.8Several Popular Antacids

Figure 13.2Acid rain occurs when reactions in theatmosphere lead to the formation of variousacids lowering the pH below the natural pHof rain, which is approximately 5.5. Theacidic pH of natural rain water is due to theformation of carbonic acid by the reaction:CO2(g) � H2O H2CO3.

166 Acids and Bases

York and New England are devoid of fishdue to the effects of acid rain. Indirecteffects of the low pH values associated withacid rain also affect organisms. As noted inTable 13.1, one of the properties of an acidis the ability to dissolve certain metals. Thishas a profound effect on soil subjected toacid rain. Acid rain can mobilize metal ionssuch as aluminum, iron, and manganese inthe basin surrounding a lake. This not onlydepletes the soil of these cations disruptingnutrient uptake in plants, but also introducestoxic metals into the aquatic system.

The detrimental effects of acid rain area major reason why legislation such as theClean Air Act places strict limitations onsulfur and nitrogen emissions. It is also areason why low sulfur coal is preferred overhigh sulfur coal. To reduce sulfur dioxideemissions, industry also uses a techniquecall scrubbing. Industrial scrubbers employa variety of physical and chemical processesto remove sulfur dioxide from emissions.Another technique used to combat acidifi-cation of lakes is to treat these systems withlime. The lime acts to neutralize the acid, butsuch techniques are usually costly and areonly a temporary remedy for combating theproblem.

Liming an acidic lake is similar to theprocess many people use to maintain a pHbalance in their soil for lawn maintenance.Plants have an optimum pH range in whichthey strive. Acidic conditions often developin soils for several reasons. Rain tends toleach away basic ions, weak organic acidsdevelop from the carbon dioxide producedby decaying organic matter, and strongacids, such as nitric acid, can form whenammonium fertilizers oxidize. To neutralizethese acids, different forms of lime such asquicklime, CaO, and slaked lime, Ca(OH)2,are used to neutralize the acid and increasethe pH of the soil. Table 13.9 shows howmuch fertilizer is wasted when applied to

lawns with different soil pH values. Muchof a fertilizer is wasted unless the soil’s pHis maintained at close to a neutral pH 7value.

BuffersA buffer is a solution that resists

changes in pH when an acid or base is addedto the solution. Buffers are important forlimiting the change in pH within specificlimits. Many commercial and natural prod-ucts contain buffers for pH control, forexample, shampoo, medicines, and blood.The most common buffers consist of a weakacid and its conjugate base. The acid part ofthe buffer neutralizes excess base added tothe solution while the conjugate base neu-tralizes acid added to the solution. Buffersmay also consist of a weak base and its con-jugate acid. Buffers are prepared by com-bining a weak acid (or base) and a saltcontaining its conjugate base.

To illustrate how a buffer works, con-sider the simple buffer consisting of aceticacid and the acetate ion: HC2H3O2/C2H3O2

�. This buffer is prepared by com-bining acetic acid and sodium acetate. Insolution the two components give the fol-lowing reactions:

Table 13.9Percentage of Fertilizer Wasted

Acids and Bases 167

HC2H3O2(aq) � H2O(l) H3O�

(aq) � C2H3O2�

(aq)

NaC2H3O2(s) Na�(aq) � C2H3O2

�(aq)

Acetic acid partially dissociates to givehydronium and acetate ions, and sodiumacetate dissociates completely into sodiumand acetate ion. Acid added to this solutionis neutralized by the acetate ion, while anybase added is neutralized by the hydroniumion. When these neutralization reactionstake place, the equilibrium in the first reac-tion shifts to replenish the reacted hydro-nium or acetate ions according to LeChâtelier’s Principle. Therefore, the pHremains constant. The ability of buffers toresist changes in pH is limited. Buffers doactually experience very slight changes inpH when small amounts of acid or base areadded to the buffer. Eventually, when a cer-tain amount of acid or base is added, a bufferis no longer able to control the pH withincertain limits. Buffers are generally pre-pared to control pH over a limited range.This range is dictated by the concentrationsof acid and salt used in preparing the buffer.The amount of acid or base that a buffer canabsorb before its pH moves out of its bufferrange is known as the buffering capacity.

One of the most important buffer systemsin the human body is that which keeps the pHof blood around 7.4. If the pH of blood fallbelow 6.8 or above 7.8, critical problems andeven death can occur. There are three primarybuffer systems at work in controlling the pHof blood: carbonate, phosphate, and proteins.The primary buffer system in the bloodinvolves carbonic acid, H2CO3 and its conju-gate base bicarbonate, HCO3

�. Carbonic acidis a weak acid that dissociates according tothe following reaction:

H2CO3 � H2O H3O� � HCO3

�

The equilibrium in this reaction lies to theright and the amount of bicarbonate is

approximately ten times the amount of car-bonic acid. The bicarbonate ion neutralizesexcess acid and the carbonic acid neutral-izes excess base in the blood:

H3O� � HCO3

� H2O � H2CO3

OH� � H2CO3 H2O � HCO3�

The excessive amount of bicarbonate in theblood means that blood has a much greatercapacity to neutralize acids. Many acidsaccumulate in the blood during strenuousactivity, for example lactic acid. Excretionof bicarbonate through the kidneys and the removal of carbon dioxide through respiration also regulate the carbonic acid/bicarbonate blood buffer.

Many aspirins are advertised as buf-fered aspirin, and this has led to the misun-derstanding that these aspirin are buffers. Inreality, as mentioned previously, bufferedaspirins actually contain an antacid toreduce the problems brought on by theeffects of aspirin, which is the acid calledacetylsalicylic acid. Because of its wide-spread use and importance to humans, let’stake an extended look at the history andchemistry of aspirin.

Aspirin is the most widely used medi-cation. Over 10,000 tons of aspirin areused in this country annually, and world-wide the annual consumption is 35,000tons. The history of acetylsalicylic acidactually goes back thousands of years.Hippocrates (460–377 B.C.) and the ancientGreeks used powdered willow bark andleaves to reduce fever (antipyretic) and asa pain reliever (analgesic). Native Ameri-can populations also used willow and oilof wintergreen for medication. The chem-icals responsible for the medicinal proper-ties in willow and oil of wintergreen areforms of salicylates. Willows (genus Salix)contain salicin and oil of wintergreen con-tains methyl salicylate.

168 Acids and Bases

While the use of willow bark and oil ofwintergreen as an accepted antipyretic andanalgesic has been around for at least 2,000years, by the nineteenth century medicineswere starting to be synthesized in chemicallaboratories. In 1837, Charles Frederic Ger-hardt (1819–1856) prepared salicylic acidfrom salinin, but abandoned work in thearea because of difficulties he encounteredin trying to synthesize this compound. TheGerman Adolph Wilhelm Hermann Kolbe(1818–1884) synthesized salicylic acid in1853. Salicylic acid was used for pain relief,fever, and to treat rheumatism, but it causedgastrointestinal problems and had anunpleasant taste so its use was limited. In1860, salicylic acid was identified as theagent in plants that resulted in pain relief.Chemists sought to improve medicines bysynthesizing different compounds contain-ing salicylic acid. Sodium salicylate wasfirst used around 1875, and phenyl salicy-late, known as salol, appeared in 1886, butboth these produced the undesirable gas-trointestinal side effects.

The discovery of aspirin or acetylsali-cylic acid as a pain reliever is credited toFelix Hoffman (1868–1946) who was look-ing for a substitute for sodium salicylate totreat his father’s arthritis. Hoffman uncov-ered and continued the work of Gerhardtfrom forty years before. Hoffman reactedsalicylic acid with acetic acid to produceacetylsalicylic acid in 1897 Figure 13.3.

Hoffman’s acetylsalicylic acid com-peted with the other salicylate compoundsafter its initial synthesis. In 1899 Hoffman’semployer, the Bayer Company, founded in1861 by Friedrich Bayer (1825–1880),began to market acetylsalicylic acid underthe name Aspirin. The term “aspirin” wasderived by combining the letter “a” fromacetyl and “spirin” from spiric acid. Spiricacid is another name for salicylic acid foundin plants of the genus Spirea. Bayer pro-

moted aspirin widely at the turn of the cen-tury and rapidly established a large marketfor its use. Bayer also acquired a trademarkto the name Aspirin, but relinquished theserights to England, the United States, andFrance as part of the Treaty of Versaillesending World War I. After Bayer’s patent onaspirin expired, other companies started toproduce aspirin.

Aspirin’s original use as an analgesic,antipyretic, and to reduce inflammation con-tinues to this day, and more recently someevidence has been found that it may lessenthe chance of heart attacks due to its effect asa blood “thinner.” Just as aspirin continues toprovide the same benefits as a century ago, italso produces some of the same problems.The major problem is that it can upset thestomach. In the acidic environment of thestomach, aspirin can diffuse through the pro-tective mucous lining of the stomach and rup-ture cells and produce bleeding. Undernormal doses, the amount of blood loss inmost individuals is only a milliliter or two,but in some individuals who take heavy dosesbleeding can be severe. To counterattack thisside effect, manufacturers include an antacidsuch as aluminum hydroxide and call theaspirin a buffered aspirin. As noted previ-

Figure 13.3Acetylsalicylic Acid

C6H4(OH)COOH � HC2H3O2 r C9H8O4 � H2O

Salicylic acid � Acetic acid r Acetylsalicylic acid � Water

Acids and Bases 169

ously, this term is misleading because theantacid does not buffer the solution, but neu-tralizes some of the acidic effects of theaspirin. Another type of aspirin called entericaspirin dissolves in the intestines where theenvironment is more basic.

SummaryWe have seen in this chapter that acids

and bases describe two broad classes ofcompounds that have significant impacts onour lives. Our bodies contain scores of these

compounds. We are exposed to hundreds ofacids and bases daily. It was only at the startof the twentieth century that an adequatedefinition for acids and bases was proposedby several individuals. The most commonacids and bases are those that occur in aque-ous solutions. A solution’s pH defineswhether a solution is an acid or base and isthe one of the most common measurementsmade in chemistry. As we refer to acids andbases in the pages ahead, we will have amuch better understanding of these impor-tant compounds.

IntroductionWhenever you start a car, use a battery-powered device, apply a rust inhibitor to apiece of metal, or use bleach to whiten yourclothes, you deal with some aspect of elec-trochemistry. Electrochemistry is thatbranch of science that involves the interac-tion of electrical energy and chemistry.Many of our daily activities use some formof electrochemistry. Just imagine how yourlife would be in a world without batteries.What immediately comes to mind is the lossof power for our portable electronic devices.While this would certainly be an inconve-nience, consider the more critical needs ofthose with battery-powered wheelchairs,hearing aids, or heart pacemakers. In thischapter, we examine the basic principles ofelectrochemistry and some of their applica-tions in our lives.

HistoryThe phenomenon called electricity has

been observed since the earliest days ofhumans. Lightning struck fear in our pre-historic ancestors and led to supernaturalexplanations for the displays accepted today

as a natural atmospheric occurrence. Theancient Greeks observed that when a treeresin they called elektron hardened intoamber, it possessed a special attractive prop-erty. When amber was rubbed with fur orhair, it attracted certain small objects. TheGreek Thales observed this attractive prop-erty and noted that this property also existedin certain rocks. Thales’ attractive rock waslodestone (magnetite), which is an iron-bearing mineral with natural magnetic prop-erties. The modern study of electricitycommenced with the publication of WilliamGilbert’s (1544–1603) De Magnete in 1600.Gilbert coined the word “electricity” fromthe Greek “elektron” and proposed that theEarth acted as a giant magnet.

During the seventeenth and eighteenthcenturies, several individuals made signifi-cant contributions to the science of electric-ity. Otto von Guericke (1602–1686) createdone of the first electrostatic generators bybuilding a sphere of sulfur that could berotated to accumulate and hold an electriccharge. Numerous devices were built to gen-erate static electricity through friction, andother instruments in the form of plates, jars,and probes were constructed to hold andtransfer charge (Figure 14.1). One of the

14

Electrochemistry

172 Electrochemistry

most popular of these devices was the Ley-den jar, which was perfected at the Univer-sity of Leyden by the Dutch physicist Pietervan Musschenbroek (1692–1761). The orig-inal Leyden jar was a globular-shaped glasscontainer filled with water fitted with aninsulated stopper. A nail or wire extendedthrough the stopper into the water. The Ley-den jar was charged by contacting the endof the protruding nail with an electrostaticgenerator. Leyden jars are still used in sci-ence laboratories.

Charles Du Fay (1698–1739) observedthat charged objects both attracted andrepelled objects and explained this by posit-ing that electricity consisted of two differ-ent kinds of fluids. Du Fay called the twofluids vitreous and resinous electricity andsaid each of these two different fluidsattracted each other, but each repelled itself.Benjamin Franklin (1706–1790) also con-

sidered electricity a fluid, but he consideredelectricity a single fluid. Franklin consid-ered the fluid either present, positivelycharged, or absent, negatively charged, andsaid that the electrical fluid flowed frompositive to negative. The practice of defin-ing electricity as flowing between positiveand negative poles can be attributed toFranklin.

The birth of electrochemistry paralleledthe birth of modern chemistry. Both oc-curred at the end of the eighteenth century,and both grew out of conflicting theoriessupported by eminent scientists. The battleover phlogiston between Priestley andLavoisier contributed largely to the devel-opment of modern chemistry. During thesame period that Lavosier and Priestleywere advancing their ideas, the two promi-nent Italian scientists Luigi Galvani (1737–1798) and Alessandro Volta (1745–1827)

Figure 14.1Numerous devices were used to generate and hold charge: a) Guericke’s generator, b) Leydenjar, and c) Wimhurst generator. (Rae Déjur)

Electrochemistry 173

held opposite views on the subject of ani-mal electricity. Galvani, physician and pro-fessor of anatomy at the University ofBologna, devoted most of his life to thestudy of the physiology of muscle andnerve response to electrical stimuli. Gal-vani’s work was a natural consequence ofhis profession and the period in which helived. Electrical shock was a popular thera-peutic treatment in the late eighteenth cen-tury. Patients were subjected to staticelectrical shocks as a means to treat all sortsof ailments including gout, rheumatism,infertility, toothaches, and mental disorders(Figure 14.2). The area for which this

“shock therapy” seemed to hold the mostpromise was in treating paralysis. The hopeof reversing paralysis grew out of experi-ments in which dissected body parts of ani-mals responded to electrical dischargethrough them. Within this framework, Gal-vani and other researchers sought a theo-retical explanation for the contraction ofmuscles when nerves were electrically stim-ulated.

Galvani conducted hundreds of experi-ments starting in the late 1770s. Theseexperiments consisted of using dissectedbody parts and subjecting them to chargeheld on a device such as a Leyden jar

Figure 14.2Galvani’s Study of Animal Electricity (Courtesy of the Bakken Library, Minneapolis)


(Figure 14.1). Researchers would use vari-ous conductors and configurations to trans-fer the electricity to the tissue and note theresponse of the limb. A favorite subject wasusing frog legs and the attached spinal cord.Over the years, Galvani developed his the-ory of animal electricity. He explained theresponse of limbs to static electricity byproposing that animals contained a specialelectrical fluid within their muscles ornerves. According to Galvani, limbs movedwhen the animal electricity flowed betweenareas of positive and negative accumulation.

Galvani published his theory on animalelectricity in 1792. Alessandro Volta, aphysics professor at the University of Pavia,was a respected colleague of Galvani whoinitially supported Galvani’s theory, butbegan to raise questions on the work soonafter its publication. In addition to ques-tioning Galvani’s animal electricity theory,Volta commenced his own series of experi-ments on the subject. Volta believed musclecontraction was not due to a flow of nervo-electrical fluid, but was a result of an elec-trical flow between dissimilar metals used inthe experiment. Volta’s experiments demon-strated that the stimulation of nerves causedthe contractions. A key part of Volta’s exper-iments involved placing dissimilar metalsalong small segments of nerves, providing aconnection between the metals, and observ-ing that the entire limb responded when theconnection was made. Volta’s theorybecame known as the contact theory for ani-mal electricity.

Galvani’s death in 1798 left his theoryto be defended by his proponents, but histheory was ultimately proven incorrect.Modern science, though, has shown thatGalvani was correct in believing that animalelectricity could be generated from withinthe organism. Today it’s known that electri-cal impulses are generated by chemical processes associated with neurons and

chemicals known as neurotransmitters.Nerves can also be stimulated by externalstimuli. In summary, although both Gal-vani’s and Volta’s ideas had flaws, they pro-vided fundamental knowledge on howmuscles and nerves work.

The most important result of the Galvani-Volta controversy was a device thatVolta created to study how dissimilar met-als in contact with each other generatedelectricity. Volta’s pile, which was actuallythe first battery created, consisted of stacksof metal disks about the size of a quarter(Figure 14.3). Two different metals wereused to make the disks such as silver andzinc, and the disks were arranged so that onemetal rested on another forming a pair.Between each pair of metals Volta inserted apaper or felt disk that had been soaked inwater or brine solution (Figure 14.3). Volta’soriginal pile contained about 20 Ag-Znpairs. Besides the pile configuration, Voltaalso used an arrangement called the crownof cups. The crown of cups consisted ofnonconducting cups filled with a water orbrine solution. The cups were connected bybimetallic conductors (Figure 14.4).

If a date were to be placed on the birthof electrochemistry, that date would beMarch 20, 1800. This is the date of a letterVolta sent to Sir Joseph Banks, president of

Figure 14.3The Arrangement of Disks in Volta’s Pile.


the Royal Society of London. Included inthe letter was a description of Volta’s pileand his findings on how an electrical currentcould be generated by connecting metals.Volta observed that he received a noticeableshock when he placed two wet fingers onopposite ends of the pile. The same was truefor his crown of cups, where he noted the

shock progressively increased as the num-ber of cups between his fingers increased.

Even before its official publication,Volta’s findings received considerable atten-tion. Here was a device that could provide asteady, although small, current of electricity.Until this time, scientists used the process ofgenerating static electricity and transferring

Figure 14.4Volta’s Crown of Cups (top) and Pile


this charge in pulses. Volta’s results werepublished in English (the original letter waswritten in French) in The PhilosophicalMagazine, September 1800. Other scientistsimmediately constructed their own versionsof Volta’s pile to study this new type of elec-tricity. One of the first experiments using apile was done by William Nicholson (1753–1815) and Anthony Carlisle (1768–1840).They created a pile and used it to electrolyzewater into hydrogen and oxygen.

The most important figure that emergedin the new science of electrochemistry wasHumphrey Davy. Volta had attributed thecurrent generated in his pile to the directcontact between metals. His work focused onthe physical relationship between the metals,and for the most part neglected the chemicalaspects of the electrical current. Davybelieved a chemical reaction was the basis ofthe current produced in the piles or electro-chemical cells. Davy discovered that no cur-rent was generated in the cells if pure waterwas used, and better results were obtainedwhen an acid solution was used to separatemetals. Davy constructed a variety of cells.He used cups of agate and gold and experi-mented with various metals and solutions toconstruct cells that were more powerful.Using a huge battery constructed by con-necting over two hundred cells, Davy appliedcurrent to various salts and his work led tothe discovery of a number of new elements.Potassium was discovered when Davy rancurrent through potash in 1807 and severaldays later he isolated sodium from soda in asimilar fashion. The next year he discoveredmagnesium, strontium, barium, and calcium.Davy determined oxygen was not a productin the electrolysis of hydrochloric acid lead-ing him to refute Lavoisier’s oxygen theoryof acids. Because of his work on HCl, Davyis sometimes credited with the discovery ofchlorine that had been isolated a generationearlier by Scheele.

The work of Davy was continued andexpanded upon by the great English scientistMichael Faraday (1791–1867). Faraday’sprimary studies in electrochemistry tookplace between 1833 and 1836. Faraday isresponsible for giving us much of our mod-ern electrochemical terminology. The terms“electrode,” “anode,” “cathode,” “elec-trolyte,” “anion,” “cation,” and “electroly-sis” are all attributed to Faraday. Even moreimportant than his qualitative description ofelectrochemistry, Faraday did quantitativestudies that led to his formulation of elec-trochemical laws. These laws provided a means to determine the relationshipbetween current and the amount of materialsreacting in an electrochemical reaction.Because Faraday’s major contributions arestill used today, they are covered in the prin-ciples of electrochemistry later in the chap-ter rather than in this historical section.

Oxidation and ReductionEarly pioneers in the field of electro-

chemistry had no knowledge of electrons orthe structure of the atom. In fact, early stud-ies involving electrochemistry were con-temporaneous with Dalton’s proposedatomic theory. It was not until Thomson’sdiscovery of the electron at the end of thenineteenth century that a true understandingof electrochemical processes could begin.Chemical reactions involve the breaking andformation of bonds. These bonds consist ofelectron pairs that rearrange as atoms sepa-rate and recombine during a chemical reac-tion. Many reactions involve the transfer ofelectrons from one substance to another. Thetransfer of electrons is responsible for ionicbonding. As a brief review, consider the sim-ple reaction of magnesium and oxygen toform magnesium oxide:

2Mg � O2 2MgO


In this reaction, each of the two magnesiumatoms donates two electrons to the two oxy-gen atoms making up the oxygen molecule.In the process, each magnesium atombecomes Mg2� and each oxygen atombecomes O2�

Mg Mg2� � 2e� O � 2e� O2�

Because the balanced equation involves twomagnesium and oxygen atoms, the previousequations are more appropriately written as

2Mg 2Mg2� � 4e� O2 � 4e� 2O2�

The two magnesium ions with a positivecharge are attracted to the two negativelycharged oxygen atoms to form the ioniccompound magnesium oxide, MgO.

The reaction of magnesium and oxygenis an example of an oxidation reaction. Thecombination of an element with oxygen wasthe traditional way to define an oxidationreaction. This definition of oxidation hasbeen broadened by chemists to include reac-tions that do not involve oxygen. Our mod-ern definition for oxidation is that oxidationtakes place when a substance loses elec-trons. Anytime oxidation takes place and asubstance loses one or more electrons,another substance must gain the electron(s).When a substance gains one or more elec-trons, the process is known as reduction.Reactions that involve the transfer of one ormore electrons always involve both oxida-tion and reduction. These reactions areknown as oxidation-reduction or redoxreactions.

Redox reactions always consist of oneoxidation reaction and one reduction reac-tion. The separate oxidation and reductionreactions are known as half reactions. Thesum of the two half reactions gives the over-all reaction. When the half reactions aresummed, there is an equal number of elec-

trons on each side of the equation. In ourexample, the sum of the two half reactions is

When writing the overall reaction, it is cus-tomary to cancel the number of electrons onboth side of the chemical equation.

In the formation of magnesium oxide,magnesium undergoes oxidation and oxy-gen undergoes reduction. Another way ofsaying this is that magnesium was oxidizedand oxygen was reduced. Because oxygenaccepted electrons causing magnesium to beoxidized, oxygen was the oxidizing agent inthe reaction. In a similar manner, magne-sium donated electrons and caused the oxy-gen to be reduced so it was the reducingagent. The terms associated with redoxreactions are summarized in Table 14.1.

Oxidation NumbersIn the oxidation of magnesium, each

magnesium atom loses two electrons andacquires a charge of �2, and each oxygenatom accepts these two electrons andacquires a charge of –2. The charge or ap-parent charge an atom has or acquires iscalled its oxidation number. The oxidationnumber of magnesium in MgO is �2, andthe oxidation number of O is –2. The con-cept of oxidation number is used in chem-istry as a form of electron accounting. In themagnesium oxide reaction, electrons aretransferred from magnesium to oxygen. Themagnesium and oxygen start in theiruncombined elemental forms so their oxi-dation numbers are originally both zero. Asnoted already after Mg and O combine, theiroxidation numbers become �2 and –2,respectively. Likewise, when sodium com-bines with chlorine to give NaCl, the oxida-tion number of Na changes from 0 to �1

2Mg r 2Mg2� � 4e� oxidation half reaction� O2 � 4e� r 2O2� reduction half reaction

2Mg � O2 � 4e� r 2Mg2� � � 4e��2O2


and chlorine changes from 0 to –1. Assign-ing oxidation numbers to atoms taking partin a redox reaction gives us another meansto identify the oxidizing and reducingagents. The reducing agent’s (substancebeing oxidized) oxidation number willincrease, and the oxidizing agent’s oxidationnumber will decrease.

The definition for oxidation numbersstates that it is the charge or apparent chargeof an atom. Apparent charge means thecharge of an atom if it is assumed that theelectrons are not shared but are assigned tothe more electronegative element in amolecule. By applying the concept of appar-ent charge, oxidation numbers can beassigned to atoms in covalently bondedmolecules even though the atoms’ electronsare shared. In essence, the oxidation numberconcept assumes that the bonds are ionicrather than covalent in substances. For exam-ple, in the water molecule each hydrogenatom shares its single electron with oxygento form H2O. To assign oxidation numbers tohydrogen and oxygen in water, it is assumedthat oxygen, because it is the more elec-tronegative element, possesses both elec-trons. It’s as though H2O acted like an ioniccompound and each hydrogen atom donatesits electron to oxygen. When this happens,oxygen acquires a –2 charge and each hydro-

gen atom obtains a charge of �1. It must bestressed that when oxidation numbers areassigned it is assumed that electrons are notshared. This false assumption is made for thesole purpose of being able to assign oxida-tion numbers to elements in compounds.

The determination of oxidation num-bers can be a bit tricky, but fortunately thereare some general guidelines to help us.These general guidelines are:

1. The oxidation number of an atom in its ele-mentary uncombined state is 0.

2. The oxidation of a monatomic ion isequal to the charge of the ion.

3. The oxidation number of hydrogen is usu-ally �1.

4. The oxidation number of oxygen is usu-ally –2.

5. The oxidation number of alkali metals isalways �1.

6. The oxidation number of alkali earth met-als is always �2.

7. The sum of the oxidation numbers of allatoms in a neutral compound is equal tozero. In a polyatomic ion, the sum of theoxidation numbers of all atoms is equal tothe charge of the ion.

It can now be seen how these rules apply tothe previous examples. For MgO, Mg has an

Table 14.1Terms Associated with Redox Reactions


oxidation number of �2 (rule 6) and O hasan oxidation of �2 (rule 3). In water, hydro-gen has an oxidation of �1 (rule 1) and oxy-gen has an oxidation of �2 (rule 3). Theserules can be used for a few more substancesto illustrate how to assign oxidation num-bers:

For ammonia, NH3, assign �1 to hydrogen (rule3). According to rule 7, the sum of the oxidationnumbers of all atoms must be equal to 0; there-fore, the oxidation number of N must be –3.Remember oxidation numbers are charges. Tomaintain neutral charge, the 3 hydrogen atomseach with a �1 charge must be balanced by a –3charge for nitrogen.

In sulfuric acid, H2SO4, H is assigned �1(rule 3), O is assigned –2 (rule 4). Because themolecule is neutral, the two hydrogens and fouroxygens means that S must have a charge of �6because (2 � �1) � (1 � �6) � (4 � –2) � 0.

For bicarbonate, HCO3�, H is �1 (rule 3), O

is –2 (rule 4), and C is �4. Notice in this exam-ple the sum is equal to –1, because the bicar-bonate ion has a charge of –1.

Oxidation numbers are used when writ-ing redox equations to account for the trans-fer of electrons. For example, consider thesimple reaction representing the rusting ofiron:

The oxidation numbers are written above theelements. It is seen that iron’s oxidationnumber changes from 0 to �3 indicatingthat each iron has lost three electrons. Oxy-gen’s oxidation goes from 0 to –2 indicatingeach oxygen atom gained two electrons.Because there are a total of 4 iron atoms inthe reaction, 12 electrons have been trans-ferred from iron to the 6 oxygen atoms, eachoxygen atom acquiring 2 electrons resultingin a neutral compound of iron oxide.

Electrochemical CellsThe concept of oxidation has been

expanded from a simple combination withoxygen to a process in which electrons aretransferred. Oxidation cannot take placewithout reduction, and oxidation numberscan be used to summarize the transfer ofelectrons in redox reactions. These basicconcepts can be applied to the principles ofelectrochemical cells, electrolysis, andapplications of electrochemistry.

Placing a strip of zinc in a beaker ofcopper sulfate solution, CuSO4(aq) resultsin several changes taking place. The zincimmediately darkens, and as time passed,the blue color of the CuSO4 solution light-ens. After a period, clumps of a reddishbrown precipitate are seen clinging to thezinc strip and falling and accumulating onthe bottom of the beaker. At this point, theblue color may have disappeared alto-gether, producing a clear solution. Thechanges observed result from a sponta-neous chemical reaction involving the oxi-dation of zinc and the reduction of copper.The basic concepts of oxidation and reduc-tion can be used to explain our observa-tions. The copper in solution exists asCu2� ions. These ions are reduced to solidcopper as the zinc metal is oxidized toZn2� ions. The solid copper appears as areddish-brown precipitate. The character-istic blue color of the copper solutionfades as the copper ions leave solution.This simple experiment demonstrates asimple redox reaction involving zinc andcopper:

Zn(s) � CuSO4(aq) Cu(s) � ZnSO4(aq)

In the above reaction, the exchange ofelectrons occurs directly between the cop-per and zinc on the surface of the metal.The transfer of electrons in redox reactions

0 0 � 3 –2

4Fe � 3O r 2Fe O2 2 3


can often be harnessed to do useful work.An electrochemical cell is an arrangementin which a redox reaction is used to gener-ate electricity. Electrochemical cells arealso known as voltaic or galvanic cells inhonor of Volta and Galvani. All batteriesare forms of electrochemical cells. Themain purpose of an electrochemical cell isto convert chemical energy from a sponta-neous chemical reaction into electricalenergy.

Figure 14.5 shows the basic arrange-ment of a electrochemical cell called theDaniell cell. This cell is named for JohnFrederick Daniell (1790–1845) who con-structed this type of cell in 1836. TheDaniell cell components include zinc andcopper solutions in separate containers.Between the solutions is a salt bridge

containing a solution of a strong elec-trolyte such as KCl. In each respectivesolution, there is a metal strip corre-sponding to the cation of the solution. Themetal strips are referred to as electrodes.The two metal strips are connected by awire. A voltmeter can be placed in the cir-cuit to measure the electrical potential ofthe cell.

In each compartment of the electro-chemical cell a half reaction occurs. Thetwo half reactions result in an overall reac-tion that generates a flow of electrons orcurrent. In one cell compartment, zinc isoxidized according to the reaction: Zn(s)

Zn2� � 2e�. The reduction of copper takesplace in the other cell’s compartment:Cu2�

(aq) � 2e� Cu(s). Notice that thesereactions are the same ones that take place

Figure 14.5Daniell Cell (Rae Déjur)


when a zinc strip is placed in copper sul-fate solution. In the cell arrangement, theelectrons lost by zinc when it oxidizes flowthrough the wire and react with the copperions on the surface of the copper electrode.If the masses of the copper and zinc elec-trodes were measured over time, the massof the zinc strip would decrease becausezinc is being oxidized. The mass of thecopper electrode would increase as solidcopper plates on to it. The salt bridge com-pletes the circuit and maintains a balanceof charge in each compartment. As zinc isoxidized creating positive zinc ions in onecompartment, the negative ion (Cl�) in thesalt bridge flows into this compartment. Ascopper is reduced in the other compart-ment, the positive ion (K�) diffuses intothe solution to replace this loss of positivecharge. Therefore, as the reaction contin-ues both solutions remain neutral as ionsfrom the salt bridge diffuse into both com-partments.

The Daniell cell illustrates the basic fea-tures of an electrochemical cell. Electro-chemical cells always involve a redoxreaction. Oxidation occurs at the cathode ofthe cell and reduction takes place at theanode. Electrons always flow from theanode to the cathode. Electrochemical cellscome in many arrangements. To gain anappreciation for the variety of electrochem-ical cells, consider all the types of batteriesavailable.

In the previous example, zinc was moreeasily oxidized than copper. The ability ofone element to donate electrons to anotherelement is based on the electron structure ofthe element and energy considerations. Aspontaneous redox reaction, like any spon-taneous reaction, results in a more stableconfiguration of the chemical system. Thefollowing list of elements shows how easilyan element is oxidized.

A term used to describe how easily a metalis oxidized is active. A more active metal isone that is more easily oxidized. A listing ofmetals in order of activity is known as anactivity series. The activity series is used todetermine which substances will be oxi-dized and reduced in an electrochemicalcell: the element higher on the list will beoxidized. For example, in a cell with alu-minum and silver electrodes in their appro-priate solutions, aluminum is oxidized andsilver is reduced. Therefore, aluminum is theanode and silver is the cathode. If you haveever bitten a piece of aluminum foil andexperienced discomfort, you had this elec-trochemical process occur in your mouth.Silver (or mercury) fillings and the alu-minum serve as electrodes and your salivaserves as an electrolyte between the two.The resulting current stimulates the nervesin your mouth resulting in the discomfort.

The Daniell cell generates an electricalcurrent, and hence a voltage is created whenthe cell operates. Electrical current can bethought of as the flow of electrons or charge

lithium (Li) � more easily oxidizedpotassium (K)barium (Ba)calcium (Ca)sodium (Na)magnesium (Mg)aluminum (Al)zinc (Zn)chromium (Cr)iron (Fe)cadmium (Cd)nickel (Ni)tin (Sn)copper (Cu)mercury (Hg)silver (Ag)plantium (Pt)gold (Au) less easily oxidized

�


in a circuit, but the concept of voltage is alittle less obvious. Voltage measures theelectrical potential or electromotive forceof a charge. What does this actually mean?Potential energy is the ability of a system todo work by virtue of its position or configu-ration. When referring to gravitationalpotential energy, sea level is used as a refer-ence, hills and mountains have positivegravitational potential energy. The higherthe elevation of a unit mass (1 kg), the moregravitational potential energy it has. In asimilar fashion, electrical potential can beconsidered as the potential energy of a unitcharge in electrical space. This unit chargecan be at different levels. When the charge ison an electrical hill or mountain, its electri-cal potential is higher than when in a valley.Just as the ability of a unit mass to do workis governed by its difference in elevation, thedifference in electrical potential dictates theability of a unit charge to do work. Electri-cal potential is measured using the unitsvolts.

A voltmeter connected to the circuit ofthe cell shown in Figure 14.5 would read1.10 V. This measurement assumes that theconcentrations of the solutions are both 1 Mand the temperature is 25°C. Under theseconditions (and for gases at 1 atmosphere),the measured voltage is referred to as thestandard potential of the cell, symbolizedE0. Different electrochemical cells obvi-ously give different E0 measurements.

To determine the E0 of different cellarrangements, chemists use what are calledstandard reduction potentials for half-cells. A standard reduction potential is theelectrical potential under standard condi-tions of a cell compared to the standardhydrogen electrode. The standard hydrogenelectrode is a special half-cell that has beenchosen as a reference to measure electricalpotential. Just as sea level is a logical ele-vation for measuring gravitational potential,

an electrical “sea level” must be chosen formeasuring cell potentials. The standardhydrogen electrode is shown in Figure 14.6.Hydrogen gas at 1 atmosphere pressure and25°C is bubbled through a 1 M solution ofhydrochloric acid. At the platinum elec-trode, the following reduction reaction isdefined to have a standard reduction poten-tial of 0:

2H�(aq) � 2e� H2(g) E0 � 0 V

Remember that the standard hydrogenelectrode is arbitrarily defined as having avoltage of 0; it cannot be measured. Thegravitational potential energy does not haveto be defined as 0 at sea level, but using thismeasurement as a reference allows the grav-itational potential energy at any other ele-vation to be determined. Similarly, thereduction potential of half-cell can be mea-sured with respect to the standard hydrogenelectrode.

Chemists have chosen to use standardreduction potentials as the basis of calculat-

Figure 14.6Standard Hydrogen Electrode (SHE)


ing E0. The oxidation potential of a half-cellis obtained simply by changing the sign ofthe reduction half-cell. This might sound abit confusing so let’s demonstrate how theprocedure would work to obtain the 1.10 Vof the Daniell cell. Each of the two half-cellsare connected to the standard hydrogen elec-trode as shown in Figure 14.7.

The two reduction reactions and themeasured voltages are

Cu2�(aq) � 2e� Cu(s) E0 � 0.34 V

Zn2�(aq) � 2e� Zn(s) E0 � –0.76 V

There is nothing wrong with having a neg-ative potential. Again, referring to the com-parison to gravity, there are places on earththat are below sea level and have a negativegravitational potential. In the Daniell cell,zinc is oxidized. To obtain the oxidationpotential of zinc, the reduction reactionmust be reversed and the sign of E0

changed:

Zn(s) Zn2� � 2e� E0 � �0.76

Adding the two half-cell reactions to obtainthe overall reaction and the cell E0 gives

Chemistry handbooks and many texts listthe standard reduction potentials for numer-ous half reactions. A small sampling ofthese is provided in Table 14.2.

The reduction potential increases mov-ing up the table. This means that substancesnear the top of the table are more likely tobe reduced (are better oxidizing agents) andsubstances near the bottom are more likelyto be oxidized (are better reducing agents).The substances in Table 14.2 correspond tothe previous listing of how easily substancesare oxidized. When two substances in Table14.2 take part in a redox reaction, the onehigher in the table is the substance that isreduced.

Thinking about the position of elementsin the periodic table and their valence elec-tron structure should help in understandingthe relative order of the reduction potentials.Fluorine gas is very electronegative andreadily accepts electrons to obtain a stableconfiguration. Conversely, alkalines and

Figure 14.7Two half-cells in the Daniell cell are each connected to the standard hydrogen half-cell (RaeDéjur)

Zn(s) r Zn2� � 2e E0 � �0.76V� Cu2� � 2e� r Cu(s) E0 � 0.34 V

Zn(s) � r Cu(s) � E0 � � 1.10V2� 2�Cu Zn(aq) (aq)


alkali earth metals such as sodium and mag-nesium readily donate electrons to obtain astable octet electron configuration; therefore,their reduction potentials are low or equiva-lently their oxidation potentials are high.

One more example demonstrates how touse standard reduction potentials to deter-mine the standard potential of a cell. Let’ssay you wanted to construct a cell using sil-ver and zinc. This cell resembles the Daniellcell of the previous example except that asilver electrode is substituted for the copperelectrode and a silver nitrate solution is usedin place of copper sulfate. From Table 14.2,it is determined that when silver and copperinteract silver is reduced and copper oxi-dized. The two relevant reactions are

Ag�(aq) � e� Ag(s) E0 � 0.80V

Zn2� � 2e� Zn(s) E0 � �0.76 V

Because zinc is oxidized, the zinc reactionmust be reversed and the sign of its E0

changed to give

Ag�(aq) � e� Ag(s) E0 � 0.80V

Zn(s) Zn2� � 2e� E0 � �0.76 V

The number of electrons lost by zinc andgained by silver is not equal. To balance thetransfer of charge, the silver half reactionmust be multiplied by 2:

2Ag�(aq) � 2e� 2Ag(s) E0 � 0.80V

Zn(s) Zn2� � 2e� E0 � �0.76 V

Notice that when the silver reaction is mul-tiplied by 2 its E0 value stays the same. Thepotential of a half-cell does not depend onthe coefficients of the equation because thepotential of the cell is independent of thequantities of reactants. Adding the two halfreactions gives the overall reaction and thecell voltage:

Zn(s) � 2Ag� Zn2� � 2Ag(s) E0 � 1.56 V

ElectrolysisElectrochemical cells produce electrical

energy from a spontaneous chemical reac-tion. In electrolysis, the process is reversedso that electrical energy is used to carry outa nonspontaneous chemical change. A cellarranged to do this is called an electrolyticcell. An electrolytic cell is similar to an elec-trochemical cell except that an electrolyticcell’s circuit includes a power source, forexample, a battery. The same electrochemi-cal cell terminology applies to electrolyticcells. Reduction occurs at the cathode andoxidation at the anode.

In an electrolytic process, redox reac-tions that occur spontaneously in electro-chemical cells can be reversed. One of themost common electrolytic proceduresdemonstrating this is when a battery is

Table 14.2Standard Reduction Potentials


charged. During its normal operation, arechargeable battery functions as an electro-chemical cell. When charging it, electricalenergy reverses the process. This conceptcan be applied to the previous example usingsilver and zinc. If instead of reducing silverand oxidizing zinc, we want to oxidize silverand reduce zinc, we attach a power supply tothe circuit and use a voltage greater than1.56 V to reverse the current. Electrons nowflow from the silver electrode to the zincelectrode where Zn2� is reduced. Oxidationoccurs at the solid silver electrode.

On several occasions the electrolysis ofwater has been mentioned. This was one ofthe first investigations conducted withVolta’s pile. The electrolysis of water can beobserved by placing a small 9 V battery in aglass of water and sprinkling in a little salt(to create an electrolyte to conduct current).Very soon a steady stream of bubbles willappear emerging from the positive and neg-ative terminals. The standard half reactionsrepresenting the electrolysis of water are

The reactions show that oxygen gas is pro-duced and the solution becomes acidic at theanode. Hydrogen gas and a basic environ-ment occur at the cathode. The overall reac-tion for the electrolysis of water can besimplified if it’s assumed the hydrogen andhydroxide ions combine to form water.Doing this and canceling water from bothsides of the equation gives

2H2O(l) 2H2(g) � O2(g)

Note that the actual potential to electrolyzepure water is about 1.25 volts. Remember,

the value for E0, which is –2.06 V, assumesstandard conditions where the concentrationof hydrogen and hydroxide ions are both 1M. In reality, the concentration of H� andOH� is only 10�7 M, and this changes theactual cell voltage.

BatteriesElectrochemical processes occur all

around us. We close this chapter by examin-ing a few of these processes and relatingthem to the electrochemical principles pre-viously introduced. Batteries are probablythe most common example of electrochem-ical applications associated with every-day life. While batteries come in all sizesand shapes, all batteries contain the basicelements common to all electrochemicalcells. What differentiates one battery fromanother are the materials used for cathode,anode, and electrolyte, and how these mate-rials are arranged to make a battery. Thestandard dry cell battery or alkaline cell isshown in Figure 14.8. Batteries consist of

anode:

2H2O(l)rO2(g)� �4e� E0��1.23V�4H(aq)

cathode:

�4H2O(l)�4e�r2H2(g)� E0��0.83V�4OH(aq)

6H2O(1)r2H2(g)�O2(g)� � E0��2.06� �4H 40H(aq) (aq)

Figure 14.8Dry Cell Battery (Rae Déjur)


two or more cells connected to deliver a spe-cific charge, although the terms are oftenused synonymously.

Dry cells are classified as primary cellsbecause the chemicals in them cannot beregenerated, that is, primary cells cannot berecharged. The dry cell battery was patentedby the French engineer George Leclanché(1839–1882) in 1866. The standard dry cellconsists of a zinc container that acts as theanode of the cell. The container contains anelectrolyte (this corresponds to the saltbridge) containing a mixture of ammoniumchloride (NH4Cl), zinc chloride (ZnCl2), andstarch. A carbon (graphite) rod runs downthe axis of the battery. Surrounding the rodis a mixture of manganese dioxide (MnO2)and carbon black. (Carbon black is basicallycarbon in the form of soot produced fromburning kerosene or another hydrocarbonwith insufficient oxygen.) The MnO2-carbon black mixture acts as the cathode.The MnO2-carbon black and electrolyteexist as a moist paste, so a dry cell is notactually dry. While several reactions occurin a dry cell, the primary reactions govern-ing its operation can be represented as:

Anode: Zn(s) Zn2�(aq) � 2e�

Cathode: 2MnO2(s) � 2NH4�

(aq) � 2e�

Mn2O3(s) � 2NH3(aq) � H2O(l)

Overall: Zn(s) � 2NH4�

(aq) � 2MnO2(s)

Zn2�(aq) � 2NH3(aq) � H2O(l)

� Mn2O3(s)

When the circuit containing the battery isclosed, electrons from the oxidation of zincflow from the negative terminal through theobject being powered to the negative carbonrod that serves as an electron collector. Sev-eral reduction reactions occur at the collec-tor; one main reaction is represented earlier.The dry cell with the NH4Cl electrolyte cre-ates acidic conditions in the battery. Theacid attacks the zinc, which accelerates its

deterioration. To extend the life of a dry cell,a basic substance, typically KOH or NaOH,can be used for the electrolyte. Alkalinecells derive their name from the alkalineelectrolyte they contain.

Mercury cells, like dry cells, have a zincanode and a use a mercuric oxide (HgO)cathode. The electrolyte is potassiumhydroxide, KOH. These small, flat, metalliccells are widely used in watches, calculators,cameras, hearing aids, and other applica-tions where small size is a premium. Thereactions in the mercury cell are:

Anode: Zn(s) � 2OH�(aq) ZnO(s)

� H2O(l) � 2e�

Cathode: HgO(s) � H2O(l) � 2e� Hg(l)

� 2OH�(aq)

Overall: Zn(s) � HgO(s) ZnO(s) � Hg(l)

The voltage obtained from a mercury cell isabout 1.35 V.

A common of battery is the lead storagebattery. This battery, composed of severalcells, is used mainly in cars. The lead stor-age battery, the first secondary or recharge-able cell, was invented in 1859 by theFrench physicist Gaston Planté (1834–1889). A lead storage battery consists of aseries of plates called grids immersed in asolution of sulfuric acid (approximately 6M) that serves as the electrolyte (Figure14.9). Half the plates contain lead and serveas the anode of the cell, and the remainingplates are filled with lead dioxide (PbO2)and serve as the cathode.

The following reactions represent thetwo half-cell reactions:

Anode: Pb(s) � HSO4�

(aq) PbSO4(s)

� H� � 2e�

Cathode: PbO2(s) � HSO4�

(aq) � 3H�(aq)

� 2e� PbSO4(s) � 2H2O(l)

Overall: Pb(s) � PbO2(s) � 2H�(aq) �

2HSO4�

(aq) 2PbSO4(s) � 2H2O(l)


The cell reaction generates a potential ofabout 2 V. A car battery is created by con-necting six cells in series. The reactionsshow that electrons are created when leadoxidizes at the anode. The lead ions formedin oxidation combine with the sulfate ionsin the electrolyte to form lead sulfate. Theelectrons flow through the car’s electricalsystem back to the battery’s cathode plates.Here the electrons combine with PbO2 andhydrogen ions present in the acid and alsoform lead sulfate. While driving, the car’salternator continually recharges a battery.The alternator carries out an electrolyticprocess and reverses the reactions justshown. In this manner, the lead sulfate gen-erated on the plates is converted back intolead and lead dioxide.

Advances in the past few decades haveimproved car battery technology immensely.Many lead batteries currently manufacturedare labeled as maintenance free. This refersto the fact that the acid level in them doesnot have to be checked. The addition ofwater to a lead battery is necessary becausethe charging process causes water toundergo electrolysis. This process creates

hydrogen and oxygen gas that escapes fromthe battery vents. Modern battery technol-ogy has greatly reduced and even eliminatedthis problem by perfecting new types ofelectrodes, for example, lead-calcium andlead-selenium electrodes. Maintenance-freebatteries are sealed and water does not haveto be added to them. Prior to this develop-ment, it was generally recommended thatthe fluid level in batteries be checked on aregular basis. Additionally, a hydrometercould be used to check the density of thefluid in the battery to indicate its condition.Many manufacturers now claim a 10–15year lifetime of batteries as opposed to thetypical 3–5 year lifetime of twenty yearsago. While improved technology has pro-duced batteries that are more reliable, youstill may experience the problem of startinga car on a cold day. The inability to start avehicle on a cold day results from chargebeing impeded as the electrolyte turns vis-cous in the cold conditions.

Another familiar secondary cell is ni-cad or nickel cadmium cell. In these cells,the anode is cadmium and the cathode con-sists of a nickel compound (NiO(OH)(s)).One of the problems with ni-cad cells is thehigh toxicity of cadmium. An alternative toni-cad cells currently gaining a foothold inthe market is lithium ion cells. Lithium ismuch less toxic than cadmium. Additionally,lithium ion cells do not suffer from thememory effect of ni-cad cells. This effectresults in the loss of efficiency when ni-cadcells are recharged before they completelylose their charge. Besides these factors,lithium has several properties that make it ahighly desirable material for use in batter-ies. Table 14.2 shows that lithium has thelowest reduction potential. This is equivalentto saying it has the highest oxidation poten-tial. Lithium’s high oxidation potentialmeans that a greater voltage can be obtainedusing lithium cells. Another property of

Figure 14.9Lead Storage Battery (Rae Déjur)


lithium is that it is light, which is especiallyimportant for portable applications such ascell phones and laptop computers. A finaladvantage of lithium ions is that they aresmall and move more readily through mate-rials than large ions.

Original lithium cells employed lithiummetal as the anode, and these cells havebeen used in the military and space programfor years. Lithium, though, is highly reac-tive presenting fire and explosive hazards.Lithium ion cells retain the advantages oforiginal lithium batteries, but they eliminatethe hazards associated with them. In lithiumion cells, the lithium metal anode is replacedby a substance with the ability to absorblithium ions such as graphite. The cathodeis made of a lithium compound such asLiCoO2, LiSO2, or LiMnO2. The electrolyteseparating the electrodes allows the passageof lithium ions but prevents the flow of elec-trons. The electrons move from the anode tocathode through the external circuit.

Numerous other types of cells existsuch as zinc-air, aluminum-air, sodium sul-fur, and nickel-metal hydride (NiMH).Companies are on a continual quest todevelop cells for better batteries for a widerange of applications. Each battery must beevaluated with respect to its intended useand such factors as size, cost, safety, shelf-life, charging characteristics, and voltage.As the twenty-first century unfolds, cellsseem to be playing an ever-increasing rolein society. Much of this is due to advances inthe consumer electronics and the computerindustry, but there have also been demandsin numerous other areas. These include battery-powered tools, remote data collec-tion, transportation (electric vehicles), andmedicine.

One area of research that will advancein years to come is the study of fuel cells.The first fuel cell was actually produced in1839 by Sir William Grove (1811–1896). A

fuel cell is essentially a battery in which thechemicals are continuously supplied froman external source. One simple type of fuelcell is the hydrogen-oxygen fuel cell (Fig-ure 14.10). These cells were used in theGemini and Apollo space programs. In thehydrogen-oxygen fuel cells, the reactions atthe anode and cathode are

Anode: 2H2(g) � 4OH�(aq) 4H2O(l) � 4e�

Cathode: O2(g) � 2H2O(l) � 4e� 4OH�(aq)

Overall: 2H2(g) � O2(g) H2O(l)

The electrodes in the hydrogen-oxygen cellare porous carbon rods that contain a plat-inum catalyst. The electrolyte is a hot (sev-eral hundred degrees) potassium hydroxidesolution. Hydrogen is oxidized at the anodewhere the hydrogen and hydroxide ionscombine to form water. Electrons flowthrough the external circuit.

Many different types of fuel cells arecurrently under development. Many of theseare named after the electrolyte or fuel usedin the cell. The polymer electrolyte mem-brane or proton exchange membrane cell(pem) also uses hydrogen and oxygen. The

Figure 14.10Schematic of Hydrogen-Oxygen Fuel Cell


electrolyte in this cell allows hydrogen ions(protons) to pass to the cathode where theycombine with oxygen to produce water. Thisfuel cell has several advantages over thealkaline fuel cell. Its lower operating tem-perature and relatively lightweight make itsuitable for use as a power supply for vehi-cles. Several prototype vehicles have beenbuilt using pem cells.

CorrosionCorrosion refers to the gradual natural

deterioration of a substance due to a chem-ical or electrochemical reaction. Many met-als and nonmetals, such as glass, are subjectto corrosion, but the most common form ofcorrosion is rusting of iron. In rusting, irondoes not combine directly with oxygen butinvolves the oxidation of iron in an electro-chemical process. There are two require-ments for rust: oxygen and water. Thenecessity of both oxygen and water is illus-trated by observing automobiles operated indry climates and ships or other iron objectsrecovered from anoxic water. Autos andships subjected to these conditions showremarkably little rust, the former due to lackof water and the latter due to lack of oxygen.

Figure 14.11 illustrates the electro-chemical rust formation process.

Rust is a product of electrochemicalreactions occurring on a piece of iron orsteel. Iron metals are never uniform. Thereare always minor irregularities in both thecomposition and physical structure of themetal. These minute differences give rise tothe anodic and cathodic regions associatedwith rust formation. During the rusting pro-cess, iron is oxidized at the anode accordingto the reaction:

Fe(s) Fe2� � 2e�.

The electrons from this reaction flowthrough the metal to an area of the metalwhere oxygen is reduced according to thereaction:

O2(g) � 4H�(aq) � 4e� 2H2O(l)

The hydrogen ions in this reaction are pro-vided mainly by carbonic acid from the dis-solution of atmospheric carbon dioxide inthe water. The water acts as an electrolyteconnecting the anode and cathode regionsof the metal. Iron ions, Fe2�, migratethrough the electrolyte and in the processare further oxidized by dissolved oxygen inthe water to Fe3�. Iron (III) oxide (Fe2O3),the technical term for rust, is formed at thecathode according to the equation:

4Fe2�(aq) � O2(g) � (4 � 2x)H2O(l)

4Fe2O3 • � xH2O(s) � 8H�(aq)

The “x” in the above equation indicates apositive whole number corresponding to thevariable amount of water molecules thattakes part in the reaction.

The area of rust formation is differentfrom the point where oxidation of iron takesplace. As noted, water serves as an elec-trolyte through which iron ions migrate.This explains why vehicles rust much morerapidly in regions where road salts are usedto melt winter ice. The salts improve the

Figure 14.11The formation of rust involves anelectrochemical reaction taking placebetween two different regions in the metal.


conductivity of the electrolyte, therebyaccelerating the corrosive process.

Most metals undergo corrosion of someform except the so-called noble metals ofgold, platinum, and palladium. Why thendon’t we experience rusting problems withmany other metals, for example, aluminum?The key is the layer of the oxide or otherproduct formed in the corrosion process.Iron (III) oxide is highly porous and as aresult rust does not protect the underlyingmetal from further corrosion. Metals suchas aluminum and zinc have an even greatertendency to oxidize compared to iron, butthe oxide layer on these metals forms animpervious protective layer that protects themetal below from further oxidation. Copperoxidizes in moist air and its surface changesfrom the characteristic reddish-brown colorto green patina. This patina consists of sul-fate compounds, CuSO4•xCu(OH)2 (wherex is a positive whole number).

A number of methods are used toreduce and prevent corrosion. The mostcommon method is to paint iron materialsso that the metals are protected from waterand oxygen. Alloying iron with other metalsis also a common means to reduce corro-sion. Stainless steel is an alloy of iron,chromium, nickel, and several other metals.Iron may also be protected by coating it withanother metal. Galvanizing refers to apply-ing a coating of zinc to protect the underly-ing metal. Additionally, because it is a moreactive metal, zinc oxidizes rather than iron.

Connecting iron objects to a more activemetal is called cathodic protection.Cathodic protection is widely used to pro-tect underground storage tanks, ship hulls,bridges, and buried pipes. One of the mostcommon forms of cathodic protection is toconnect the object to magnesium. Whenmagnesium is connected to an iron object,magnesium rather than iron becomes theanode in the oxidation process. In cathodic

protection, bars of magnesium are con-nected either directly or by wire to the ironstructure (Figure 14.12). Because the metalconnected to the iron corrodes over time, itis called the sacrificial anode. Sacrificialanodes must eventually be replaced.

ElectroplatingOur discussion of batteries and corro-

sion involves electrochemical cells. One ofthe most common applications of electroly-sis is in the area of electroplating. Duringelectroplating, the object to be electroplatedserves as the cathode in an electrolytic cell.The anode of the cell is often made of theplating metal. The object is immersed in abath or solution containing ions of the plat-ing metal and is connected to the negativeterminal of a power source. Electrons fromthe power source flow through the objectwhere they cause the plating metal ions tobe reduced on to the surface of the object(Figure 14.13).

Electroplating is another technique usedto protect iron from corrosion. The mostcommon plating metal used for this purposeis chromium.

Figure 14.12In cathodic protection, a metal more activethan iron such as magnesium or zinc is con-nected to the iron object. The more activemetal is oxidized protecting the iron.


Commercial Applicationsand Electrorefining

Most metals found in nature exist asores in combination with other elementssuch as oxygen and sulfur. Electrolysis canbe used to separate metals from their oresand remove impurities from the metals. Thegeneral process is known as electrorefining.Other electrolytic processes are used toobtain a number of important chemicals.

Chlorine (Cl2) and sodium hydroxide(NaOH) are two important chemicals pro-duced by electrolysis. Chlorine and sodiumhydroxide are generally among the top tenchemicals produced annually in the UnitedStates. The electrolysis of brine or aqueousNaCl solution is used to produce chlorine,sodium hydroxide, and hydrogen (Figure14.14). The chloride ion in the brine solu-tion is oxidized at the anode, while water isconverted at the cathode according to thefollowing reactions:

Anode: 2Cl� Cl2(g) � 2e�

Cathode: 2H2O(l) � 2e� H2(g) � 2OH�(aq)

The anode and cathode sides of the cell areseparated by a membrane that allows thesodium ions to migrate from the anode sideof the cell to the cathode side.

Electrorefining can be used to purify ametal by using alternate electrodes of a pureand impure metal. Impurities oxidized at theanode, which is made of the impure metal,travel into solution. By arranging the cellappropriately, the ion of the metal to be puri-fied is reduced on the pure metal cathode.For example, copper metal that contains leadand iron may be used as one electrode andpure copper as the other electrode in a cell.When the proper voltage is applied, copper,lead, and zinc will be oxidized and move intothe electrolyte. Because copper is more eas-ily reduced compared to zinc and lead, it willbe plated out at the pure copper cathode.Therefore, this process effectively removesthe zinc and lead impurities from the copper.

An important application of electrolysisis in the production of aluminum. Alu-minum is the most abundant metal in theEarth’s crust. Until the late 1800s, it wasconsidered a precious metal due to the dif-ficulty of obtaining pure aluminum. Theprocess involved the reduction of aluminumchloride using sodium. In 1886, CharlesMartin Hall (1863–1914) in the UnitedStates and Paul L. T. Héroult (1863–1914)in France independently discovered a meansto produce aluminum electrochemically.Hall actually started his work on aluminumin high school, continued his researchthrough his undergraduate years at OberlinCollege, and discovered a procedure for itsproduction shortly after graduating. Theproblem that confronted Hall, and otherchemists trying to produce aluminum usingelectrochemical means, was that aqueoussolutions of aluminum could not be used.This was because water rather than the alu-minum in solution was reduced at the cath-ode. To solve this problem, Hall embarked

Figure 14.13Electrolysis is used to electroplate an objectby making the object the cathode in an elec-trolytic cell.


on a search for a substance that could beused to dissolve alumina (Al2O3) to use inhis electrolytic cells. Hall and Héroultsimultaneously found that the substancecryolite could be used. Cryolite is an ioniccompound with the formula Na3AlF6.

The production of aluminum actuallyinvolves several steps. Bauxite is the ore thatcontains aluminum oxide (Al2O3) used toproduce aluminum. Impurities of iron, sul-fur, silicon and other elements are removedfrom bauxite using the Bayer process to pro-duce purified alumina. The Bayer process,patented in 1887 by Austrian Karl JosefBayer (1847–1904), involves pulverizingbauxite and treating it with a hot sodiumhydroxide solution to produce sodium alu-minate (NaAlO2). Sodium aluminate is thenplaced in a reactor in which temperature andpressure can be varied to precipitate outimpurities. The sodium aluminate solutionis then hydrolyzed to produce purified alu-mina:

2NaAlO2 � 4H2O Al2O3 • 3H2O � 2NaOH

The Hall-Héroult process involves takingthe purified alumina and dissolving it in cry-olite. This step requires producing a moltenmixture. One of the problems Hall had toovercome in developing his process wasproducing the molten alumina mixture. Purealumina’s melting point is 2,050°C, but Hallwas able to produce a cryolite-alumina mix-ture that melted at approximately 1,000°C.The purified alumina-cryolite melt is elec-trolyzed in a large iron electrolytic cell thatis lined with carbon. The iron containerserves as the cathode in the cell. The cellcontains a number of carbon electrodes thatserve as anodes (Figure 14.15). The exactnature of the reactions at the anode and cath-ode is not clearly known, but can be consid-ered to involve the reduction of aluminumand oxidation of oxygen:

anode: 6O2� 3O2(g) � 12e�

cathode: 4Al3� � 12e� 4Al(s)

The oxygen produced reacts with the carbonelectrodes to produce carbon dioxide; an

Figure 14.14The electrolysis of aqueous sodium chloride is used to prepare chlorine, sodium, and hydrogen.


Figure 14.15In the Hall-Héroult process, a moltenalumina is dissolved in cryolite and the mix-ture is electrolyzed to produce moltenaluminum.

overall equation involving conversion ofalumina into aluminum can be written:2Al2O3(l) � 3C(s) 4Al(l) � 3CO2(g).

Soon after Hall and Héroult developedtheir process, aluminum plants were con-structed to process bauxite into aluminum.Hall became a partner in the PittsburghReduction Company, which eventuallybecame ALCOA. Because the electrolyticproduction of aluminum was, and still is,energy intensive, aluminum plants werelocated in areas where electricity was abun-dant and relatively cheap. These plants weresituated near hydroelectric plants to takeadvantage of their electrical output. Thehigh electrical demand to produce alu-minum from bauxite (the aluminum indus-try is the largest consumer of electricalenergy in the United States) has created alarge market for recycled aluminum. Alu-minum made from raw bauxite is referred toas primary production aluminum, whilealuminum from recycled aluminum is sec-ondary production. The amount of energy

for secondary production of aluminum isonly about 5% of the primary productionrequirement. One way of thinking about theenergy saving involved with recycling ofaluminum is to consider that every time analuminum can is thrown away that it isequivalent to dumping half that can filledwith gasoline. The United States continuesto produce the most aluminum in the world,but over the last few decades its position hasfallen from 40% of the world market toaround 15%. Current world production isapproximately 20 million tons.

SummaryThe history of electrochemistry has

paralleled the history of modern chemistry.The discoveries of Galvani and Voltaoccurred as Lavoisier was proposing hisnew chemical system. As the theory of elec-trochemistry developed, so too did its appli-cations. Society’s dependence on electricityis brought to our attention whenever thereis a power outage. During these times, weturn to electrochemical energy to supplyour electrical needs. There is no reason tothink that as the twenty-first centuryunfolds that our reliance on cells and bat-teries and other forms of electrical storagedevices will lessen. In fact, there is everyreason to think just the opposite. The devel-opment of vehicles that pollute less, sourcesto power a mobile society’s use of commu-nication devices and computers, and a hostof cordless gadgets provide ample motiva-tion to continue the search for more effi-cient cells. These new technologies willaugment traditional chemical industrialprocesses that are based on electrochemicalprinciples.

Introduction and HistoryOrganic chemistry is the study of carboncompounds. While the formal developmentof this branch of chemistry began in themid-eighteenth century, organic chemicalshave been used throughout human history.All living material, the food we eat, and thefossil fuels we burn all contain organic com-pounds. Numerous organic substances wereprepared by ancient civilizations includingalcohols, dyes, soaps, waxes, and organicacids. From its modest beginning in theearly 1800s, an increasing number oforganic compounds have been isolated, syn-thesized, and identified. By far, organiccompounds dominate the number of knownchemicals. Currently, there are over twentymillion known organic compounds andthese make up over 80% of all known chem-icals.

The term “organic” was not used until1807 when Berzelius suggested it to distin-guish between compounds devised from liv-ing (organic) and nonliving (inorganic)matter. Before Berzelius made this distinc-tion, interest in organic compounds wasbased primarily on the medicinal use ofplants and animals. Plant chemistry involved

the study of chemicals found in fruits,leaves, saps, roots, and barks. Animal chem-istry focused on compounds found in theblood, saliva, horn, skin, hair, and urine. Incontrast to plant and animal chemistry, sci-entists distinguished those compounds iso-lated from nonliving mineral material. Forexample, mineral acids, referred to in Chap-ter 13, were made from vitriolic sulfur com-pounds.

Until the mid-eighteenth century, scien-tists believed organic compounds came onlyfrom live plants and animals. They reasonedthat organisms possessed a vital force thatenabled them to produce organic com-pounds. The first serious blow to this theoryof vitalism, which marked the beginning ofmodern organic chemistry, occurred whenFriedrich Wöhler (1800–1882) synthesizedurea from the two inorganic substances leadcyanate and ammonium hydroxide:

Pb(OCN)2 � 2NH4OH 2(NH2)2CO � Pb(OH)2

lead cyanate � ammonium hydroxide urea

� lead hydroxide

Wöhler was actually attempting to synthe-size ammonium cyanate when he discoveredcrystals of urea in his samples. He first

15

Organic Chemistry

196 Organic Chemistry

prepared urea in 1824, but he did not iden-tify this product and report his findings until1828. In a note written to Berzelius, he pro-claimed: “I must tell you that I can makeurea without the use of kidneys, either manor dog. Ammonium cyanate is urea.” WhileWöhler’s synthesis of urea signaled the birthof organic chemistry, it was not a fatal blowto vitalism. Some argued that because Wöh-ler used bone material in his preparationsthat a vital force could still have beenresponsible for Wöhler’s urea reaction. Overthe next twenty years, the vitalism theoryeroded as the foundation of organic chem-istry started to take shape. Hermann Kolbe(1818–1884) synthesized acetic acid fromseveral inorganic substances in 1844.Kolbe’s work and that of Marcellin Berth-elot (1827–1907), who produced numerousorganic compounds including methane andacetylene in 1850, marked the death of thevitalism theory.

Much of the development of modernchemistry in the first half of the eighteenthcentury involved advances made in organicchemistry. At the time of Wöhler’s synthe-sis of urea, there was much confusion aboutthe proper values of atomic masses, the roleof oxygen in acids, and how atoms com-bined. Some chemists felt organic and inor-ganic compounds followed differentchemical rules and sought a theory toreplace vitalism. One of the early explana-tions of how organic compounds formedinvolved the radical theory. Lavoisier con-sidered a radical an atom or groups of atomsthat did not contain oxygen. Radicals com-bined to give different substances. Organicsubstances were formed from the combina-tion of a carbon radical and a hydrogen rad-ical. Berzelius also believed radicalscombined to form compound radicals thatwere equivalent to organic compounds.Berzelius’ radicals were held together byelectrical forces. According to Berzelius,

organic compounds contained a negativecarbon radical and a positive hydrogen rad-ical. Compound radicals could incorporateother atoms such as oxygen and nitrogeninto their structures. Chemists employingthe radical theory sought to identify radi-cals. In 1815 Gay-Lussac identified cyano-gens gas, (CN)2, as a radical and explainedhow this radical was present in many otherorganic compounds. In 1828, Jean-BaptisteAndré Dumas (1800–1884) and PierreFrançois Guillaume Boullay (1806–1835)proposed ethylene (C2H4) as the radical thatformed the base for many different organiccompounds. For example, ethyl alcohol(C2H5OH) could be formed by adding waterto the ethylene radical. Berzelius named theethylene radical etherin and Dumas andBoullay’s idea came to be known as theetherin theory. In the same year, Wöhler andhis close friend Liebig proposed the benzoylradical (C14H10O2). Like the ethylene radi-cal, a number of different organic com-pounds could be formed by addition to thisradical.

Throughout the 1830s, chemists identi-fied other radicals. Dumas, Liebig, andBerzelius believed that organic compoundscould be classified based on their compo-nent radicals. While the radical was a meansto organize the different types of organiccompounds, it also led to much confusion.Chemists would often propose alternate rad-icals as the base for different compounds.Adding to the problem was the disagree-ment over atomic masses; therefore, the trueformulas of many compounds were stillunknown. Berzelius continued to believeradicals were held together by electrostaticforces. Some chemists believed that radicalswere true substances, while others believedthat they were merely hypothetical con-structs to explain their observations.

One serious problem with Berzelius’radical theory occurred when Dumas and

Organic Chemistry 197

his student Auguste Laurent (1808–1853)discovered that chlorine replaced hydrogenin a number of reactions. Dumas’ work inthis area originated when he was presentedwith a problem by the French royalty. Evi-dently, the emissions from candles burnedduring a royal ball irritated the guests.Dumas determined the reason for this wasthat the gas hydrogen chloride (HCl) wasemitted when the candles burned. The chlo-ride was introduced when the candles werebleached with chorine solution. Dumas pro-ceeded to conduct numerous experiments onthe chlorination of wax. As Dumas, Laurent,and others reported that chlorine could substitute for hydrogen in organic com-pounds, Berzelius voiced strong objections.Berzelius and his supporters Liebig andWöhler believed a negative chlorine atomcould not replace a positive hydrogen in anorganic compound, because this wouldmean there was an attraction between twonegative atoms: carbon and chlorine.

During the 1840s as more scientistsdemonstrated that substitution occurred inorganic chemical reactions, Berzelius’ elec-trochemical theory of organic radicals wasseverely questioned. This period also sawseveral alliances and bitter rivalries developamong the chemists attempting to develop acoherent theory of organic compounds. Amain source of controversy was a theorydeveloped by Laurent called the nuclear the-ory. Laurent believed organic compoundswere based on three-dimensional nucleiinstead of radicals. According to Laurent,atoms occupied positions on the corners andedges of geometrical shapes. Substitutioncould take place at these positions or addi-tion of other atoms took place on the face ofthe figure (Figure 15.1).

Dumas initially supported his studentLaurent, but then under pressure fromBerzelius and Liebig turned against Lau-rent. Wöhler and others ridiculed Laurent in

scientific journals. At the same time, Dumasdeveloped his own theory of organic substi-tution called type theory. Dumas’ theory hadmany similarities to Laurent’s nuclear the-ory. Laurent accused Dumas of stealing hisideas. Laurent, ostracized by the establishedchemical community and unable to obtainan academic position, joined forces withanother chemical outcast and brilliantchemist named Charles Gerhardt (1816–1856). Gerhardt’s idea on organic com-pound formation was called the theory ofresidues. According to Gerhardt, organicradicals decomposed into residues. Some ofthese residues recombined to form organiccompounds, while other residues formedstable inorganic compounds such as waterand ammonia.

The collaboration of Gerhardt and Lau-rent resulted in the type theory for organiccompounds. A type according to Gerhardtand Laurent was an inorganic compound thatproduced an organic compound when sub-stitution of a hydrogen atom in the inorganiccompound occurred. Gerhardt and Laurentdefined four types of organic compounds:water type, ammonia type, hydrogen type,and hydrogen chloride type. These four types

Figure 15.1August Laurent believed organic structurewas based on geometric nuclei. Substitutionand addition take place at differentpositions in the figure.


could be used to explain the formation of allorganic compounds. For example, alcoholswere a water type that formed when one ofthe hydrogen atoms in water was replaced bya radical, for example, methyl alcoholformed when the methyl radical replacedone of the hydrogen atoms in water: H2O �CH3 CH3OH. A vast array of organiccompounds such as ethers, aldehydes,ketones, amines, and so on were definedusing the four types. Other chemists pro-posed additional types to explain the ever-increasing array of organic chemicals beingdiscovered. Laurent and Gerhardt’s methodof classification for organic compounds wasbased on a specific type that unified a groupof organic compounds. This method formedthe basis of the modern classification oforganic chemicals based on functionalgroups. These functional groups are exam-ined later in this chapter. The death ofBerzelius in 1848 and the advancement ofthe work in organic chemistry contributed tothe acceptance of the work of Laurent andGerhardt. Unfortunately, both of thesechemists died just as their work was beingaccepted and they were being vindicated bythe established chemical community.

In the late 1850s, Frankland introducedthe concept of valence (see Chapter 7). Thisset the stage for Archibald Scott Couper(1831–1892) and Kekulé to explain thestructure of organic compounds. Couperwrote a paper in 1858 entitled On a NewChemical Theory. Couper’s paper showedthat carbon acted as a tetravalent (capable ofcombining with four hydrogen atoms) ordivalent atom (capable of combining withtwo hydrogen atoms). Couper also explainedhow tetravalent carbon could form chains toproduce organic molecules. He also pre-sented diagrams in his paper depicting thestructure of organic compounds and evenproposed ring structures for some organiccompounds. Unfortunately for Couper, the

presentation of his paper was delayed, and inthe meantime Kekulé’s similar, but lessdeveloped, ideas appeared in print. Kekulé,who originally studied architecture butswitched to chemistry, had close contactswith the great chemists of his day includingLiebig, Gerhardt, and Dumas. His work wasreadily accepted as a solid theory to explainorganic structure. Kekulé showed thatorganic molecules could be constructed bycarbon bonding to itself and other atoms. In1865, Kekulé explained how the propertiesof benzene (C6H6) could be explained usinga ring structure. Benzene’s structure came toKekulé during a dream in which he sawatoms as twisting snakes and was inspiredwhen one snake bites its own tail.

While Kekulé received the credit for themodern theory of tetravalent carbon duringhis day, Couper suffered from a mentalbreakdown and spent the last thirty years ofhis life in ill-health. It was not until the1900s after his death that chemists appreci-ated the depth of his work and acknowl-edged he deserved to share credit withKekulé for the theory of modern organicstructure. The work of Couper and Kekulémarked the start of modern organic theory.Radicals, types, and residues were no longerneeded to explain organic structure. Therevival of Avogadro’s hypothesis by Canniz-zaro enabled chemists to agree on atomicmasses, and the acceptance of atoms as thebuilding blocks of molecules advancedmodern organic theory. With this brief his-tory as a background, we can now examinethe basic groups of organic compounds,important organic reactions, and the roleorganic chemicals play in modern society.

Organic CompoundsOf all the existing chemical compounds,

the number that contains carbon is over-whelming. Due to the preponderance of


organic chemicals produced naturally, thenumber of organic compounds that exist cannot be determined. Approximately sixmillion organic compounds have beendescribed. The reason for the tremendousnumber of organic compounds is carbon’sability to covalently bond with itself andmany other atoms. This ability is due to car-bon’s electron structure. The electron con-figuration of carbon, atomic number 6, is 1s2

2s2 2p2. Carbon needs four more electrons tocomplete its outer shell and acquire a stableoctet electron configuration. Of all theatoms, only carbon has the ability to formlong chains of hundreds and even thousandsof carbon atoms linked together. The carbon-carbon bonds in these molecules exist as sta-ble single, double, and triple bonds. Carbonalso has the ability to form ring structuresand other atoms and groups of atoms can beincorporated into the molecular structuresformed from carbon atoms.

Organic compounds contain covalentbonds. In general, compared to inorganiccompounds organic compounds have lowmelting and boiling points, tend to beflammable, have relatively low densities, donot dissolve readily in water, and are pri-marily nonelectrolytes.

Because of the number and variety oforganic chemicals, an elaborate set of rulesexists for classifying and naming them.These rules have been established by IUPAC(International Union of Pure and AppliedChemistry). The IUPAC rules governing thenaming of organic compounds forms anentire language comparable to a foreign lan-guage. All you have to do is look at a labelof many commercial products to gain anappreciation for the variety and complexityof organic nomenclature. Like any lan-guage, the language of chemistry, and espe-cially organic chemistry, requires daily useto become proficient in its vocabulary andsyntax. Some of the major groups of organic

compounds will be examined with empha-sis on the most common compounds. Chem-ical nomenclature will be incorporated intoour discussion as needed. Common namesfor organic compounds will be usedthroughout our discussion, but some generalIUPAC rules will be given. These generalrules are helpful when you know the nameof a compound and want to know to whichfamily it belongs.

Aliphatic HydrocarbonsOrganic compounds containing only

carbon and hydrogen are known as hydro-carbons. There are several major groups ofhydrocarbons. Aliphatic hydrocarbons areprimarily straight and branched chains ofhydrocarbons consisting of alkanes,alkenes, and alkynes. Alkanes are hydro-carbons that contain only carbon-carbon(C––C) single bonds. When only singleC––C bonds exist in a hydrocarbon, it iscalled a saturated hydrocarbon. Unsatu-rated hydrocarbons contain at least onedouble or triple bond. Alkanes have thegeneral formula CnH2n�2, where n � 1,2,3. . . . The simplest alkane, defined when n� 1 in this equation, is CH4, methane. Table15.1 lists the first eight alkanes and givestheir molecular and structural formulas.The first four alkanes have common namesand the alkanes with more than four carbonatoms take the numerical Greek prefix fortheir root names.

The structural formulas shown in Table15.1 are written in their expanded form. Theexpanded form shows all atoms and allbonds in the formula. Structural formulascan also be written using a condensed form.In the condensed form, groups of atoms areused to convey how the molecule isarranged. For hydrocarbons, essential infor-mation is shown using the carbon-carbonbonds and grouping the hydrogen atoms

Table 15.1The First Eight Alkanes


with carbon atoms. For instance, the con-densed formulas of ethane and propane is:

Throughout this chapter, both expanded andcondensed structural formulas are used.Molecules may even be shown using boththe condensed and expanded forms for dif-ferent parts of the molecule. Rather thanadhere to using one structural formula, ourgoal is to represent the basic structure of themolecule.

Alkyl groups are a type of hydrocarbonformed by removing a hydrogen atom froman alkane. Their names are formed by usingthe alkane prefix and attaching “yl.” Thecorresponding alkyl groups for the firsteight alkanes are shown in Table 15.2.

The alkyl groups attach themselves toother groups, and their names often arise inorganic compounds, for example, methylalcohol and ethyl alcohol.

The structural formula for butane showsall the carbons linked in a straight chain. Inreality, the carbons actually line up in azigzag fashion, so the term “straight chain”simply refers to a continuous arrangementof carbon atoms. Butane and alkanes that

exist as straight chains are referred to asnormal alkanes. The butane in Table 15.1is called normal butane or n-butane. It isalso possible for butane to exist in abranched form:

This second branched molecule is calledisobutane. Compounds sharing the samemolecular formula but having differentstructures are called structural isomers.Normal butane and isobutane have differentphysical properties. The number of struc-tural isomers for the alkanes is included inTable 15.1. It can be seen in this table thatas the number of carbon atoms increasesthat the number of possible isomers alsoincreases. The fact that numerous isomersexist for most organic compounds is anotherreason why there are so many organic com-pounds.

Alkanes are relatively unreactive or-ganic compounds. They do undergo com-bustion, and several of the alkanes listed inTable 15.1 are familiar fuel sources. Natu-ral gas is primarily methane, propane is usedfor cooking and heating, butane is used inlighters, and octane is a principal compo-nent of gasoline. The octane number ofgasoline is an indication of the fuel’s abilityto resist premature ignition and cause knock in the engine. A fuel’s octane number isdetermined by comparing the fuel’s perfor-mance in a standard engine to that when an isooctane-heptane mixture is used in theengine. The isooctane used to determineoctane rating is 2,2,4 trimethylpentane:

Table 15.2Alkyl Groups

H3C—CH3 H3C—CH2—CH3

Ethane PropaneCH3

|H C—CH—CH3 3


The trade name Freon is used for the com-mon CFCs.

Alkenes and alkynes are unsaturatedhydrocarbons. Alkenes contain at least onecarbon-carbon double bond and alkyneshave at least one carbon-carbon triple bond.The names of alkenes and alkynes use thealkane prefixes, but add “ene” and “yne”endings, respectively (Table 15.3).

Alkenes have the general formulaCnH2n, where n � 2,3,4,. . . . Alkenes werecalled olefins, a term primaryly associatedwith the petrochemical industry. Alkyneshave the general formula CnH2n�2, where n� 2,3,4,. . . .

Alkenes are very reactive organic com-pounds. They readily undergo addition reac-tions in which atoms or groups of atoms areadded to each carbon in a double bond. Dur-ing this process, the double bond is con-verted into a single bond. In additionreactions, there is no loss of atoms from theoriginal compound. A common additionreaction is hydrogenation in which hydro-gen is added to each carbon in a doublebond. Hydrogenation occurs under increasedtemperature and pressure generally using anickel or platinum catalyst. Hydrogenationis used to convert unsaturated liquid vegeta-ble oils (cottonseed, corn, and soybean) intosaturated solid margarine, shortening, andvegetable spreads. Hydrogenation allowsmore healthy vegetable oils to be substitutedfor animal fats. If you examine the ingredi-ent label of many of these products, you will

This isooctane is one of the eighteen iso-mers of octane, and its octane number isset at 100. Normal heptane’s octane num-ber is 0. In general, straight chain alkaneshave low octane numbers and branchchain alkanes have high numbers. In fact,normal octane has an octane rating of –20. A gas with a 92 octane rating wouldhave an anti-knock performance equal to ablend that is 92% isooctane and 8% n-heptane. Often on gas pumps the for-mula R � M/2 is posted. The R and M inthis formula stand for research and motor,respectively. Research octane and motoroctane refer to different methods to deter-mine octane numbers. Motor octane isfound with the engine operating under aheavier load and running at higher rpm,more reflective of driving on a freeway.Research octane is found at lower rpm andbetter approximates acceleration fromrest. The octane rating of the gasoline isan average of these two as indicated by theformula.

Alkanes also undergo substitutionwhere hydrogen atoms in the alkane arereplaced by other atoms. Substitutionreactions are the most common type ofreaction in organic chemistry. Substitutionoccurs when an atom or group of atomsreplaces an atom or group of atoms in anorganic compound. Halogenation is acommon substitution reaction in whichhydrogen atoms are replaced by halogenatoms. One group of organic compoundsthat result from substitution reactions ischlorofluorocarbons. Chlorofluorocar-bons or CFCs have received widespreadattention in recent years because of theirrole in stratospheric ozone destruction.CFCs are commonly used as refrigerants.The two most common CFCs are tri-chlorofluoromethane and dichlorodifluo-romethane:

Cl Cl| |

Cl—C—F Cl—C—F| |Cl F

Trichlorofluoromethane Dichlorodifluoromethane

Freon-11 Freon-12


Other examples of addition polymerizationof alkenes are the production of polypropy-lene from propylene, polyvinyl chloride(PVC) from vinyl chloride, and Teflon fromtetrafluoroethylene. The structure of thethree monomers is depicted in Figure 15.2.

The characteristics of different polymersmay be modified by the inclusion of different

see the term “partially hydrogenated” refer-ring to this process.

Ethene or ethylene is the most impor-tant organic chemical used in commercialapplications. Annual production of ethylenein the United States was over twenty-fivemillion tons in the year 2000. Propylene isalso used in large quantities with an annualproduction of over thirteen million tons.Alkenes such as ethylene and propyl-ene have the ability to undergo addition polymerization. In this process, multiple addition reactions take place and manymolecules link together to form a polymer.A polymer is a long chain of repeating unitscalled monomers. For example, the addi-tion of two ethylene molecules can be rep-resented as

In the production of polyethylene, this pro-cess is repeated thousands of times. Thebasic repeating structure can be pictured as

Table 15.3Basic Alkenes and Alkynes

Figure 15.2The polymers polyethylene, polyvinylchloride, and Teflon are made by repeatedaddition of their respective monomers.


wraps, commonly called plastic wrap, aremade of LDPE.

A variety of polymers are classifiedunder the general heading of plastics. Whilethis classification may suffice for everydayusage, recycling of plastics requires them tobe separated according to their chemicalcomposition. Failure to separate plastics bychemical composition results in contaminat-ing one type of plastic with another. Becauseof this, international agreement requiresplastic containers greater than eight ouncesto contain a symbol designating their chem-ical composition. The symbol consists of atriangle made of three arrows containing anumber. The symbol is generally moldedright into the bottom of the container duringmanufacturing. Table 15.4 summarizes thesymbols, their composition, and some usesof several common polymers.

atoms or different reaction conditions, forexample, varying temperature and pressure.Additives affect properties such as flamma-bility, flexibility, thermal stability, and resis-tance to chemicals and bacteria. Oneexample of the difference in polymer char-acteristics is demonstrated by two main typesof polyethylene. High density polyethylene(HDPE) consists mostly of straight chainmolecules, while low density polyethylene(LDPE) consist of molecules containingmany branches. For this reason, HDPEmolecules can pack together more tightlyproducing a hard rigid material. HDPE plas-tics are use to produce materials where flex-ibility is not critical such as T.V. cabinets orrigid bottles. Conversely, the branching inLDPE plastics results in a significantly moreflexible material. Items such as plastic gro-cery bags, squeeze bottles, and various food

Table 15.4Plastic Recycling Symbols


Aromatic Hydrocarbons—Benzene

Aromatic hydrocarbons are unsatu-rated cyclic compounds that are resistant toaddition reactions. The aromatic hydrocar-bons derive their name from the distinctiveodors they exhibited when discovered. Ben-zene is the most important aromatic com-pound. Because many other aromaticcompounds are derived from benzene, it canbe considered the parent of other aromaticcompounds. Benzene molecular formula isC6H6 and its structural formula is

This structure is usually represented with-out the carbon and hydrogen atoms in eitherof two ways as shown in Figure 15.4.

Benzene was discovered in 1825 byMichael Faraday who identified it from aliquid residue of heated whale oil. Faradaycalled the compound bicarburet of hydrogenand its name was later changed to benzin byEilhardt Mitscherlich (1794–1863) who iso-lated the compound from benzoin. Ben-zene’s formula indicates it is highlyunsaturated. This would suggest benzene

Approximately 90% of the plastic recy-cled by consumers falls into categories 1and 2.

Alkynes contain a triple bond. The sim-plest alkyne is ethyne (C2H2) and goes bythe common name acetylene:

Acetylene is the most widely used alkyne.Alkynes undergo the same reactions asalkenes. Because a triple bond connects thetwo carbon atoms, addition of an atom ini-tially forms a double bond. The addition ofa second atom converts the double bondinto a single bond. For example, hydro-genation of propyne to propane is repre-sented as:

The alkanes, alkenes, and alkynes dis-cussed thus far are acyclic. This means thatthe hydrocarbon chains are linear and donot form rings that close on themselves.Alkanes, alkenes, and alkynes have theability to form closed rings. At least threecarbon atoms are needed to form a ring sothe simplest cyclic alkane is a form ofpropane called cyclopropane. Several sim-ple cyclic alkanes are shown in Figure15.3.

Cyclic compounds have different gen-eral formulas compared to their acycliccounterparts. The general formula for cyclicalkanes for instance is CnH2n.

Figure 15.3Alkane, alkenes, and alkynes can fold themselves to form cyclic compounds. Those representedabove are cycloalkanes.


structure of benzene as a hybrid of the tworesonance structures. Each carbon atom in thebenzene ring is bonded to two other carbonatoms and a hydrogen atom in the same plane.This leaves six delocalized valence electrons.These six delocalized electrons are shared byall six carbon atoms. This is demonstrated bythe fact that the lengths of all carbon-carbonbonds in benzene are intermediate betweenwhat one expects for a single bond and a dou-ble bond. Rather than consider a structure thatexists as one resonance form or the other, benzene should be thought of as existing asboth resonance structures simultaneously. Inessence, 1.5 bonds are associated with eachcarbon in benzene rather than a single anddouble bond. Using the hexagon symbol witha circle inside is more representative of thedelocalized sharing of electrons rather thanusing resonance structures, although benzeneis often represented using one of its resonancestructures.

Because of its structure, benzene is avery stable organic compound. It does notreadily undergo addition reactions. Additionreactions involving benzene require hightemperature, pressure, and special catalysts.The most common reactions involving ben-zene involve substitution reactions. Numer-ous atoms and groups of atoms may replacea hydrogen atom or several hydrogen atomsin benzene. Three important types of sub-stitution reactions involving benzene arealkylation, halogenation, and nitration. Inalkylation, an alkyl group or groups substi-tute for hydrogen(s). Alkylation is the pri-mary process involving benzene in thechemical industry. An example of alkylationis the production of ethyl benzene by react-ing benzene with ethylchloride and an alu-minum chloride catalyst:

and other aromatic compounds should read-ily undergo addition reactions like thealiphatic compounds. The fact that benzenedid not undergo addition puzzled chemistsfor a number of years. In 1865 Kekulé, asdiscussed earlier in this chapter, determinedthe correct structure of benzene to explainits behavior. Kekulé proposed that benzeneoscillated back and forth between two struc-tures so that all carbon-carbon bonds wereessentially equivalent. His model for thestructure of benzene is represented in Fig-ure 15.5. In Figure 15.5, benzene is shownas changing back and forth between twostructures in which the position of the dou-ble bonds shifts between adjacent carbonatoms. The two structures are called reso-nance structures.

A true picture of benzene’s structure wasnot determined until the 1930s when LinusPauling produced his work on the chemicalbond. Benzene does not exist as either of itsresonance structures, and its structure shouldnot be considered as either one or the other. Amore appropriate model is to consider the

Figure 15.4Two Alternate Ways of RepresentingBenzene

Figure 15.5Kekulé proposed that the structure for ben-zene resonated between two alternate struc-tures in which the position of the double andsingle bonds switched positions.


In halogenation, halogen atoms are sub-stituted for hydrogen atoms. The process ofnitration produces a number of nitro com-pounds, for example, nitroglycerin. Nitra-tion involves the reaction between benzeneand nitric acid in the presence of sulfuricacid. During the reaction, the nitronium ion,NO2�, splits off from nitric acid and substi-tutes on to the benzene ring producingnitrobenzene:

Just as an alkyl group is formed byremoving a hydrogen atom from an alkane,removing a hydrogen atom from an aromaticring produces an aryl group. Removing asingle hydrogen atom from benzene pro-duces phenyl, C6H6. The phenyl groupbonds with carbon in many organic com-pounds. Benzene rings may fuse together toform fused ring aromatic compounds. Acommon fused ring is naphthalene. Naph-thalene is used in mothballs and in the man-ufacture of many chemical products.Naphthalene has the distinctive odor of anaromatic compound. Its structural formulais

Some fused benzene rings have been impli-cated as carcinogens (cancer-causingagents). Several of these are found intobacco smoke.

Millions of organic chemicals arederived from benzene. Many of these arefamiliar chemicals such as aspirin (Chapter13) and several already mentioned in this sec-tion. Figure 15.6 shows some other commonbenzene derivatives. Besides the benzene

derivatives mentioned in this section, count-less other products are based on the benzenering. Cosmetics, drugs, pesticides, andpetroleum products are just a few categoriescontaining benzene-based compounds. Thou-sands of new benzene compounds are addedannually to the existing list of millions.

The field of organic synthesis is likeworking with an atomic Lego set. Using justa few atoms, millions of compounds can bebuilt. This process has occurred naturallythroughout Earth’s history, and in the lastcentury chemists have been able to mimic it.This section dealt with the two main classesof hydrocarbons called the aliphatic and aro-matic hydrocarbons. In the next section theconcept of the functional group is examined.Functional groups are used to define organicfamilies.

Functional Groups andOrganic Families

A functional group is a small chemicalunit consisting of atoms or a group of atoms

Figure 15.6Several Common Benzene Derivatives


of methanol in humans results in the pro-duction of toxic formaldehyde.

Alcohols may contain more than onehydroxyl group. An alcohol containing twohydroxyl groups is called dihydric, and onewith three is called trihydric. Ethylene gly-col and glycerol (glycerin) are two examplesof common alcohols with more than onehydroxyl group (Figure 15.8). Ethylene gly-col is the primary ingredient used inantifreeze, and glycerol is used in soaps,shaving cream, and other cosmetic products.

Aromatic alcohols are called phenols.The simplest phenol, also called phenol,forms when a hydroxyl group replaces ahydrogen atom in the benzene ring. Phenol(carbolic acid) was used as an antiseptic inthe 1800s. Today other phenol derivativesare used in antiseptic mouthwashes and incleaning disinfectants such as Lysol. Phe-nols are easily oxidized, and this makesthem ideal substances to use as antioxidants.By adding phenols such as BHT (butylatedhydroxy toluene) and BHA (butylatedhydroxy anisole) to food, the phenols oxi-dize rather than the food.

Ethers contain an oxygen atom singlybonded to two hydrocarbon units. The func-

and their bonds that is part of a largerorganic molecule. The functional group isgenerally the part of the molecule wherereactions take place. Organic compounds inthe same family have the same functionalgroup. The functional group for alkenes isthe carbon-carbon double bond, and foralkynes it is the carbon-carbon triple bond.Aromatic compounds share the hexagonalring functional group (Figure 15.7). A num-ber of functional groups define several otherlarge families of organic compounds.

Alcohols contain the hydroxyl group,OH. An alcohol can be represented as R-OH, where R is the rest of the moleculeattached to the functional group (forinstance an alkyl or aryl group). Alcoholstake their name from the correspondingalkane molecule and using the ending “ol”.For example, methanol (CH3OH) is methanewith the hydroxyl group replacing one of thehydrogen atoms (Figure 15.8). Ethanol,often called ethyl alcohol, has the formulaC2H5OH. Ethanol results when sugar fer-ments and is sometimes referred to as grainalcohol. It is the alcohol found in alcoholicbeverages. The hydroxyl group gives alco-hols polar characteristics, and it is oftenused as a solvent in many chemical pro-cesses. Denatured alcohol refers to “spik-ing” ethanol with a substance making itunfit for human consumption. The sub-stances used to spike ethanol are calleddenaturing agents. A common denaturingagent for ethanol is methanol. The oxidation

Figure 15.8Some Common Alcohols

Figure 15.7Functional Groups Defining Alkenes,Alkynes, and Aromatic Hydrocarbons


odor and taste. More recently, EPA has listedMTBE as a potential carcinogen. While theEPA has not banned MTBE use, many statesare regulating and even eliminating its usein gasoline. California, for instance, bannedits use after December 31, 2002.

Aldehydes and ketones both containthe carbonyl functional group:

Aldehydes have at least one of their hydro-gen atoms bonded to the carbon in the car-bonyl group. In ketones, the carbonylcarbon bonds to other carbon atoms and nothydrogen atoms (Figure 15.9).

Aldehydes use the parent name of theircorresponding alkanes and the ending “al.”Therefore, the simplest aldehyde is meth-anal, but it goes by its common nameformaldehyde. Aldehydes and ketones areoften produced by the oxidation of alcohols.For example, formaldehyde is produced bythe oxidation of methanol according to thefollowing reaction:

Formalin is an aqueous solution of ap-proximately 40% formaldehyde. It was traditionally used to preserve biologic specimens, but its identification as a mildcarcinogen has curtailed its use as a preser-vative. Formaldehyde is widely used to pro-duce synthetic resins. Resins are sticky,liquid organic compounds that are insolublein water. They often harden when exposedto air. Many commercial types of glue areresins. Natural resins are produced by plantsas a response to damage. When a plant suf-fers external damage, natural resins flow tothe area and harden to protect the underlying

tional group is the oxygen in the general for-mula R�-O-R. The R and R� in this generalformula may be the same or different. Ethersare identified by naming the attached Rgroups follow by the word “ether.” Forexample, CH3––O––CH2––CH3 would bemethyl ethyl ether. The most common etheris diethyl ether: CH3––CH2––O––CH2––CH3. This is sometimes referred to as ethylether or simply ether. Ether has been used asa general anesthetic since 1846 when Wil-liam T. G. Morton, a Boston dentist, demon-strated its use at Massachusetts GeneralHospital. Morton anesthetized a patient whowas having a tumor surgically removed fromhis jaw. The demonstration was so success-ful that ether was hailed as a miraculoussubstance that could alleviate pain. Todayether has largely been replaced by modernanesthetics. One of the major problems withether is its extreme flammability.

An ether called methyl tert butyl ether(MTBE) has traditionally been used as afuel oxygenate to reduce toxic air emissionsfrom automobiles:

Oxygenates are added to gasoline to givemore complete combustion. This in turnreduces the amount of carbon monoxideproduced. MTBE was widely used through-out the last part of the twentieth century forthis purpose, but environmental concernshave curtailed its use in recent years. Theproblem stems from contamination ofgroundwater from leaking undergroundstorage tanks and surface water from boats.At first, concerns revolved around odor andtaste problems. MTBE has a low concentra-tion threshold (20 to 40 parts per billion) for

CH3

|H C—O—C—CH3 3

|CH3

MTBE

O| |

—C—


Acetone is widely used as an organic sol-vent, as a paint remover (nail polish removeris mostly acetone), and to dissolve syntheticresins.

Acetone and several other ketones areproduced in the liver as a result of fatmetabolism. Ketone blood levels are typi-cally about 0.001%. Lack of carbohydratesin a person’s diet results in greatermetabolism of fats, which leads to a greaterconcentration of ketone in the blood. Thiscondition is called ketosis. People on lowcarbohydrate diets and diabetics have prob-lems with ketosis because of their increasedependence on fats for metabolism.

Carboxylic acids are organic acids thatcontain the carboxyl functional group,––COOH––. The carboxyl group is a com-bination of the carbonyl group and thehydroxyl group:

IUPAC names of carboxylic acids comefrom their parent alkane and use the ending“oic.” The simplest carboxylic acid containstwo carbons, and therefore, is calledethanoic acid, CH3COOH. Many carboxylicacids, though, are known by their commonnames, which reflects their natural origin.Some of these traditional names date backhundreds of years. The simplest carboxylicacid (IUPAC name methanoic acid) isformic acid. “Formica” is the Latin word for“ant.” Formic acid, which is responsible forthe stinging sensation of red ant bites, wasoriginally isolated from ants. Acetic acid(ethanoic acid) comes from the Latin word“acetum” meaning “sour” and relates to thefact that acetic acid is responsible for thebitter taste of fermented juices. Pure aceticacid is sometimes called glacial acetic acid.

structure. An interesting note is that theresin lac, which is used in shellac, is a resinsecreted by the insect Laccifer lacca. Somefamiliar natural resins are aloe, amber, andmyrrh. Synthetic resins include a variety offamiliar materials such as polystyrene, ure-thane, and epoxies.

The widespread use of formaldehyde toproduce resins has led to some concernsabout its affect on indoor air quality. Resinsare used in plywood, particleboard, syntheticwoods, and other building materials. Thesweet pungent odor of formaldehyde canoften be smelled in a lumber yard. Whenbuilding materials containing formaldehydeare incorporated into structures, theformaldehyde enters the air through a processcalled outgassing. Formaldehyde has beenidentified as a mild carcinogen, and peoplesusceptible to this chemical may experiencea variety of reactions including rashes andflu-like symptoms. The use of formaldehydefoam insulation, which was used widely inhomes, was discontinued in 1983.

The functional group for ketones showsthat the basic ketone contains three carbonatoms. Ketones have the ending “one” so thesimplest ketone is called propanone, but itgoes by the common name acetone:

Figure 15.9Both aldehydes and ketones contain the car-bonyl group, but in aldehydes, the carbonylcarbon bonds to at least one hydrogen atom.In ketones, the carbonyl carbon bonds totwo other carbons.


ganate (KMnO4). For example, ethanol canbe converted to acetic acid:

Similar to inorganic acids, the reaction ofcarboxylic acids and bases produces car-boxylic acid salts. Several of these salts arecommonly used in foods and beverages aspreservatives. The most common are saltsfrom benzoic, propionic, and sorbic acids.The salts of these acids have names endingwith “ate,” and can often be found in the listof ingredients of baked goods and fruitdrinks. Several common preservatives areshown in Figure 15.11.

Fatty acids are carboxylic acids con-taining an unbranched carbon chain andusually an even number of carbon atoms.Fatty acids do not occur freely in nature, butgenerally come from esters (esters are dis-cussed later). A few common fatty acids andtheir sources are shown in Figure 15.12.Fatty acids are important in the production

This is because pure acetic acid freezes at17°C, which is slightly below room temper-ature. When bottles of pure acetic acid frozein cold labs, snow-like crystals formed onthe bottles; the term glacial was thus asso-ciated with the pure form. Butyric acidcomes from the Latin word “butyrum” for“butter”; butyric acid is responsible for thesmell of rancid butter.

As discussed in Chapter 13, organicacids are some of the oldest known com-pounds. Salicylic acid is found in the plantfamily Salix and is the acid associated withaspirin (see Chapter 13). Several otherfamiliar carboxylic acids are acetic acid, car-bonic acid, and citric acid (Figure 15.10)More than one carboxyl group may be pres-ent in carboxylic acids, for example, citricacid contains three carboxyls. Oxalic acid,associated with the plant genus Oxaliswhich contains cabbage, rhubarb, andspinach, is a dicarboxylic acid containingjust two carboxyl groups bonded together:

Oxalic acid is poisonous to humans, but itsconcentrations are generally too low infoods to be of concern, although rhubarbleaves are quite poisonous. Lactic acid isproduced from the fermentation of lactose,which is the principal sugar found in milk.The taste and smell of sour milk is due tothe production of lactic acid from bacterialfermentation. Lactic acid accumulates inour muscles during exercise and strenuousphysical activity. It is responsible for thesore, aching feeling often associated withthese activities. Benzoic acid is the simplestaromatic carboxylic acid.

Carboxylic acids are prepared by oxi-dizing alcohols and aldehydes using an oxidizing agent such as potassium perman-

Figure 15.10Some Common Carboxylic Acids

KMnO4

CH CH OH ——r CH COOH3 2 3


Esters are named by first using the namederived from the alcohol and then the stemof the acid name with the ending “ate.”Acetate is derived from the common nameacetic acid. The IUPAC name would bemethyl ethanoate, with the ethanoatederived from acetic acid’s IUPAC nameethanoic acid.

Many of the fragrances and tastes fromplants are due to esters. The smells we per-ceive are generally due to a combination ofesters, but often one ester fragrance willdominant. By mimicking these naturalesters, the food industry has synthesizedhundreds of different flavoring agents. Threeof these are shown in Figure 15.14. Manypheromones are esters. Pheromones arechemical compounds used by animals forcommunication. Many medications are alsoesters. Aspirin is an ester of salicylic acid(see Chapter 13).

Esters can react with water in the pres-ence of acid catalysts to produce an alcohol

of soap. Before discussing soap production,let’s examine another family of organiccompounds called esters.

Esters are derived from carboxylic acidswhen these acids react with an alcohol in thepresence of strong acid catalysts. The pro-cess of ester formation is called esterifica-tion. In esterification, the hydrogen atom inthe alcohol’s hydroxyl group combines withthe hydroxyl portion of the carboxyl groupto form water. The remaining segments ofthe acid and alcohol form the ester. Figure15.13 represents the esterfication process.The functional group of esters is a carbonbonded to two oxygen atoms, one with adouble bond and one with a single bond:

Figure 15.11Some Common Carboxylic Acid Salts Usedas Preservatives

Figure 15.12Some common fatty acids and their sources.Note that the expanded structure is shownfor butyric acid and condensed structuresfor the other acids.

Figure 15.13In esterfication, an alcohol reacts with a car-boxylic acid to produce an ester and water.The resulting ester contains an alcohol partand an acid part.


fatty acids then react with the base to pro-duce a carboxylic acid salt known com-monly as soap (Figure 15.16).

The cleaning power of soaps can beattributed to the polar characteristics of thesoap molecule. Long chain hydrocarbons arerepresented by the “R”s in Figure 15.16. Thisend of the molecule is nonpolar and tends todissolve in nonpolar oils and grease (remem-ber the rule from Chapter 11 that “like dis-solves like”). The other end of the soapmolecule is ionic and dissolves in water. Theprocess of how soap works can be thought ofas one end of the soap molecule attractingthe grease and oil particles and the other end

and a carboxylic acid. This process is calledhydrolysis. Hydrolysis is the reverse ofesterification. When hydrolysis takes placein a basic solution, a carboxylic acid saltforms in a process called saponification.Saponification is the procedure used to pro-duce soap. Saponification is a two-step pro-cess in which an ester is initially hydrolyzedinto a carboxylic acid and alcohol. Duringthe second step, the carboxylic acid formedreacts with a base, such as sodium hydrox-ide or potassium hydroxide, to form a car-boxylic acid salt. The traditional productionof soap involves the use of animal fats andvegetable oils. The type of soap produceddepends on both the type of fat or oil used(beef tallow, lard, coconut oil, palm oil, lin-seed oil, etc.) and the base. Fats and oils areesters formed from the combination of fattyacids and glycerols. Three fatty acids, eitherthe same or different ones, attach to thethree hydroxyl groups in the glycerol toform a triester (Figure 15.15). Numerous tri-esters may be formed in this process, andthe general term for these triesters is triglyc-erides. During saponification of animal andplant products, the triglycerides are con-verted back to fatty acids and glycerol. The

Figure 15.14Esters are used as flavoring agents. Notice the similarity in the structure of the three flavors.

Figure 15.15A triglyceride forms when three fatty acidsreact with glycerol. When a triglycerideforms, three molecules of water areproduced as a byproduct.


dissolving in water. This differential solubil-ity pulls the oil molecules apart and sus-pends them in water in a process calledemulsification. Because the oil and greaseare largely responsible for holding the dirt tothe fabric, once they have been broken up,the dirt also readily dissolves.

Soaps react with the calcium and mag-nesium ions in hard water to produce soapcurd that greatly reduces its effectiveness.The curds are actually insoluble calcium andmagnesium salts. Synthetic laundry deter-gents have replaced soap for cleaningclothes in the last half century. Syntheticdetergents are made from petroleum. Theywork like soap except they do not react withmagnesium and calcium ions to form insol-uble precipitates and salts.

Polymerization of esters to producepolyesters is an important commercial pro-cess. Polyethylene terephthalate or PET isone of the most common plastics used infood containers (Table 15.4). This ester isformed by the reaction of ethylene glycoland terephthalic acid (Figure 15.17). PETand other polyesters consist of esters linkedtogether. Notice that both terephthalic acidand ethylene glycol have two carboxyls andtwo hydroxyls, respectively. When a polyes-ter such as PET is formed, a monomer con-

sisting of the ester is sandwiched betweenthe remaining alcohol and carboxyl acidgroups. Another acid group can attach to thealcohol and another alcohol group canattach to the acid end (Figure 15.18). Twonew acid and alcohol ends are then availableto continue the polymerization process. PETgoes by a number of commercial names, themost generic of these is polyester. It is alsothe materials known as Dacron. Mylar isPET in the form of thin films. It is used inrecording tapes and floppy disks.

The families of organic compounds dis-cussed thus far have been limited to thosecontaining carbon, hydrogen, and oxygen.

Figure 15.16In saponification, triglycerides are heated in a basic solution to give glycerol and soap. R repre-sents long carbon chains.

Figure 15.17PET is a polyester formed from terphthalicacid and ethylene glycol.


Amines are nitrogen-containing compoundsthat contain the amino functional group, –NH2. Amines can be viewed as compoundsin which an organic group(s), such as analkyl or aryl group(s), is substituted for ahydrogen(s) in ammonia (NH3). A primaryamine is one in which only one of thehydrogen atoms in ammonia has been sub-stituted for, secondary amines have twohydrogen atoms substituted, and tertiaryamines have all three ammonia hydrogenatoms substituted.

Many amines, like other organic com-pounds, are known by their common names.Their IUPAC names follow that of namingof alcohols, but the ending amine is used. Acommon practice is to list the names of thegroups attached to the nitrogen followed bythe ending “amine.” Thus, CH3-NH2 wouldbe methylamine. The simplest aromatic iscalled aniline:

Amines may also be present in other ringstructures similar to the cyclohydrocarbons.Ring structures that contain atoms other

than carbon in the ring are called hetero-cyclic compounds. Heterocyclic amines arethe most common type of heterocyclic com-pounds.

A number of common amines arefamiliar to all of us. Alkaloids are nitrogen-containing compounds extracted fromplants. Caffeine is found in varying quanti-ties in coffee beans (1 to 2% dry weight),tea leaves (2–5%), and cola nuts (3–4%).Caffeine is generally associated with cof-fee, tea, and soft drinks, but it is also pres-ent in many medications such as coughsyrups. Nicotine (4–5% dry weight) comesfrom tobacco leaves, and cocaine is derivedfrom the coca leaves (Figure 15.19).Cocaine was used as a local anesthetic dur-ing the early twentieth century, but in thelast half cenutry other amine derivativessuch as novacaine and benzocaine replacedits use. Morphine, codeine, and heroin arealkaloids that come from opium, which isobtained from poppy plants.

Histamine is an amine found in the cellsof all animals (Figure 15.20). The release ofhistamine in the body is associated withinvoluntary muscle contraction and dilationof the blood vessels. Allergic reactionsrelease histamine in the body and producethe accompanying side effects associatedwith hay fever such as coughing, sneezing,and runny nose. These effects can be allevi-ated with the use of antihistamines.

Antihistamines block the release of his-tamines, but also tend to produce drowsi-ness.

Figure 15.18In the formation of PET, a monomer esterforms with a remaining alcohol (OH) endgroup and carboxyl group (COOH). Theserepresent points to form two more estersand the process continues building the poly-mer.

Figure 15.19Caffeine and Nicotine


Epinephrine, commonly referred to asadrenaline, is an amine secreted inincreased amounts during times of stress(Figure 15.21). Adrenaline increases theheart rate and blood pressure, releasessugar stored in the liver, and constrictsblood vessels. It is sometimes administeredto people in shock or during periods ofacute asthma attacks.

The organic families and compoundspresented in this section give only aglimpse of the vast array of organic chem-icals present in our world. While eachorganic family is characterized by a uniquefunctional group, many compounds containmore than one functional group and classi-fying a compound into a specific family canbe difficult. Table 15.5 summarizes thosefamilies presented in this chapter. It can beused for a quick reference to the basicorganic families.

PetroleumPetroleum is the principal starting mate-

rial for the vast array of organic chemicalsthat dominate modern society. Chemicalsproduced from petroleum are collectivelytermed petrochemicals. The most prevalentuse of petrochemicals is as a fuel source, butthey are also the basic building blocks ofplastics, synthetic fabrics, detergents, pesti-cides, and many other products. Petroleumis created over a process lasting millions ofyears as microscopic marine organisms aredeposited in sediments and transformed byhigh pressure and temperature into oil andgas. Over time, the trapped oil and gasmolecules migrate through sedimentaryrock formations to areas called reservoirswhere they become trapped. These reser-voirs may be close to or far beneath the sur-face of the Earth. In certain regions, oil andgas seeps to the surface. Throughout history,surface seeps have been a source of lubri-cants and tars used for caulking and water-proofing. Medicinal products were distilledfrom oil obtained from surface sources bythe alchemists. The industrial revolution cre-ated a greater demand for oil. In the mid-1800s, whale oil was the principal source ofoil used in lamps, candles, and industry. Thedifficulty of obtaining whale oil, and its lim-ited supply, made it and its derivatives quiteexpensive. This motivated the search foralternative sources of oil. It was known min-eral oil could be obtained from the Earth.Besides known surface seeps, oil was some-times found mixed with water when diggingwater wells. Individuals in both this countryand in Europe dug wells in hope of findingoil that could be distilled into kerosene forlamps. Edwin L. Drake (1819–1880) drilledthe first successful commercial well, strik-ing oil on August 27, 1859, in TitusvillePennsylvania. This marked the start of thepetrochemical industry.

Figure 15.20Histamine

Figure 15.21Adrenaline


Crude oil is a mixture of hundreds oforganic chemicals. Through a process offractional distillation, the different com-ponents of petroleum can be segregated(Figure 15.22). Fractional distillation, orrefining, involves separating the compo-nents according to their boiling points as thecrude oil is heated. The lightest components,

those containing between one and four car-bon atoms, are gases at room temperature.Methane, ethane, propane, and butane areused as fuels and are converted into otheralkenes such as ethylene and propylene forindustrial use. Propane (boiling point–42°C) and butane (boiling point 0°C) canbe liquified under pressure, and the term LP

Table 15.5Organic Families


is used for liquified petroleum. Propane isreadily available at many gas stations for usein barbecues and camp stoves. Butane isused in lighters. Compounds containingbetween 4 and 10 carbon atoms, C5 to C10,include those that form petroleum ethersand gasoline. These compounds are ob-tained at temperatures between approxi-mately 30°C and 200°C. The nextcomponent includes compounds rangingfrom C11 to C16. Kerosene, the primary prod-uct in this range, is obtained between tem-peratures of 175°C and 275°C. Heating oils,C14 to C18 are obtained between 250°C and350°C, and lubricating oils and paraffin waxare obtained above 300°C. The solid residueremaining after distillation consists of heavyasphalt and tar components.

It is interesting to note that refiners in the1800s were primarily interested in obtainingkerosene and heavier components. Beforethe auto industry and the development of themodern petrochemical industry, there wasrelatively little need for gasoline and its asso-

ciated hydrocarbons. One of the residuesobtained from refining oil is petroleum jelly.Petroleum jelly’s, or petrolatum’s, use as askin conditioner was developed by an earlypioneer in the oil industry named RobertChesebrough. Chesebrough noticed that oilworkers in western Pennsylvania applied athick oily residue deposited on the drillingrods to their skin as a salve. Chesebroughcollected the material and started to marketit under the name Vaseline. Evidently, Vase-line was derived from the words “vase” and“line” denoting the fact that Chesebroughstored the product in vases.

Other processes are combined with frac-tional distillation to obtain the multitude ofchemicals used in the petrochemical indus-try. Cracking is a process where largeorganic molecules are broken up into smallermolecules. Thermal cracking involves theuse of heat and pressure. Catalytic crackingemploys various catalysts to reduce theamount of heat and pressure required duringthe cracking process. Alkylation is a process

Figure 15.22Fractionating Column for Oil (Rae Déjur)


in which smaller molecules are combinedinto larger molecules. Alkylation is used toproduce branched hydrocarbons that havemuch higher octane numbers than straightchain hydrocarbons.

Organic ProductsThe accelerated use of synthetic chem-

icals in society, which commenced afterWorld War II, has brought about a concernfor their overuse. This has created a marketfor so-called natural or organic products; inreality, all food can be considered organic.The perception among many consumers ofthese products is that natural or organicproducts are safer and healthier, decreasedamage to the environment, and are morenutritious. Companies have capitalized onthese perceptions by packaging and adver-tising products as “natural” or “organic.”The labels of these products often portraythe perceived benefits by using a naturalscene or natural colors, or proclaiming thepresence or absence of certain substances.The words “includes” or “doesn’t include”are often prominently displayed to conveythe message that this product is acceptableor better than the competition.

While new markets have been createdand new companies founded to meet thedemand for natural products, the consumeris often misled or confused regarding theseproducts. When the ingredient labels ofmany advertised “natural products” areexamined, they are often found to include anumber of synthetic chemicals. The words“natural” or “organic” should not be inter-preted as meaning good or better, and thepresence of synthetic chemicals does notmean bad. There is no difference in chemi-cals derived from natural sources or pro-duced synthetically. A chemical is the samechemical whether it comes from a natural orsynthetic source. Synthetic ascorbic acid,

vitamin C, and natural ascorbic acidobtained from rose hips are identical.Although synthetic and natural chemicalsare identical, there may be an advantage inusing one over the other, for example, price,ease of preparation. It should also be notedthat individuals may be allergic to certainchemicals, both natural and synthetic, sothat it is important in many instances toknow the actual ingredients in a product.

The federal government takes an activerole in product labeling, and one area ofconcern during the last decade has been thatof organic foods. The organic food marketis a $9 billion industry in this country grow-ing at a rate of 20% per year (compare to theaverage retail growth of 3%). The OrganicFoods Production Act (OFPA) of 1990 waspassed to establish standards for the grow-ing, handling, processing, and distributionof organic food. A part of the OFPA isintended to address truth in advertisingregarding organic food.

During the decade that followed this act,the government has tried to clarify the mean-ing of the term “organic” in agriculture. Theregulations defining this act were set forth onApril 21, 2001. The industry has eighteenmonths to comply with the regulations. Thefact that it took the National Organic Stan-dards Board, a panel of fifteen individualsoverseeing the implementation of the OFPA,eleven years to develop the standards indi-cates the difficulty in defining what it meansto be “organic.” During the 1990s, over fortyorganizations had their own criteria fororganic food. These organizations haddiverse interests that included farmers, han-dlers, processors, and distributors.

According to the 2001 regulations,organic is taken to mean grown and preparedwithout the use of synthetic chemicals,sewage sludge (as fertilizers), or ionizingradiation. Synthetic pesticides, defoliants,fertilizers, and herbicides are forbidden.


Instead, natural methods such as tillage,mulching, crop rotation, mechanical weed-ing, and natural pest management must beemployed. The regulations would lead one tobelieve that foods labeled as “organic” con-tain no synthetic chemicals, but there areexemptions and loopholes that allow syn-thetic chemicals to be used. Specifically,thirty-five synthetic chemical additives andstabilizers are exempt from regulation.Another aspect of the regulations regardshow soon organic crops can be planted onsoil to which synthetic chemicals had previ-ously been applied. The new regulationsspecify a three-year time span has to pass.Opponents of this rule argue that certain syn-thetic chemical residues do not degrade inthat time and remain in the soil after threeyears. To make sure that organic growers areadhering to the regulations, they must be cer-tified by United States Department of Agri-culture (USDA) certified agents.

The most apparent aspect of the OFPAto consumers regards product labeling. ByOctober 2002, USDA labels will specifywhether a food is organic. To be labeledorganic, the food must be made with at least95% organic material. Percentages can beincluded, so some producers may opt tolabel their foods as 100% organic. Whenfoods are produced with at least 70%organic material, the label “made withorganic material” will be used.

The OFPA also regulates animal pro-duction. Feed for organic livestock must be100% organic. No hormones or antibioticsare allowed. Certain vaccinations are per-mitted. If an animal is treated for a condi-tion with medication, it is no longerconsidered organic.

The OFPA defines organic food accord-ing to the USDA. There are calls by certaingroups, such as those producing naturalproducts, that similar regulations are neededin the cosmetic and health care industry. Sofar the Food and Drug Administration(FDA) has shown no inclination to followthe lead of the USDA to establish organicstandards for cosmetics and health careproducts.

SummaryThe material in this chapter traced the

history of organic chemistry from Wöhler’ssynthesis of urea through Kekulé’s structureof benzene. The millions of organic chem-icals known to exist can be classified into arelatively small number of families, eachdefined by a common functional group.During the last century, chemists have dis-covered how to mimic nature to synthesizeorganic chemicals. A multitude of familiarproducts are natural or synthesized organicchemicals.

Although organic substances withunique properties tailored to meet the needsof the producer dominate modern society,new organic compounds that are safer,stronger, cheaper, and perform better arecontinually being sought. Some of theorganic substances produced have createdunpredictable health and environmentalconsequences, but that is part of the pricefor the benefits accrued from these sub-stances. As we search for cures to diseases,seek safer and better products, and try tocorrect our past mistakes, organic chemicalsand organic chemistry will continue to playa vital role in modern society.

IntroductionBiochemistry is the study of chemicals asso-ciated with life and how these chemicalsinteract. Living matter consists primarily ofthe elements carbon, hydrogen, oxygen, andnitrogen. Many other elements, such as sul-fur and phosphorus, are essential to life, butthey do not comprise a large portion of livingmatter. For instance, the elemental mass com-position of humans is approximately 65%oxygen, 18% carbon, 10% hydrogen, 3%nitrogen, 1.6% calcium, and 1.2% phospho-rus, with the rest of the elements found in thehuman body totaling the remaining 1.2%.Elements combine to form the molecules thatmake up living cells. The molecules associ-ated with life are called biomolecules, andthe main categories of biomolecules includecarbohydrates, proteins, lipids, and nucleicacids. Biomolecules form cells that compriseliving matter, while biomolecular reactionsprovide the energy needed to sustain life.Biomolecules are constantly broken down ina process called catabolism and put togetherthrough anabolism as organisms carry outlife processes such as growth, locomotion,and reproduction. The sum total of catabolicand anabolic reactions is known as metab-

olism. In this chapter, we examine the mainclasses of biomolecules and emphasize theirrole in human metabolism.

CarbohydratesCarbohydrates are the predominant

class of biomolecules. Carbohydrates derivetheir name from glucose, C6H12O6, whichwas considered a hydrate of carbon with thegeneral formula of Cn(H2O)n, where n is apositive integer. Although the idea of waterbonded to carbon to form a hydrate of car-bon was wrong, the term “carbohydrate”persists. Carbohydrates consist of carbon,hydrogen, and oxygen atoms, with the car-bon atoms generally forming long un-branched chains. Carbohydrates are alsoknown as saccharides, derived from theLatin word for sugar “saccharon.”

Carbohydrates contain either an alde-hyde or a ketone group and may be eithersimple or complex. Simple carbohydrates,known as monosaccharides, contain a sin-gle aldehyde or ketone group and cannot bebroken down by hydrolysis reactions. Glu-cose and fructose are simple carbohydrates(Figure 16.1).˚ Glucose contains an alde-hyde group and fructose a ketone group.

16

Biochemistry

222 Biochemistry

Although shown as open chain struc-tures in Figure 16.1, most common sugarsin nature exist as closed five- and six-carbonringed structures known as pentoses andhexoses, respectively. Figure 16.2 illustratesthe common ring structures for glucose andfructose. The glucose and fructose struc-tures shown in Figure 16.2 are only one ofseveral forms possible depending on thearrangement of atoms in the molecule. Theforms shown in Figure 16.2 are the betaforms and are referred to as �-glucose and�-fructose, respectively. There are alsoalpha forms as shown in Figure 16.3. Noticethat the alpha forms are identical to the beta

forms except that the atoms on the right sideof the molecules have been flipped. Addi-tionally, carbohydrates, as well as otherbiomolecules, may exist as L and D forms.The “L” and “D” come from the Latin lae-vus and dexter meaning left and right,respectively. The L and D refer to mirrorimages of each molecule and are related tothe concept of chirality. A chiral object isone in which mirror images cannot besuperimposed on one another. In achiralobjects, mirror images are superimposable.For example, consider the mirror images ofthe heart and moon shown in Figure 16.4.The heart is achiral because the two imagesare identical when superimposed on eachother. The moon is chiral, because the

Figure 16.1Glucose and fructose are simplecarbohydrates. Notice that in glucose thecarbonyl C is bonded to a hydrogen atomcharacteristic of an aldehyde, while in fruc-tose the carbonyl C is bonded to two carbonatoms characteristic of a ketone.

Figure 16.2Most carbohydrates exist as closed ringsrather than open chains.

Figure 16.3Alpha Forms of Glucose and Fructose

Figure 16.4The heart is achiral and the moon is chiral.The mirror images of chiral objects do notcoincide when superimposed on each other.

Biochemistry 223

moon’s mirror image does not coincide withthe original object when superimposed on it.

Chiral molecules are often referred to asleft handed and right handed. The conceptof chirality can be seen if you hold your twohands in front of you with palms facingaway. The left and right hands are similarand can be considered mirror images, butwhen you place one hand on top of theother, they do not match. In a similar fash-ion, molecules can have identical formulasbut differ according to how they are orientedin space. Molecules that differ according totheir arrangement in space are calledstereoisomers. Stereoisomers can react dif-ferently and display different solubilities.This has significance in biochemical reac-tions in which one form of the isomer maybe biologically active and another form maynot, for example, enzymes. The sugarsshown in Figures 16.2 and 16.3 are D forms.We will not concern ourselves with the rulesfor determining D and L forms of carbohy-drates and other biomolecules, and will generally refer to biomolecules without ref-erence to their specific form. You should beaware, though, that other books might usethe designations �, �, D, and L when nam-ing biomolecules.

D-Glucose is the most important andpredominant monosaccharide found innature. It also goes by the names dextroseand blood sugar. The term “blood sugar”indicates that glucose is the primary sugardissolved in blood. Glucose is the principalenergy source derived by cells. D-fructose,as its name implies, is abundant in fruits.Honey also has a lot of fructose. Invertsugar is a mixture of glucose and fructose.

Two monosaccharides bonded togethercomprise a disaccharide. Three commondisaccharides are sucrose (common tablesugar), maltose, and lactose. Sucrose iscomposed of glucose and fructose. Maltoseis composed of two glucose molecules. Lac-

tose is made up of glucose and galactose.When sucrose, maltose, and fructose arehydrolyzed, the monosaccharides that com-prise these disaccharides are obtained:

sucrose � H2O glucose � fructosemaltose � H2O glucose � glucoselactose � H2O glucose � galactose

Hydrolysis is very important in biochemicalreactions and refers to a reaction in which asubstance reacts with water causing the sub-stance to break into two products. The struc-ture of common table sugar, sucrose, isshown in Figure 16.5.

Polysaccharides are composed of manymonosaccharides bonded together. Com-mon polysaccharides are cellulose, starch,and glycogen. Cellulose forms the structuralmaterial of the cell walls of plants. Celluloseis a polymer of glucose and consists of thou-sands of glucose molecules linked in anunbranched chain (Figure 16.6).

Humans do not produce the enzymes,called cellulases, necessary to digest cellu-lose. Bacteria possessing cellulase inhabitthe digestive tracts of animals such as sheep,goats, and cows giving these animals theability to digest cellulose. Cellulase bacte-ria also exist in the digestive systems of

Figure 16.5Common table sugar, sucrose, is a disaccha-ride composed of a hexose form of glucosebonded to a pentose form of fructose.

224 Biochemistry

certain insects, such as termites, allowingthese insects to use wood as an energysource. Although humans cannot digest cel-lulose, this carbohydrate plays an importantrole in the human diet. Carbohydrates thathumans cannot digest are called fiber orroughage. Fruit, vegetables, and nuts are pri-mary sources of fiber. These foods seem tohave a cleansing effect on the large intestine,and such action is thought to speed the pas-sage of cancer-causing materials through theintestine and reduce the risk of colon can-cer. Fiber helps retain water in the digestivesystem, aiding the overall digestion process.Another benefit of fiber is believed to be itsability to lower blood cholesterol, reducingheart and arterial disease.

Starch is a �-glucose polymer with thegeneral formula (C6H12O5)n, where n can bea number from several hundred to severalthousand. Starch is the form that plants storecarbohydrate energy. Grains, such as rice,corn, and wheat; potatoes; and seeds are richin starch. Two principal forms of starchexist. One form is amylose and it comprisesabout 20% of all starches. Amylose is astraight chain form of starch containing sev-eral hundred glucose units. In the otherform, called amylopectin, numerous glu-cose side branching is present. Starch is theprimary food source for the human popu-lation, comprising nearly 70% of all foodconsumed. Starch is broken down in thedigestive system starting in the mouth where

the enzyme amylase breaks the bonds hold-ing the glucose units together. This processceases in the stomach because stomach acidcreates a low pH environment, but resumesin the small intestine. Glucose, maltose, andother small polysaccharides result from thedigestion of starch.

In animals, most glucose is producedfrom starch and is not immediately neededfor energy; therefore, it is stored as glyco-gen. Because glycogen’s function in animalsis similar to that of starch in plants, it issometimes referred to as animal starch.Glycogen is similar in structure to amy-lopectin, but it is larger and contains moreglucose branching. When glucose blood lev-els drop, for instance during fasting or phys-ical exertion, glycogen stored in the liverand muscles is converted into glucose andused by cells for energy. During periods ofstress, the release of the hormone epi-nephrine (adrenaline) also causes glucose tobe released from glycogen reserves. Whenblood glucose is high, such as immediatelyfollowing a meal, glucose is converted into glycogen and stored in the liver andmuscles.

The transformation of glucose to glyco-gen and glycogen back to glucose enableshumans to regulate their energy demandsthroughout the day. This process is disruptedin individuals who have a form of diabetesknown as diabetes mellitus. Diabetes melli-tus occurs in individuals whose pancreas

Figure 16.6Cellulose consists of repeating glucose units bonded together.

Biochemistry 225

produces insufficient amounts of insulin orwhose cells have the inability to use insulin.Insulin is a hormone responsible for signal-ing the liver and muscles to store glucose asglycogen. In Type 1 diabetes, called insulin-dependent diabetes mellitus (IDDM), thebody produces insufficient amount ofinsulin. This type of diabetes occurs inyouth under the age of 20 and is the lessprevalent form (about 10% of diabeticscarry this form). Type 1 diabetes is con-trolled using insulin injections and regulat-ing the diet. Type II diabetes, known asinsulin-independent diabetes mellitus(IIDM), is the most common form of dia-betes and is associated with older individu-als (generally over 50) who are overweight.In this form of diabetes, individuals produceadequate supplies of insulin, but the cells donot recognize the insulin’s signal, and there-fore, do not capture glucose from the blood.This type of diabetes is regulated with drugsand a strictly controlled diet.

In recent years, the consumption of car-bohydrates in the form of refined sugars hasreceived significant attention from thehealth professionals. The annual consump-tion of sugar in the United States is about 50pounds per person or just under one poundper week. The adverse effects of excessivesugar in the diet include obesity, increaserisk of heart disease, diabetes, tooth decay,and disruptive behavior such as hyperactiv-ity in children. Because of these problems,the food industry has employed a number ofsynthetic or artificial sweeteners in place ofsugar. While these artificial sweeteners mayreduce the use of sugars, they have also beenlinked to adverse health problems such ascancer. Saccharin is the oldest syntheticsweetener. It was discovered in 1879 and hasbeen used since 1900. During the 1970s,rats feed large doses of this substance devel-oped bladder tumors, and the federal gov-ernment proposed banning this substance.

Public resistance prevented its ban, butfoods containing saccharin must carry awarning label. Aspartame, known com-monly as Nutra-Sweet, is the predominantsynthetic sweetener in use today. It consistsof two amino acids: aspartic acid and phe-nylalanine (see section on proteins). It isapproximately 160 times sweeter thansucrose. The health issues associated withaspartame are associated with the diseasephenylketonuria (PKU). PKU results whenexcess amounts of phenylalanine occur inindividuals due to the lack of the enzymeneeded to catalyze the essential amino acidphenylalanine. This genetic condition, if nottreated, causes brain damage resulting inmental retardation. Because aspartame con-tains phenylalanine, it is suggested thatinfants under two years of age not ingest anyfoods containing synthetic sweeteners.Cyclamates were once a popular sweetenerused in this country, but they were bannedin 1969 because of possible carcinogeniceffects. Their use is permitted in Canada. A synthetic sweetener approved for use in the United States in 1998 is sucralose.Sucralose is a chlorinated substituted chem-ical prepared from sucrose itself. It is 600times sweeter than sucrose, passes unmodi-fied through the digestive system, and unlikeaspartame, is heat stable so it can be used incooking.

Although the use of synthetic sweeten-ers is increasing annually, the quantity ofsucrose is also increasing. This apparentcontradiction seems to be related to theincreased production of low fat foods inrecent years. It seems the fat calories ofmany low-fat foods have been replaced bysugars. Questions persist regarding thesafety of artificial sweeteners, but the con-sensus is that moderate intake of these sub-stances do not pose a risk to human health.The federal government in approving syn-thetic sweeteners considers guidelines

226 Biochemistry

regarding safe daily limits. The accept-able daily intake (ADI) is considered theamount of sweetener that can be consumeddaily over a lifetime without adverse effects.For example, the ADI of aspartame is 50milligrams per kilogram of body weight. NoADI has been established for saccharin, butthe FDA limits the amount of saccharin perserving to 30 milligrams.

LipidsLipids are biomolecules characterized

by their solubility properties. Lipids aremolecules that are insoluble in water andsoluble in organic solvents. Triglyceridesare the major class of lipids, comprising95% of all lipids. Triglycerides include fatsand oils. Fats and oils are classified accord-ing to their state at room temperature. Fatsare solid at room temperature, while oils areliquid at room temperature. Two otherclasses of lipids are the phospholipids andthe steroids. Lipids serve a number of func-tions in the organisms. They provide anenergy source to supplement stored glyco-gen when glucose supplies are low, are prin-cipal components of cell membranes, formprotective waxy coatings in plants, andserve as the starting material for the synthe-sis of other biomolecules such as hormones.

Triglycerides are made from three fattyacids attached to a glycerol molecule. Thebasic structure of fatty acids and triglyc-erides was discussed in Chapter 15 in thesection on organic acids. The fatty acids ina triglyceride determine the triglyceride’snature. One important characteristic of fattyacids is the degree of saturation. Saturatedfatty acids contain only single bonds andcontain the maximum number of hydrogenatoms. A monounsaturated fatty acid hasa single double bond. The carbon atoms atthe double bond are called unsaturated.Polyunsaturated fatty acids have morethan one double bond per molecule, and at

each of these double bonds, the carbons areunsaturated (Figure 16.7). The melting pointof triglycerides is dependent on the degreeof saturation. Saturated triglycerides exist assolid fats. As the fat becomes unsaturated,its melting point decreases; therefore,polyunsaturated triglycerides exist as oilsrather than fats.

Health professionals concerned aboutcardiovascular disease recommend monoun-saturated and polyunsaturated fatty acids inthe diet as opposed to saturated fats. Oilsthat have high degrees of unsaturation areperceived as healthier compared to thosewith greater saturation. To rank the degreeof unsaturation, a measure called the iodinenumber can be used. The iodine number isthe mass of iodine in grams that can beadded to 100 grams of a triglyceride.Because the degree of unsaturation isrelated to the number of double bonds themolecule contains, and the double bondsrepresent points where iodine atoms can beadded, a higher iodine number means agreater degree of unsaturation. Table 16.1gives iodine numbers and relative percent-ages of the different types of fatty acids inseveral common fats and oils.

Figure 16.7Fatty acids can be saturated,monounsaturated, or polyunsaturateddepending on whether the molecule containsdouble bonds. Where the double bond(s)exists, the carbon atoms are unsaturated.

Biochemistry 227

Phospholipids are lipids that containphosphorus. They are similar in structure totriglycerides, except rather than having threefatty acids bonded to a glycerol one of thefatty acids in a triglyceride is replaced byphosphoric acid and an amino alcoholgroup. Replacing a single fatty acid with thephosphorus group changes the solubilitycharacteristics of the molecule. The fattyacid end of a phospholipid is soluble in fatsand oils, while the phosphorus end is solu-ble in water. The differential solubility char-acter of the opposite ends of phospholipidsmakes them ideal emulsifiers. An emulsifieris a substance that mixes with both fats andwater, and allows emulsion. Emulsion is theability of a lipid and water to mix. The abil-ity of phospholipids to mix in both waterand oil makes them important substancesfor transporting lipids across cell walls inorganisms. Phospholipids are a key sub-stance in the cell walls of plants and ani-mals. Common phospholipid groups found

in all living tissue are the lecithins. Lecithinsare widely employed in the food industry toassist in the mixing of fats and oil into water.Commercial lecithin is extracted from soy-beans and egg yolks, and can be found as aningredient in many food products, for exam-ple, margarine, candy bars, cookies, choco-late chips, and so on.

The third class of lipids is steroids.Included in this category of lipids arecholesterol, bile salts, and sex hormones.Steroid structures contain fused rings con-sisting of three six-carbon rings and a five-carbon ring:

Table 16.1Fatty Acid Composition and Iodine Numbers

228 Biochemistry

Cholesterol is a steroidal alcohol orsterol. (Figure 16.8). It is the most abundantsteroid in the human body and is a compo-nent of every cell. Cholesterol is producedin the liver, and a number of other sub-stances are synthesized from it includingvitamin D, steroid hormones, and bile salts.Cholesterol is commonly associated withcardiovascular disease. High blood serumcholesterol levels are often correlated withexcessive plaque deposits in the arteries, acondition known as arteriosclerosis orhardening of the arteries. Genetics plays asignificant role in blood cholesterol levels,but reducing the amount of saturated fat inthe diet can often control elevated choles-terol. Drugs are also employed to reducecholesterol levels. Bile is produced in theliver from cholesterol and stored in the gall

bladder. Bile acts as an emulsifying agentenabling triglycerides to mix with water inthe small intestine where enzymes act todigest fats. Steroidal hormones, includingthe sex hormones, are also steroids. The sexhormones fall into three major groups.These are estrogens, which are female sexhormones; androgens, which are male sexhormones; and progestins, which are preg-nancy hormones. The basic steroid structureis present in the three main hormones ineach of these groups (Figure 16.9).

Biochemists have succeeded in produc-ing synthetic steroids that have found wideuse in a number of applications. Oral con-traceptives used for birth control are basedon structures that resemble progesterone.Oral contraceptives were first marketed in1960. Synthetic progestins act by producinga state of false pregnancy that prevents thefemale user from ovulating. Synthetic sexhormones have also been employed, mostlyillegally, to boost athletic performance.Anabolic steroids are testosterone deriva-tives that serve to build muscle mass andspeed recovery from strenuous physicalactivity. The advantage gained by the use ofthese performance-enhancing drugs is notwithout cost. Side effects include cancer,heart disease, atrophy of testicles, disruption

Figure 16.8Cholesterol

Figure 16.9Estradiol, testosterone, and progesterone are the principal sex hormones from the estrogen,androgen, and progestin groups, respectively. These hormones are synthesized in the humanbody starting with cholesterol.

Biochemistry 229

of the menstruation cycle, kidney damage,and the appearance of masculine traits infemales.

ProteinsProteins are the third major class of

biomolecules. The word “protein” comesfrom the Greek word “proteios” meaningprimary. Protein is the primary materialcomposing cells. Fifty percent of the humanbody is composed of proteins, and it is esti-mated that roughly 100,000 different pro-teins are found in humans. Protein is a majorcomponent of structural and connective tis-sue found in skin, ligaments, bones, mus-cles, and tendons. Digestive enzymes,insulin, and other hormones are proteins.

Proteins are polymers made up ofamino acid units that form long chains.Amino acids are molecules that contain an amino group, NH2, and a carboxyl group,COOH, which characterize carboxylicacids. All amino acids contain a simplechemical backbone consisting of a carbonatom with an amino group, a carboxylgroup, and a side chain group attached tothe carbon:

The side chain differentiates one aminoacid from another. The simplest amino acidis glycine in which the side group consistsof a single hydrogen atom. Proteins found innature are composed of 20 standard aminoacids (Figure 16.10).

Humans have the ability to produceabout half of the twenty basic amino acids.The other ten that must be obtained fromfood are called essential amino acids.

A sequence of up to 50 amino acids iscalled a peptide, and the bonds between

individual amino acid units are called pep-tide bonds. Peptide bonds form between theamino group of amino acid and the carboxylgroup of an adjacent amino acid in a peptidechain. The process of peptide formation isillustrated in Figure 16.11, which showshow two amino acids bond together to forma dipeptide. Two amino acids form a dipep-tide, three form a tripeptide, several aminoacids make an oligopeptide, and more thanten comprise a polypeptide. Proteins can beconsidered polypeptides of more than 50amino acids. Peptides are named by usingthe abbreviations for their constituent aminoacids separated by dashes. By convention,the peptides start at the nitrogen-containingend and the sequence is read from left toright, for example, Tyr-Gly-Bly-Phe-Met.Numerous hormones secreted by the body’spituitary, hypothalamus, and endocrineglands are polypeptides. Insulin is a hor-mone that stimulates glucose uptake fromthe bloodstream. Another polypeptideglucagon stimulates the release of glucosefrom the liver. Other polypeptides releasegrowth hormones, stimulate production ofsteroidal hormones, and block nerve mes-sages to control pain.

When the number of amino acids in apolypeptide chain reaches more than fifty, aprotein exists. The structure of bothpolypeptides and proteins dictate how thesebiomolecules function. There are severallevels of structure associated with polypep-tides and proteins. The sequence of theamino acids forming the backbone of theprotein is referred to as the primary struc-ture. A different order or even a minorchange in an amino acid sequence creates anentirely different molecule. Just reversingthe order of amino acids in a dipeptidechanges how the dipeptide functions. Anexample of this is sickle-cell anemia.Sickle-cell anemia is a genetic disorder thatoccurs when the amino acid valine replaces

Figure 16.10The twenty basic amino acids that can be grouped to form proteins. Abbreviations are given inparentheses. Those shown in bold are essential amino acids for humans. Arginine is essential inchildren, but not adults.

Biochemistry 231

glutamic acid an eight sequence peptide.The presence of valine produces a form ofhemoglobin called hemoglobin S.

Sickle-cell anemia causes red blood cells toassume a sickle shape. The abnormal bloodcells disrupt blood circulation, which canlead to infections, organ damage, andchronic pain.

In addition to the primary structure,proteins also exhibit secondary, tertiary,and quaternary structure. The overall struc-ture of proteins is related to several factors.Primary among these factors is the electro-static nature of amino acids. The structuresdisplayed in Figure 16.10 do not show thecharge distribution displayed by aminoacids. In neutral solutions, the carboxylgroup tends to donate a proton (hydrogenion) to the amino group. The transfer of aproton means the amino end of the molecule

acquires a positive charge and the carboxylend a negative charge (Figure 16.12). Thetransfer of the proton within the moleculecreates a zwitterion, which is a moleculewith a positive charge on one atom in themolecule and a negative charge on anotheratom. The overall charge of the moleculeremains neutral. In acidic solutions, a pro-ton can add to the COO� creating a posi-tively charged ion, while in basic solutions aproton can be donated from the nitrogen endcreating a negatively charged ion. In a sim-ilar fashion, the side chains of certain aminoacids can take on a polar character and haveeither an acidic or a basic character.

The difference in the charge character-istics of different parts of an amino acidmolecule results in attractive and repulsiveforces within a polypeptide or protein. Theseforces cause the molecule to coil, fold, andacquire various shapes that produce differ-ent structural properties. These structuralproperties dictate, largely, how proteins func-tion. Secondary structure of a protein refersto how the atoms comprising the backboneare arranged in space around a central axis.Secondary structure results from the hydro-gen bonding between hydrogen atoms of theamino group and the oxygen on the carbonylgroup of another amino acid. Two commonsecondary structures exhibited by proteinsare alpha helices and beta sheets. An alphahelix structure resembles a coiled spring andslinky beta sheets arise between parallel pro-tein chains arranged along a sheet form-ing an alternating zigzag pattern (Figure

Figure 16.11The amino and carboxyl ends of two aminoacids unite to form a peptide bond. In theprocess, water is produced. Notice that thedipeptide form has an amino and a carboxylgroup at its end allowing other amino acidsto bond to the structure.

Figure 16.12In neutral solutions, amino acids acquire apositively charged nitrogen end and a nega-tively charged carbon end.

232 Biochemistry

16.13). Tertiary structure refers to the three-dimensional shape acquired by the proteindue to the forces between amino acid sidechains within a single polypeptide chain. Thetertiary three-dimensional structure resultsfrom the folding of polypeptide chains. Qua-ternary structure results from the interactiv-ity of different polypeptide chains within aprotein. The different levels of protein struc-ture can be compared to a telephone cord.The linear sequence of amino acids com-prises the primary structure, the coil the sec-ondary structure, and folding of the coilwhen the receiver is hung up gives the ter-tiary structure. The quaternary structurewould result when two or more cordsbecame intertwined. The overall structure of proteins leads to two general shapes.Fibrous proteins are long and slender fibers.These proteins are present in muscles, ten-dons, ligaments, and other structural tissue.The most prevalent protein in the humanbody is the fibrous protein called collagen.Collagen is a major constituent of tendon,skin, bone, ligaments, and cartilage. Boiling

it in water denatures it, and when allowed tocool, gelatin is produced. Globular proteinshave a spherical shape and move through thecirculatory system. Examples of globularproteins include insulin, hemoglobin, andmyoglobin.

As mentioned previously, structure iscritical to how proteins function. Denatura-tion refers to the process that disrupts pro-tein structure. Heat is a primary cause ofprotein denaturation. When eggs are cooked,the transformation of clear transparent eggwhite to a solid white material is due to thedenaturation of albumin. Heat causes thehydrogen bonds to break and unravels the globular albumin to produce the fibrouswhite cooked egg. Denaturation of proteinenzymes presents a severe problem whenbody temperature reaches levels above105°F. These temperatures inactivateenzymes that control body process, andmany systems no longer function. Thisproblem was tragically illustrated during theAugust 2001 training camp of the Min-nesota Vikings with the heat-related deathof an offensive lineman whose body tem-perature rose above 108°F. X-rays or otherforms of radiation may also cause denatura-tion. Damage from radiation may lead tocancer or genetic disorders. Yet, anothermeans of denaturation is due to toxic chem-icals such as benzene, heavy metals, or chlo-rinated hydrocarbons.

Proteins composed solely of amino acidunits are referred to as simple proteins.Many proteins contain other chemicalgroups in addition to amino acids. Theseproteins are called conjugated proteins. Pro-teins containing a carbohydrate unit arecalled glycoproteins, those with a nucleicacid are called nucleoproteins, and thosecombined with a lipid are called lipopro-teins. High-density lipoproteins (HDLs) andlow-density lipoproteins (LDLs) areincluded in a cholesterol test. LDLs, which

Figure 16.13The alpha helix and beta sheets are twotypes of secondary structures exhibited bypolypeptides and proteins. (Rae Déjur)

Biochemistry 233

are mainly lipids, carry fats and large quan-tities of cholesterol made in the liver to bodycells. HDLs, which are primarily protein,transport cholesterol back to the liver. LDLand HDL are often referred to as bad andgood cholesterol, respectively. Althoughboth these contribute to the overall bloodcholesterol, it is more important for indi-viduals to focus on how cholesterol is dis-tributed between HDL and LDL rather thana total cholesterol level. For this reason, theratio of either the total cholesterol or LDLto HDL is a better indicator of health riskthan the number indicating total cholesterol.For men an acceptable LDL:HDL choles-terol ratio is 4:5 and for women it is 3:5.

Nucleic AcidsNucleic acids are biomolecules that

pass genetic information from one genera-tion to the next. Nucleic acids contained inDNA, deoxyribonucleic acid, and RNA,ribonucleic acid, are responsible for how allhigher organism develop into uniquespecies. DNA is present in the nucleus of allcells where segments of DNA comprisegenes. DNA carries the information neededto produce RNA, which in turn producesprotein molecules. A simplified view of therole of nucleic acids is to produce RNA, andthe role of RNA is to produce proteins.

Just as the basic unit of proteins is theamino acid, the basic unit of nucleic acids isthe nucleotide. Nucleic acids are polymersof nucleotides. Nucleotides consist of threeparts: a sugar, hydrogen phosphate, and anitrogen or amine base. The sugar unit inDNA is 2-deoxyribose and in RNA it isribose. The “2-deoxy” means an oxygenatom is absent from the second carbon ofribose (Figure 16.14). Hydrogen phosphatebonds to the sugar at the number 5 carbon,and in the process a water molecule isformed. A nitrogen-containing base bonds

at the number 1 carbon position also to pro-duce water. The combination of the sugar,hydrogen phosphate, and base produces anucleotide (Figure 16.15).

The sugar and phosphate groups formthe backbone of a nucleic acid, and theamine bases exist as side chains. A nucleicacid can be thought of as alternating sugar-phosphate units with amine base projec-tions:

In DNA the sugar molecule in the nucleo-tide is 2-deoxyribose, and in RNA it isribose. The amine bases in DNA are ade-nine, thymine, cytosine, and guanine, sym-bolized by A, T, C, and G, respectively. RNAcontains adenine, cytosine, and guanine, butthymine is replaced by the based uracil (Fig-ure 16.16). The primary structure of nucleicacids is given by the sequence of the amineside chains starting from the phosphate endof the nucleotide. For example, a DNAsequence may be –T-A-A-G-C-T.

Like proteins, DNA and RNA exhibithigher order structure that dictates howthese molecules function. Determining thestructure of DNA challenged the world’sforemost scientists during the middle of the

Figure 16.14Shows that 2-deoxyribose and ribose areidentical except for an oxygen atom absenton the second carbon. Carbons arenumbered for future reference.

234 Biochemistry

strand. In DNA, cytosine is always hydro-gen bonded to guanine, and adenine isalways hydrogen bonded to thymine. Thefact that cytosine always lies opposite gua-nine and adenosine always opposite thyminein DNA is referred to as base pairing, andthe C-G and A-T pairs are referred to ascomplementary pairs. Base pairing can becompared to two individuals shaking hands.Both individuals can easily shake handswhen both use either their right or lefthands. However, when one person extendsthe left hand and another the right, the handsdo not match. In a similar fashion, the basepairing in DNA is necessary to produce a

twentieth century. The puzzle of DNA’sstructure was solved by James Watson(1928–) and Francis Crick (1916–) in 1953.Watson and Crick arrived at a double helixmodel consisting of two nucleotide strandsheld together by hydrogen bonds (Figure16.17). One way to picture the DNAmolecule is to consider a ladder twisted intoa helical shape. The sides of the ladder rep-resent the sugar-phosphate backbone, whilethe rungs are the amine base side chains.The amine bases forming the rungs of theladder exist as complementary pairs. Theamine base on one strand hydrogen bonds toits complementary base on the opposite

Figure 16.15A nucleotide is formed when hydrogen phosphate derived from phosphoric acid combines witha sugar and nitrogen base. Several different nitrogen bases may form a nucleotide, and thehexagon is used as a general symbol for any one of these.

Figure 16.16The amine bases making up DNA and RNA. Thymine generally occurs only in DNA and uracilonly in RNA.

Biochemistry 235

neatly wound double helix. In DNA replica-tion, the two strands unwind and freenucleotides arrange themselves according tothe exposed amine bases in a complemen-tary fashion. DNA replication can be com-pared to an unzipping a zipper. As the DNAstrands unzips, complementary nucleotidesmatch up with the unpaired nucleotide. Inthis manner DNA replicates itself (Figure16.18).

Genes, residing in the chromosomes, aresegments of the DNA molecule. Thesequence of nucleotides, represented bytheir letters, corresponding to a specific genemay be hundreds or even thousands of let-ters long. Humans have between 50,000 and100,000 genes contained in their 46 chro-

mosomes, and the genetic code in humansconsists of roughly five billion base pairs.The process by which the information inDNA is used to synthesize proteins is calledtranscription. Transcription involves turningthe genetic information contained in DNAinto RNA. The process starts just like DNAreplication with the unraveling of a sectionof the two strands of DNA. A special proteinidentifies a promoter region on a singlestrand. The promoter region identifies wherethe transcription region begins. An enzymecalled RNA polymerase is critical in thetranscription process. This molecule initiatesthe unwinding of the DNA strands, producesa complementary strand of RNA, and thenterminates the process. After a copy of theDNA has been made, the two DNA strandsrewind into their standard double helixshape. The RNA strand produced by RNApolymerase follows the same process as inDNA replication except that uracil replacesthymine when an adenosine is encounteredon the DNA strand. Therefore, if a DNAsequence consisted of the nucleotides: C-G-T-A-A, the RNA sequence produced wouldbe G-C-A-U-U. The transcription processoccurs in the cell’s nucleus, and the RNAproduced is called messenger RNA ormRNA. Once formed, mRNA moves out ofthe nucleus into the cytoplasm where themRNA synthesizes proteins. The transfer of

Figure 16.17DNA Molecule (Rae Déjur)

Figure 16.18In DNA replication, the two strands makingthe double helix unravel and nucleotidespair against their complementary nucleotidein the original strands. In this manner, theDNA can replicate itself.

236 Biochemistry

genetic information to produce proteinsfrom mRNA is called translation. In thecytoplasm, the mRNA mixes with ribo-somes and encounters another type of RNAcalled transfer RNA or tRNA. Ribosomescontain tRNA and amino acids. The tRNAtranslates the mRNA into three lettersequences of nucleotides called codons.Each three letter sequence corresponds to aparticular amino acid. Because there are fournucleotides (C, G, A, and U), the number ofdifferent codons would be equal to 43 or 64.Because there are only twenty standardamino acids, several codons may producethe same amino acid. For example, thecodons GGU, GGC, GGA, and GGG allcode for glycine. Three of the codons serveas stop signs to signal the end of the gene.These stops also serve in some cases to ini-tiate the start of a gene sequence. Thesequential translation of mRNA by tRNAbuilds the amino acids into the approxi-mately 100,000 proteins in the human body.

BiotechnologySince Watson and Crick’s discovery of

the structure of DNA in 1953, advances inbiochemistry, molecular biology, and genet-ics have had tremendous impacts on mod-ern society. No area has been more affectedthan human health and modern medicine.Scientists continue to develop more precisemethods for recognizing genetic diseases,and gene therapies hold promise for curingsome of these diseases. In addition toadvances in human health, advances in bio-chemistry and molecular biology affectnumerous other fields. These include agri-culture with genetically engineered plantsand animals, forensics that uses DNA fin-gerprinting in crime analysis, and bioreme-diation that uses genetically modifiedmicroorganisms to clean contaminatedareas. While advances in biotechnology

have created new opportunities, they havealso presented society with difficult choicesconcerning the use of this technology. Con-troversial issues such as human cloning,stem cell research, and alternative forms ofreproduction such as artificial inseminationreceive wide coverage in the popular media.These issues will increasingly capture thepublic’s attention and provide a source ofdebate as further scientific discoveries aremade as the twenty-first century unfolds. Toconclude this chapter, we touch upon a fewmodern biochemical applications, empha-sizing the science behind the applications.

The Human Genome Project is an inter-national project to identify all the genes inthe nucleus of a human cell. In short, theHuman Genome Project seeks to read thebiochemical blueprint that is responsible formaking a human being. This blueprint,termed the human genome, consists of thesequence of base pairs on all 46 chromo-somes. By mapping the human genome, sci-entists hope to understand the relationshipbetween different human characteristics anddiseases and genes. For example, a genelocated on the seventh chromosome carriescystic fibrosis.

A technique that has revolutionized thefield of molecular biology during the lastdecade is the polymerase chain reaction(PCR). This method, conceived by Kary B.Mullis (1944–) in 1983, involves makingmultiple copies of fragments of DNA in ashort time. Mullis received the 1993 NobelPrize in chemistry for his work. PCR mim-ics DNA’s natural ability to replicate itself. Itinvolves three basic steps conducted at dif-ferent temperatures. In the first step, a mixture of DNA and other basic PCR ingre-dients is heated in a test tube to approxi-mately 90°C. At this temperature, the DNAstrands unwind. The second step involveslowering the temperature to around 55°C,which allows special enzymes called

Biochemistry 237

primers to mark the section of DNA to beduplicated. The primers serve as bookendsthat identify the section of DNA to becopied. Once the primers have attachedthemselves to the strand, the temperature israised to 75°C. At this temperature, thepolymerase makes a copy of the two DNAsegments. The entire process takes only afew minutes, and once completed the pro-cess is repeated to produce more copies. Inthis manner millions of copies of the origi-nal DNA material can be produced in a fewhours. One of the keys in perfecting thismethod was the use of polymerase isolatedfrom the bacterium Thermus aquaticus. Thisorganism resides in hot springs such asthose located in Yellowstone National Park.The polymerase used in the PCR is calledTaq polymerase to signify its origin. Theability of Taq polymerase to function overrapid temperature changes and at elevatedtemperatures made this polymerase ideal forthe PCR method. Most polymerases, such asthose extracted from humans, were not sta-ble at elevated temperatures.

The ability to produce multiple seg-ments of DNA segments is referred to asDNA amplification. The PCR method gavescientists a quick, simple, and inexpensivemeans to amplify DNA. Amplification isimportant because the amount of DNAavailable often is limited for a particularprocedure. Using PCR, the amount of DNAnecessary for certain analyses can be pro-duced from very small original samples.One application where this has been used isin genetic fingerprinting in crime cases. Asmall sample of DNA found in a drop ofblood or a piece of hair can be amplified,and then sequenced for comparison to asample of DNA obtained from a suspect orvictim. While most of the DNA sequence inall humans is identical, there are regions inthe DNA where individual variation exists.By focusing on these sections of DNA,

forensic experts now have the ability todetermine whether evidence gathered at acrime scene can be linked to a particularindividual. DNA fingerprinting can be usedto prove guilt or innocence; in recent yearsindividuals on death row have been freedwhen PCR and DNA fingerprinting wasapplied to evidence saved in police files.

PCR amplification has led to moresophisticated and accurate diagnostic tech-niques regarding diseases. This allows ear-lier detection of the disease compared toconventional methods, making earlier treat-ment possible. For example HIV (humanimmunodeficiency virus) may be detectedby searching for the DNA sequence uniqueto this virus. Amplifying samples andsearching for DNA associated with the bac-teria responsible for the condition has iden-tified infectious bacterial diseases. Lymedisease, certain stomach ulcers, and middleear infections have been detected in thismanner.

One interesting case of the use of PCRinvolved analysis of tissue preserved fromthe eyes of John Dalton. Dalton, who diedin 1844, requested that an autopsy be con-ducted after his death to determine the rea-son for his color blindness. The modernanalysis amplified DNA from Dalton’s eyetissue and discovered he lacked a gene nec-essary for detecting the color green. Thisexplains why Dalton suffered red-greencolor blindness and could not differentiatebetween these two colors.

Current PCR technology requires adevice about the size of a small televisionand costs several thousand dollars.Advances are reducing both the size andcost of PCR analyses. This is making PCRscreening standard practice in diagnosticmedicine and available in numerous otherareas. These include such diverse fields asarchaeology, ecology, art history, paleontol-ogy, and conservation biology.

238 Biochemistry

Genetic engineering involves the alter-ation of an organism’s genetic material inorder to introduce desirable effects or elim-inate undesirable effects. The most basicform of genetic engineering has been prac-ticed for thousands of years ever since ourancestors used selective breeding to culti-vate heartier plants and animals. While thepractice of selective breeding continues toflourish, genetic engineering in the twenty-first century involves the modern practice ofsynthesizing recombinant DNA. Recombi-nant DNA is DNA in which genetic mate-rial from one organism is spliced into thegenetic material of another. Using restric-tion enzymes to cut a section out of plasmid,DNA obtained from an organism, typicallya bacterium or yeast, is used in the initialstep to create recombinant DNA. Plasmidsare small circular DNA molecules that con-tain a few genes. The same restrictionenzyme is then used to cut a matching sec-

tion of chromosomal DNA (gene) fromanother organism. The exposed ends of theplasmid and chromosomal DNA containunpaired bases that are complementary toeach other. Using enzymes called DNA lig-ases, the ends of the DNA from the two dif-ferent organisms can be joined. Because theends of chromosomal DNA stick to the endsof the plasmid DNA, the ends are calledsticky ends (Figure 16.19). Once chromoso-mal DNA is inserted into the plasmid, thebacterium containing the recombined plas-mid is allowed to reproduce. The reproduc-tion of the bacteria produces clones. Clonesare cells with identical DNA that originatedfrom a single cell. In this manner multiplecopies of genes coded for specific proteinscan be made. The protein produced in thebacterium can then be extracted for humanuse.

Today thousands of products in med-icine, agriculture, and industry are geneti-

Figure 16.19Recombinant DNA (Rae Déjur)

Biochemistry 239

cally engineered using recombinant DNAtechnology. For example, plants have beenengineered to resist viruses, make themmore frost-resistant, and delay spoilage.Growth hormones are regularly used toincrease dairy and beef yield from cows. Inmedicine, recombinant DNA is used togenetically engineer proteins that can beused to treat various diseases. The technol-ogy is also used to produce vaccines thatare used to treat certain diseases, for exam-ple, hepatitis and influenza. Insulin was thefirst recombinant DNA substance approvedfor human use in 1982. Before this time,

diabetics depended on insulin extractedfrom hogs and cattle. Current methodsrequire the injection of insulin producedusing recombinant DNA. The hope for thefuture is that gene therapy can be used tocure diabetics. Gene therapy involvesinserting a healthy gene for an abnormal ormissing gene to cure a condition. While thistechnique is still being perfected, there ishope that gene therapy will eventually pro-vide a cure for many diseases. Currently,gene therapy is being attempted on a lim-ited basis to treat cystic fibrosis and mus-cular dystrophy.

IntroductionChemical reactions involve the interactionof the outer electrons of substances. As onesubstance changes into another, chemicalbonds are broken and created as atoms rear-range. Atomic nuclei are not directlyinvolved in chemical reactions, but they playa critical role in the behavior of matter. Typ-ical chemical reactions involve the interac-tion of electrons in atoms, but nuclearreactions involve the atom’s nucleus. Thenucleus contains most of the atom’s massbut occupies only a small fraction of its vol-ume. Electrons have only about 1/2000 themass of a nucleon. To put this in perspec-tive, consider that if the nucleus were thesize of a baseball, the mean distance to thenearest electrons would be over two miles.

At the start of the twentieth century, itwas discovered that the nucleus is composedof positively charged protons and neutralneutrons (see Chapter 4). These particles arecollectively called nucleons. During the lasthalf of the same century, scientists learnedhow to harness the power of the atom. Thedeployment of two atomic bombs brought aquick and dramatic end to World War II.This was followed by nuclear proliferation,the Cold War, and the current debate over

developing a nuclear defense shield. The useof nuclear power as a source of clean, effi-cient energy continues to be debated. Whilenuclear power is used by many countries tofulfill their energy needs, accidents such asChernobyl in Ukrainia and Three MileIsland in this country have dampened thepublic’s enthusiasm for this energy source.Because of nuclear weapons and reactoraccidents, people often perceive the term“nuclear” negatively, but there are manypositive aspects to nuclear science. Nuclearmedicine is used extensively for both diag-nostic and therapeutic procedures. Radio-metric dating is an invaluable tool toscientists who use this method for applica-tions such as dating relics, determining theage of the Earth, and studying climatechange. In this chapter, we extend our exam-ination of the nucleus started in Chapter 4,and examine in detail the scientific princi-ples behind human use of the atom.

Nuclear Stability andRadioactivity

Each element exists as several isotopes.Isotopes of an element have identicalatomic numbers, but different mass num-bers (see Chapter 4). Therefore, isotopes of

17

Nuclear Chemistry

242 Nuclear Chemistry

an element differ in the number of neutronsin their nuclei. Every element has at leastone unstable or radioactive isotope, butmost have several. The nuclei of unstableisotopes undergo radioactive decay.Radioactive decay is the process where par-ticles and energy are emitted from thenuclei of unstable isotopes as they becomestable.

Nuclear stability is related to the ratioof protons to neutrons. Protons packed intothe atom’s nucleus carry a positive chargeand exert a repulsive force on each other.For the nucleus to remain intact, a forcecalled the strong nuclear force must balancethe electrostatic repulsion between protons.The strong nuclear force is the “nuclearglue” responsible for holding the nucleustogether. This force is related to the ratio ofneutrons to protons in the nucleus. A nor-mal hydrogen atom’s nucleus contains a sin-gle proton, and therefore, no neutrons areneeded because there is no repulsive forcebetween protons. All other elements havemore than one proton and require neutronsto supply the strong nuclear force. Heliumatoms have two protons and two neutrons intheir nuclei. When moving up the periodictable, the atomic number increases as moreprotons exist in each element’s nucleus.When more protons are packed into thenucleus, more neutrons are required toovercome the resultant repulsive force. Forelements with an atomic number less than20, stable nuclei have either an equal num-ber of protons and neutrons or one moreneutron than proton. For example, carbonhas 6 protons and 6 neutrons and fluorinehas 9 protons and 10 neutrons. For elementswith atomic numbers greater than 20, theratio of neutrons to protons becomesincreasingly greater than 1. All isotopesabove bismuth, atomic number 83, areradioactive. The relationship betweennuclear stability and the ratio of neutrons to

protons is depicted in Figure 17.1. Figure17.1 shows a zone of stability. Stable nucleitend to be found in the zone of stability. Theneutron-to-proton ratio for isotopes outsidethe zone of stability characterizes unstablenuclei, and these undergo various forms ofradioactive decay.

Stable isotopes prefer a certain combi-nation of neutrons and protons. By far,most stable isotopes have an even numberof both protons and neutrons. A smallernumber of stable isotopes have either aneven number of protons and an odd num-ber of neutrons or vice-versa, and only afew have both an odd number of protonsand neutrons (Table 17.1). Stable nuclei arealso associated with a specific number ofprotons or neutrons. These islands of sta-bility occur when the number of protons orneutrons is 2, 8, 20, 28, 50, 82, and 126. Toillustrate this, consider tin with an atomicnumber of 50. Tin has ten stable isotopes,but antimony with an atomic number of 51has only two.

Figure 17.1The ratio of the number of neutrons to pro-tons determines the stability of atomicnuclei. As the number of protons in thenucleus increases, the number of neutronsmust increase at a greater rate to be stable.

Nuclear Chemistry 243

Radioactive Decay andNuclear Reactions

An unstable nucleus emits particles orelectromagnetic radiation, or both, until astable nucleus results. The main forms ofradioactive decay include alpha (�), beta(�), and gamma (�) decay. These threeforms of radiation are named for the firstthree letters of the Greek alphabet. An alphaparticle is equivalent to the nucleus of ahelium atom and consists of 2 protons and2 neutrons. A beta particle is equivalent toan electron. Gamma radiation is electro-magnetic energy with wavelengths thatrange from 10�14 to 10�11 m. When alphaand beta particles are emitted from anatom’s nucleus, the atomic number changes.This means the atom changes from one ele-ment to another. Alpha particles have a mass

number of 4, while beta particles have amass number of 0. Therefore, the emissionof an alpha particle changes both the atomicand mass number, but the emission of a betaparticle only changes the atomic number.Because gamma radiation is a form of light,the atomic number and mass number remainconstant when gamma radiation is emitted.The emission of gamma radiation does nottransform an element. Table 17.2 summa-rizes the properties and symbols used for thethree main types of radioactive emissions.

Conservation of mass and charge areused when writing nuclear reactions. Forexample, let’s consider what happens whenuranium-238 undergoes alpha decay. Ura-nium-238 has 92 protons and 146 neutronsand is symbolized as 238

92 U. After it emits analpha particle, the nucleus now has a massnumber of 234 and an atomic number of 90.

Table 17.1Nuclear Stability and Numbers of Neutrons and Protons

Table 17.2Main Forms of Radioactive Emissions


Because the atomic number of the nucleus is90, the element becomes thorium, Th. Theoverall nuclear reaction can be written as

Notice that in this equation mass is con-served, 238 � 234 �4, and charge is con-served, 92 � 90 � 2. In a chemical reaction,reactants are transformed into products. Ina nuclear reaction, the terms “parent” and“daughter” correspond to reactants andproducts, respectively. In our example, theU-238 was the parent that decayed into theTh-234 daughter.

During beta decay, a neutron is trans-formed into a proton. If Th-234 were to emita beta particle, it would be transformed intoprotactinium-234 according to the equation:

Again, both mass and charge are conserved.Gamma emission often accompanies bothalpha and beta decay, but because gammaemission does not change the parent ele-ment it is often emitted when writingnuclear reactions.

An unstable parent nucleus may decayinto either a stable or an unstable daughter.When the daughter is unstable, which isoften the case, the daughter will decay.Often the journey from an unstable nucleusto a stable nucleus involves a long series ofsteps referred to as a radioactive decayseries. One example is the decay series forradium (Figure 17.2).

Half-Life and RadiometricDating

In Chapter 12, the concept of half-lifewas used in connection with the time it tookfor reactants to change into products duringa chemical reaction. Radioactive decay fol-lows first order kinetics (Chapter 12). Firstorder kinetics means that the decay rate

depends on the amount of parent materialpresent at any given time. In nuclear decay,the number of parent nuclei decreases in anexponential fashion (Figure 17.3). Thedecay rate or activity of a radioactive sub-stance is equal to the decrease in parentnuclei over time

where �N equals the change in number ofparent nuclei and �T is the change in time.The standard unit for activity is the bec-querel. One becquerel is equal to 1 decayper second.

At any given time, a parent nucleus maydecay into a daughter. Although it cannot beknown when any individual nucleus maydecay into a daughter, the half-life can beused as a collective measure of the time it

238 234 4U r Th � He92 90 2

234 234 0Th r Pa � e90 91 �1

�Nactivity �

�T

Figure 17.2One possible decay series for radium. Ra-226 is transformed into stable Pb-206through a series of alpha and betaemissions. Horizontal arrows represent betaemissions and diagonal arrows alpha emis-sions.


takes for radioactive decay to take place.Half-life, often symbolized using t1/2 is thetime required for half the parent nuclei in asample to decay into daughter nuclei. Anequivalent definition is that it is the time ittakes for the activity of a substance to be cutin half. The concept of radioactive decayand half-life can be compared to poppingcorn. The kernels represent the parent nucleiand the popped corn the daughters. Duringthe popping process, some kernels popquickly after heat is applied, while othersnever pop. Although we cannot say whenany one kernel will pop, we could charac-terize the popping time using half-life. Forexample, the half-life of a bag of microwavepopcorn might be 2 minutes.

During each consecutive half-life, theamount of material remaining is one half ofthe amount present at the start of the half-life (Figure 17.4). The half-lives of radioac-tive isotopes vary over a wide range, from afraction of a second to over billions of years.Table 17.3 lists the half-lives of some com-mon isotopes.

The half-life of isotopes provides scien-tists with a nuclear clock that can be used todate objects. The concept is based on know-ing the fraction of original material that ispresent in a sample. For instance, if half ofthe original isotope is present in the sample,then the sample’s age is equivalent to theisotope’s half-life. If one-fourth of the orig-inal material is present, then the sample’sage is 2 half-lives. Because the use of radio-metric dating involves making accuratemeasurements of parent and daughter activ-ities in the sample, it is assumed that parent

Figure 17.3Exponential Decay (Rae Déjur)

Figure 17.4The amount of a radioactive substance is cutin half after each consecutive half-life. Theamount of the original material remainingafter n half-lives is (1⁄

2)n.

Table 17.3Half-Lives of Common Isotopes


material is not lost from the sample or thatthe sample has not been contaminated.

Radiometric techniques have beenapplied in a variety of disciplines includingarchaeology, paleontology, geology, clima-tology, and oceanography. The isotopicmethod used for dating depends on thenature of the object. Geologic samples thatare billions of years old would require usingan isotope with a relatively long half-lifesuch as potassium-40 (t1⁄2 � 1.25 � 109

years). Conversely, the movement of oceanwater masses or groundwater might use tri-tium (tritium is hydrogen containing 2 neu-trons, 3

1H) that has a half-life of 12.3 years.In general, an isotope can be used to date anobject up to ten times its half-life; for exam-ple, tritium could be used for measurementsof up to 125 years.

Carbon-14, t1⁄2 � 5,730 years, is one ofthe most prevalent methods used for datingancient artifacts. Williard Libby (1908–1980) developed the method around 1950and was awarded the Nobel Prize in chem-istry in 1960 for this work. The carbon-14method is based on the fact that livingorganisms continually absorb carbon intotheir tissues during metabolism. Most of thecarbon taken up by organisms is in its sta-ble carbon-12 form, but unstable carbon-14(as well as other carbon isotopes) is assim-ilated along with carbon-12. No distinctionis made between the isotopes of carbonassimilated, and so the ratio of carbon-14 tocarbon-12 in the organism remains thesame as that in the organism’s environment,for example, the atmosphere or ocean. Thenatural occurrence of C-14 in the environ-ment is about one C-14 atom for every 850billion C-12 atoms. As long as an organismis alive, the carbon-14 to carbon-12 ratioremains constant. Once the organism dies,the carbon-14 stopwatch starts, and the C-14:C-12 ratio begins to decrease due tothe decay of C-14. Comparing the activity

of C-14 in a sample to that in living mate-rial can date the sample. For example, if theactivity of C-14 in an excavated bone is halfthat of bone from a living organism, thenthe bone must be approximately 5,730 yearsold. Carbon-14 dating is used extensivelywith materials such as wood, seeds, fabrics,and bone.

A popular method used to date rocks isthe potassium-argon method. Potassium isabundant in rocks such as feldspars, horn-blendes, and micas. The K-Ar method hasbeen used to date the Earth and its geologicformations. It has also been applied to deter-mine magnetic reversals that have takenplace throughout the Earth’s history.Another method used in geologic dating isthe rubidium-strontium, Rb-Sr, method.Some of the oldest rocks on Earth have beendated with this method, providing evidencethat the Earth is approximately 5 billionyears old. The method has also been used todate moon rocks and meteorites.

Nuclear Binding Energy—Fission and Fusion

Because the nucleus consists of a col-lection of nucleons, it would be expectedthat the mass of the nucleus is equal to thesum of its constituent nucleons. In fact, themass of the nucleus is always slightly lessthan the sum of its parts. This decrease inmass of the nucleus is called the massdefect. What happens to this missing masswhen the nucleus of an atom is created outof nucleons? The answer is related to thestrong nuclear force that holds the nucleustogether. The mass defect is converted toenergy and released when the nucleonscombine to form a nucleus. The amount ofenergy released is related to the mass defectaccording to Einstein’s famous equation: E� mc2, where m is mass and c is the speedof light. Because the speed of light is so


high (c � 3 � 108 m/s or 6.7 � 108 mph), itdoes not take much mass to produce atremendous amount of energy. The energycalculated from Einstein’s equation usingthe mass defect for m is called the nuclearbinding energy. To separate the nucleus ofan atom into its individual protons and neu-trons would take an amount of energy equiv-alent to the binding energy.

The binding energy provides a measureof the stability of atomic nuclei, the greaterthe binding energy per nucleon of a nucleus,the greater its stability. When the amount ofbinding energy per nucleon is calculated forthe different elements, it is found that theFe-56 has the highest binding energy (Fig-ure 17.5). The fact that the highest bindingenergy curve peaks at the element iron hasimportant consequences for obtainingnuclear energy. When a heavier nucleussplits in a process called nuclear fission,lighter nuclei are produced. The lighternuclei are more stable. Whenever a processinvolves moving to a more stable state orconfiguration, energy is released. More sta-ble nuclei can also be obtained when lighternuclei combine to form a heavier nucleus ina process called nuclear fusion. The mostcommon example of a fusion reaction takesplace inside the sun, where hydrogen nucleifuse to form helium. The continual fusion of

solar hydrogen provides the energy thatmakes life on Earth possible.

Modern nuclear power is based on har-nessing the energy released in a fission reac-tion. The development of atomic energystarted in the 1930s with the discovery thatatoms could be split with neutrons. This dis-covery laid the foundation for building thefirst atomic bombs during World War II. Abasic reaction representing the fission ofuranium can be represented as:

The equation shows that uranium-235absorbs a neutron. After absorbing the neu-tron, the excited uranium nucleus splits andforms barium-141, krypton-92, and threeneutrons. Energy is also produced in thereaction. This reaction is only one of a num-ber of different ways that U-235 may split.Several hundred different isotopes have beenidentified when U-235 undergoes fission.

The three neutrons produced when ura-nium splits have the ability to split other U-235 nuclei and start a self-sustaining chainreaction. Whether a chain reaction takesplace depends on the amount of fissionablematerial present. The more fissionable mate-rial that is present, the greater the probabil-ity that a neutron will interact with anotherU-235 nucleus. The reason for this involvesthe basic relationship between surface areaand volume as mass increases. If a cubewith a length of 1 unit is compared to a cubeof 2 units, it is found that the surface area tovolume ratio of the 1 unit cube is twice thatof the 2 unit cube (Figure 17.6). This showsthat volume increases at a greater rate thansurface area as size increases. The probabil-ity that neutrons escape rather than reactalso depends on the surface area to volumeratio. The higher this ratio is the more likelyneutrons escape. When a U-235 nucleuscontained in a small mass of fissionable uranium is bombarded by a neutron, the

Figure 17.5Binding Energy Curve. Those elements suchas iron, cobalt, and nickel have the highestbinding energy per nucleon, and therefore,are the elements with the most stable nuclei.

1 235 141 92 1n � U r Ba � Kr � 3 n � energy0 92 56 36 0


probability is that less than one additionalfission will result from the three productneutrons. The amount of fissionable mate-rial at this stage is called a subcritical mass.As the mass of fissionable materialincreases, a point is reached where exactlyone product neutron causes another U-235fission, which in turn produces three moreneutrons, and one of these splits another U-235 nucleus. The quantity of fissionablematerial at the point where the reaction isself-sustaining is termed the critical mass.Once the critical mass is reached, a chainreaction can take place. If more than oneproduct neutron reacts with other U-235nuclei, a supercritical mass exists. In thiscase, the process accelerates, creating anuncontrolled chain reaction. The develop-ment of the first atomic bombs involvedbringing together simultaneously a numberof subcritical masses to form a supercriticalmass in a fraction of a second. Once thesupercritical mass was assembled, a streamof neutrons detonated the device.

Traditional nuclear power involvesusing the heat generated in a controlled fis-sion reaction to generate electricity. Aschematic of a nuclear reactor is shown inFigure 17.7. The reactor core consists of aheavy-walled reaction vessel severalmeters thick that contains fuel elementsconsisting of zirconium rods containingenriched pellets of U-235 in the form of

UO2. Natural uranium is 0.7% U-235 and99.3% U-238. Uranium-238 is nonfission-able, and therefore, naturally occurringuranium must be enriched to a concentra-tion of approximately 3% U-235 to be usedin common nuclear reactors. The reactor isfilled with water. The water serves two dif-ferent purposes. One purpose is to serve asa coolant to transport the heat generatedfrom the fission reaction to a heatexchanger. In the heat exchanger, theenergy can be used to generate steam toturn a turbine for the production of elec-tricity. Water also serves as a moderator.A moderator slows down the neutrons,increasing the chances that the neutronswill react with the uranium. Interspersedbetween the hundreds of fuel rods are con-trol rods. Control rods are made of cad-mium and boron. These substances absorbneutrons. Raising and lowering the controlrods controls the nuclear reaction.

While the basic design of most nuclearreactors is similar, several types of reactorsare used throughout the world. In the UnitedStates, most reactors use plain water as thecoolant. Reactors using ordinary water arecalled light water reactors (LWR). Lightwater reactors can be pressurized to approx-imately 150 atmospheres to keep the pri-mary coolant in the liquid phase attemperatures of around 300°C. The heatfrom the pressurized water is used to heatsecondary water to generate steam. In aboiling water reactor (BWR), water in thecore is allowed to boil. The steam producedpowers the turbines directly. Heavy waterreactors employ water in the form of D2O asthe coolant and moderator. Heavy water getsits name from the fact that the hydrogen inthe water molecule is the isotope deu-terium, D (2

1H), rather than ordinary hydro-gen. The use of heavy water allows naturaluranium rather than enriched uranium to beused as the fuel.

Figure 17.6As the size increases, volume increases at agreater rate than the surface area.


Breeder reactors were developed toutilize the 97% of natural uranium thatoccurs as nonfissionable U-238. The ideabehind a breeder reactor is to convert U-238into a fissionable fuel material, plutonium.A reaction to breed plutonium is

The plutonium fuel in a breeder reactorbehaves differently than uranium. Fast neu-trons are required to split plutonium. Forthis reason, water cannot be used in breederreactors because it moderates the neutrons.Liquid sodium is typically used in breederreactors, and the term liquid metal fastbreeder reactor (LMFBR) is used todescribe it. One of the controversies associ-ated with the breeder reactor is that it results

in the production of weapon grade pluto-nium and could contribute to nuclear armsproliferation.

The world use of nuclear power to sup-ply a nation’s electricity varies widely bycountry. France, for example, gets around75% of its electricity from nuclear power,and several other European countries getover half of their energy from this source.Approximately 20% of the electricity in theUnited States comes from 103 operatingnuclear power plants. Nuclear is second onlyto coal, 50%, and ahead of natural gas, 15%,hydropower, 8%, and oil, 3%, as a source ofelectrical energy. Although once hailed byPresident Eisenhower in the 1950s as a safe,clean, and economical source of power, theU.S. nuclear industry has fallen on hardtimes in the last twenty-five years. Nuclearaccidents at Three Mile Island, Pennsylvania,

Figure 17.7Schematic of Nuclear Power Plant (Rae Déjur)

238 1 239 239 0U � n r U r Np � e92 0 92 93 �1

239 0r Pu � e94 �1


this is related to the small size of hydrogennuclei. Many more hydrogen nuclei can bepacked into an H-bomb compared to thenumber of uranium nuclei in an atomicbomb. Because the energy produced in anuclear weapon depends on the number offission or fusion reactions that take place ina given volume, H-bombs are far moredestructive. The quest to produce the firstsuperbomb, as the H-bomb was known, ini-tiated the arms race that characterized therelationship between the United States andthe Soviet Union during most of the last halfof the twentieth century. It has only been inthe last decade that this race has abated withthe fall of the Soviet Union, but PresidentBush’s talk of a nuclear defense shield is areminder that the threat of mutually assureddestruction remains.

Temperatures in the range of 20 to 100million degrees Celsius are required forfusion reactions. For this reason, a hydrogenbomb is triggered by a conventional fissionatomic bomb. The atomic bomb producesthe tremendous heat necessary to fusehydrogen nuclei; therefore, fusion bombsare often referred to as thermonuclear. TheUnited States exploded the first hydrogenbomb on Eniwetok Atoll in the PacificOcean on November 1, 1952. This bombwas based on the fusion of deuterium:

The fusion of deuterium produces anotherhydrogen isotope called tritium, 3

1H, alongwith common hydrogen, 1

1H. The fusion ofdeuterium may also produce helium accord-ing to the reaction:

There is hope that nuclear fusion mayone day prove to be a practical energysource. The U.S. government has spent bil-lions of dollars on fusion energy research,and several prototype reactors have beenbuilt. Researchers attempting to produce

in 1979 and Chernobyl, Ukrainia, in 1986raised public concerns about the safety ofnuclear power. Utility companies, facingmore stringent federal controls and bur-geoning costs, have abandoned nuclearpower in favor of traditional fossil fuelplants. The last nuclear power plant orderedin the United States was built in 1978. Sev-eral nuclear plants under construction havebeen left abandoned, while other operatingplants have been shut down and decommis-sioned.

The lack of emphasis on nuclear powerduring the last twenty-five years does notnecessarily mean it will not stage a come-back in the future. Concerns about green-house gases and global warming from theburning of fossil fuels are one of thestrongest arguments advanced in favor of agreater use of nuclear power. New develop-ments in reactor design hold the promisethat safe reactors can be built. Yet, concernsabout the safe handling and transport ofnuclear fuels, nuclear waste disposal,nuclear proliferation, and a nimby (not inmy back yard) attitude present formidableobstacles to the resurrection of the industryin the United States.

The other type of nuclear reaction is afusion reaction. The binding energy curveindicates that fusion of light atomic nucleito form a heavier nucleus is an exothermicreaction. Although in theory many light ele-ments can combine to produce energy, mostpractical applications of fusion technologyfocus on the use of hydrogen isotopes. Thefirst application of fusion technologyinvolved development and testing of hydro-gen bombs (H-bombs) in the early 1950s.The United States and the former SovietUnion feared that the other would developthe H-bomb. This would pose an unaccept-able threat to the country without the bomb.Leaders of each country knew that hydrogenbombs would be much more destructivethan atomic (fission) bombs. The reason for

2 2 3 1H � H r H � H1 1 1 1

2 2 3 1H � H r He � N1 1 2 0


reactor. At a precise point, a pellet and thelaser beams all intersect to create a con-trolled thermonuclear explosion. The heatfrom the explosion can then be used to pro-duce work.

Fusion and Stellar EvolutionThe fusion of hydrogen to form helium

takes place throughout the universe. Starsare natural thermonuclear machines that areresponsible for the formation of all naturallyoccurring elements. The universe is pre-dominantly composed of hydrogen. Hydro-gen is not scattered uniformly throughoutthe universe. Galaxies are regions of highdensity, while interstellar space contains lit-tle. Some regions of space contain clouds ofcool hydrogen gas called nebula, and it is inthese regions that stars are born. The birthof a star begins when gravitational attractioncauses a nebula to contract. As it contracts,the mass of hydrogen generates heat. Thenebula continues to contract, and as it does,its temperature rises and its interior pressureincreases, forming a protostar. Eventually,a point is reached where the protostar’s inte-rior temperature reaches several milliondegrees. At this point, the temperature issufficient to initiate fusion of hydrogennuclei. Two primary hydrogen fusion reac-tions are:

In the first reaction a positively chargedelectron is formed. This particle is called apositron. In addition to the fusion of hydro-gen nuclei, the helium created can also enterinto fusion reactions:

Fusion reactions signify the formation of atrue star. The fusion reactions balance thegravitational contraction and an equilibrium

energy from fusion face several major prob-lems. One problem is how to produce thetremendous temperatures required forfusion. Another is how to confine a fuel,such as deuterium, in a manner to produce asustained reaction. Yet, another difficulty isobtaining more energy out of the reactionthan is required to initiate fusion. Onedesign used to produce temperatures ofaround 50,000,000°C involves using adonut-shaped coiled electromagnet called atokamak (Figure 17.8). The electromagneticfield within the cavity of the magnet acts onhydrogen isotopes, stripping them of theirelectrons and producing a plasma. Increas-ing the electric current intensifies the elec-tromagnetic field and raises the temperatureto a point where the deuterium nuclei canfuse. Confining the fuel in a suspendedplasma state within an electromagnetic fieldisolates the reactor materials from theextreme temperatures required for fusion. Inanother design called inertial confinement,a series of lasers is focused on glass pelletsfilled with a mixture of deuterium and tri-tium at a pressure of several hundred atmo-spheres. The pellets are then injected into a

Figure 17.8Tokamak Design (Rae Déjur)

1 1 2 0H � H r H � e1 1 1 1

1 2 3H � H r He1 1 2

3 1 4 0He � H r He � e2 1 2 1

3 3 4 1He � He r He � 2 H2 2 2 1


were created. This would imply that all theelements in our bodies were once part ofstars, and the heavier elements came from asupernova that exploded billions of yearsago.

TransmutationTransmutation is the process where

one element is artificially changed intoanother element. Rutherford conducted thefirst transmutation experiment in 1919 whenhe bombarded nitrogen atoms with alphaparticles. The nitrogen was transmuted intooxygen and hydrogen according to the reac-tion:

The process of transmutation produces mostof the known isotopes. In fact, only about10% of the approximately 3,000 known iso-topes occur naturally. The rest are synthe-sized in large instruments called particleaccelerators.

Particle accelerators are devices used toaccelerate charged particles and ions tospeeds approaching the speed of light. Bycolliding accelerated particles with otherparticles, scientists study the ultimate natureof matter. Collisions produced in particleaccelerators are like breaking open a piñata.Accelerated particles serve as a club tobreak open atomic nuclei and reveal theircontents. Hundreds of different elementaryparticles have been detected during colli-sion experiments. The term “elementary”was originally used to define a few basicindivisible particles such as the electron, butduring the last century hundreds of differentelementary particles have been discovered.The term “elementary” has been expandedto include a number of different classes ofparticles associated with matter. Protons,neutrons, and electrons are the most famil-iar elementary particles.

is reached. Hydrogen burning in the star’sinterior balances gravitational contraction aslong as a sufficient supply of hydrogen isavailable. Depending on the size of the star,hydrogen burning may last less than a millionyears or for billions of years. Larger starsexpend their hydrogen faster than smallerstars. The sun is considered a normal star,with an expected life expectancy of 10 billionyears. Its current estimated age is 5 billionyears, so it is about halfway through itshydrogen-burning stage.

Once hydrogen burning stops, there isno longer a balance between gravitationalcontraction and the nuclear energy releasedthrough fusion. The star’s interior at thisstage is primarily helium. The core willagain start to contract and increase in tem-perature. Eventually, a temperature isreached where helium fusion is sufficient tobalance the gravitational contraction.Depending on the size of the star, severalcontractions may take place. It should berealized that at each stage in a star’s lifecycle numerous fusion reactions take place.In addition to hydrogen and helium, manyother light elements such as lithium, carbon,nitrogen, and oxygen are involved.

Once a star has expended its supply ofenergy, it will contract to a glowing whiteember called a white dwarf. The elementsproduced in the interior of a star depend onthe size of the star. Small stars do not havesufficient mass to produce the temperaturesrequired to create the heaviest elements. Themost massive stars, though, may go througha series of rapid contractions in their finalstages. These massive stars have the abilityto generate the temperatures and pressuresnecessary to produce the heaviest elementssuch as thorium and uranium. The final fateof these massive stars is a cataclysmicexplosion called a supernova. It is in thismanner that scientists believe all the natu-rally occurring elements in the universe

14 4 17 1N � He r O � H7 2 8 1


mass of the particle due to relativisticeffects.

Particle accelerators are some of thelargest scientific instruments. The largestparticle accelerator is located at CERN(European Organization for NuclearResearch) outside of Geneva on the borderof France and Switzerland. It has a circum-ference of 25 kilometers and is buried 100meters below the surface of the Earth. Thisaccelerator is known as LEP for large electron-positron collider. Fermi Labs out-side of Chicago contains an acceleratorcalled a Tevatron. The Tevatron has a four-mile circumference and is used to acceler-ate protons and antiprotons in oppositedirections to 99.9999% of the speed of light.These particles are then smashed togetherand detectors connected to computers areused to analyze the results. The bottom andtop quark (Chapter 4) were first detected atFermi Labs. The charm quark was discov-ered at the Stanford Linear Accelerator Cen-ter. To advance the frontiers of science andlearn more about the nature of matterrequires building more powerful and largerparticle accelerators. The United Statesstarted construction on what would havebeen the largest particle accelerator in 1983.Two billion dollars was spent on the designand initial construction of the Supercon-ducting Super Collider (SSC). This wouldhave been a giant accelerator built in Texasusing thousands of superconducting mag-nets. The SSC would have accelerated pro-tons and antiprotons around an oval track 54miles in circumference. The proton energieswould be twenty times that produced by cur-rent accelerators. In 1993, Congress can-celled funding of this project.

Particle accelerators are used to pro-duce most isotopes by transmutation. Allelements greater than uranium, known asthe transuranium elements, have been pro-duced in particle accelerators. For example,

Several basic types of particle acceler-ators are used, but they are all based on theprinciple that charged particles can beaccelerated in an electromagnetic field. Thefirst accelerators built in the 1930s were lin-ear accelerators, or linacs. These consistedof a series of hollow tubes in which thealternating current regulated down thelength of the tube was used to accelerate theparticle (Figure 17.9). The speed a particlecan obtain using a linear accelerator isdirectly related to the length of the acceler-ator. The largest linear accelerator is locatedat Stanford University and is slightly longerthan two miles.

Another design for particle acceleratorsis based on a circular arrangement. Acyclotron is similar to a linear acceleratorwound into a spiral. A series of electromag-nets causes the particles to move in a circleas they are accelerated by the electric field.According to Einstein’s theory of relativity,an object’s mass increases as it accelerates.In particle accelerators this is a problembecause as the mass increases, the particleslows down and becomes out of sync withthe changing electric field. A synchrotron isa cyclotron in which the electric fieldincreases to compensate for the change in

Figure 17.9In a linear accelerator, a positive proton isaccelerated through a series of tubes byusing alternating current. The proton ispushed by the positive charge on the firsttube and pulled by the negative charge onthe second tube, causing it to accelerate. Bytiming the change in voltage applied to thetubes, the proton continues to acceleratedown the length of the tube.


tons are negatively charged protons, andantineutrons have different magnetic proper-ties than regular neutrons. Antiparticles havebeen detected in particle accelerators, andantiatoms were detected at CERN in 1995.When matter meets antimatter, mutual anni-hilation results with the mass converted intoshort-lived particles and gamma rays. Thisprinciple forms the basis of PET.

In PET, positron-emitting isotopes suchas C-11, N-13, O-15, and F-18 are used.Each of these unstable isotopes is charac-terized by lacking a neutron compared to itsstable form; for example, C-11 needs onemore neutron to become C-12. Theyundergo positron emission when a protonchanges into a neutron:

When a PET isotope is administered totissue, it emits positrons. These positronsencounter electrons found in tissue, theinteraction of an electron and positronresults in mutual annihilation and the pro-duction of beams of gamma rays directlyopposite to each other (Figure 17.10). The

Pu-241 is produced when U-238 collideswith an alpha particle:

Tritium is produced in a nuclear reactionbetween lithium and a neutron:

Nuclear MedicineWhile our discussion on transmutation,

particle accelerators, and elementary parti-cles has focused on basic research, nucleardiscoveries have been applied in manyareas. Radioactive isotopes produced in par-ticle accelerators have found wide use innuclear medicine. Radioactive nuclei areused to diagnose and treat a wide range ofconditions. Radioactive tracers are iso-topes that can be administered to a patientand followed through the body to diagnosecertain conditions. Most tracers are gammaemitters with short half-lives (on the orderof hours to days) that demonstrate similarbiochemical behavior as their correspondingstable isotopes. Because tracers are radioac-tive, radiation detectors can be used to studytheir behavior in the body. Several examplesof common tracer studies in medicineinclude iodine-131 to study the thyroidgland, technetium-99 to identify braintumors, gadolinium-153 to diagnose osteo-porosis, and sodium-24 to study blood cir-culation.

A recent development in nuclearmedicine that illustrates how advances inbasic research are transformed into practicalapplications is positron emission tomogra-phy or PET. PET creates a three-dimensionalimage of a body part using positron emittingisotopes. Positrons, positively charged elec-trons, are a form of antimatter. Antimatterconsists of particles that have the same massas ordinary matter, but differ in charge orsome other property. For example, antipro-

238 4 241 1U � He r Pu � n92 2 94 0

6 1 3 4Li � n r H � He3 0 1 2

1 1 0H r n � e1 0 1

Figure 17.10In positron emission tomography, positronsinteract with electrons in body tissue to pro-duce gamma rays that are detected and con-verted into an image.


When matter absorbs radiation, a cer-tain quantity of energy is imparted to thematter. The dose is the amount of radiationenergy absorbed by matter. With respect tohuman health, the matter of concern ishuman tissue. The traditional unit of dose isthe rad. One rad is defined as the 100 ergsof energy absorbed per gram of matter. TheSl unit for dose is the gray, Gy. One gray isequal to 1.0 joule of energy absorbed perkilogram of matter. The relationshipbetween rads and grays is 1 gray is equal to100 rads. Similar to activity measurements,prefixes can be attached to the gray unit.

The most important measure of radia-tion in terms of health effects is the doseequivalent. Dose by itself does not considerthe type of radiation absorbed by tissue.Radiation effects on human tissue dependon the type of radiation. For instance, a doseof alpha radiation is about twenty timesmore harmful compared to equivalent dosesof beta or gamma radiation. Dose equivalentcombines the dose with the type of radiationto give dose equivalent. It is found by mul-tiplying the dose by a quality factor. Thequality factors for gamma, beta, and alpharadiations are 1, 1, and 20, respectively.Dose equivalent is the best measure of theradiation effects on human health. The remis the dose equivalent when dose is mea-sured in rads. Rem is the most frequentlyused unit for measuring dose equivalent.The unit rem is actually an acronym forroentgen equivalent for man. Millirems(mrems or 1/1,000 rem) are most often usedto express everyday exposures (Table 17.4).When expressing dose in grays, the unit fordose equivalent is the sievert, Sv. One remis equal to 0.01 Sv. The various units asso-ciated with radiation measurements aresummarized in Table 17.4.

Radiation comes in many differentforms. Most of this chapter has dealt withnuclear radiation, those forms of radiation

beams strike gamma detectors, and the twosignals are translated into an image by acomputer. PET has been used to diagnoseneurological disorders such as Parkinson’sdisease, Alzheimer’s disease, epilepsy, andschizophrenia. PET is also used to diagnosecertain heart conditions and examine glu-cose metabolism in the body.

Radioisotopes are also used in radiationtherapy to treat cancer. The goal in radiationtherapy is to kill malignant cells, while pro-tecting healthy tissue from radiation effects.Radioisotopes such as yttrium-90, a betaemitter, may be placed directly in the tumor.Alternatively, the diseased tissue may besubjected to beams of gamma radiation.Cobalt-60 used in radiation therapy is pre-pared by a series of transmutations:

Radiation Units andDetection

Radiation measurements are expressedin several different units depending on whatis being measured. The strength or activityof a radioactive source is the number of dis-integrations that occurs per unit time in asample of radioactive material. This is mostoften expressed as the number of disinte-grations per second. The Sl unit for activityis the becquerel, abbreviated Bq. One bec-querel is equivalent to one disintegration persecond. The curie, Ci, is often used whenthe activity is high. One curie is equal to 3.7� 1010 becquerels. Common prefixes suchas pico (10�12), micro (10�6), milli (10�3)are often used in conjunction with curies toexpress activity values. For example, EPAconsiders radon levels in a home high if themeasured activity in the home is greaterthan 4.0 picocuries per liter of air.

58 1 59Fe � n r Fe26 0 26

59 59 0Fe r Co � e26 27 �1

59 1 60Co � n r Co27 0 27


that come from the nuclei of atoms. Nuclearradiation originates from the Earth. Whennuclear radiation comes from an extrater-restrial source, it is referred to as cosmicradiation. Nuclear and cosmic radiationconsist of particles such as protons, elec-trons, alpha particles, and neutrons, as wellas photons such as gamma rays and x-rays.When nuclear and cosmic radiation strikematter and free ions, the term ionizing radi-ation is used. Alpha, beta, and gamma radi-ation are forms of ionizing radiation.Nonionizing radiation includes ultravioletradiation, responsible for sunburns, andinfrared radiation. Humans sense infraredradiation as heat.

Radiation is impossible to detect usingour unaided senses. It cannot be felt, heard,smelled, or seen. The fact that radiation ion-izes matter as it passes through it provides abasis for measuring radiation. A commondevice used to measure radiation is theGeiger counter. A Geiger counter consists ofa sealed metal tube filled with an inert gas,typically argon. A wire, protruding into the

center of the tube, connected to the positiveterminal of a power source serves as oneelectrode. The negative terminal is con-nected to the wall of the tube and serves asthe negative electrode (Figure 17.11). Thevoltage difference between the wire andtube wall is roughly 1,000 volts. No currentflows between the wire and tube wallbecause the inert gas is a poor conductor ofelectricity. One end of the sealed tube has a

Table 17.4Common Radiation Units

Figure 17.11A Geiger Counter detects radiation by ioniz-ing argon atoms. The flow of argon ions andelectrons in the tube creates a current that isamplified and detected by the counter.


thin window that allows radiation to enter.When radiation enters the tube, it ionizes theatoms of the inert gas. The free electronsproduced during ionization move toward thepositive electrode, and as they move, theyionize other gas atoms. The positivelycharged argon ions move toward the nega-tive walls. The flow of electrons and ionsproduces a small current in the tube. Thiscurrent is amplified and measured. There isa direct relationship between the currentproduced and the amount of radiation thatenters the window so activity can be mea-sured.

Scintillation counters use materials thatproduce light when stimulated by radiation.Scintillation materials include sodiumiodide crystals and special plastics. Radia-tion is measured by exposing the scintilla-tion material to radiation and using aphotomultiplier tube to count the number ofresulting flashes. Photomultiplier tubes arephotocells that convert light into electricalsignals that can be amplified and measured.

Biological Effects ofRadiation

Radiation is ubiquitous in the environ-ment. Life exists in a sea of nuclear, cosmic,and electromagnetic radiation. Our bodiesare constantly exposed to radiation and theeffect of radiation on cells has severaleffects. Radiation may 1) pass through thecell and not do any damage; 2) damage thecell, but the cell repairs itself; 3) damage the cell and the damaged cell does not repairitself and reproduces in the damaged state,or 4) kill the cell. Fortunately, most radia-tion does not cause cellular damage, andwhen it does, DNA tends to repair itself.Under certain circumstances the amount ofradiation damage inflicted on tissue over-whelms the body’s ability to repair itself.Additionally, certain conditions disrupt

DNA’s ability to repair itself. Figure 17.12shows several different ways DNA may bedamaged by ionizing radiation.

Low levels of ionizing radiation gener-ally do not affect human health, but healtheffects depend on a number of factors.Among these are the types of radiation, radi-ation energy absorbed by the tissue, amountof exposure, type of tissue exposed to radia-tion, and genetic factors. Alpha, beta, andgamma radiation display different character-istics when they interact with matter. Alphaparticles are large and slow with very littlepenetrating power. A sheet of paper or theouter layer of skin will stop alpha particles.Their energy is deposited over short dis-tances, resulting in a relatively large amountof damage. Beta particles are smaller thanalpha particles and have a hundred times

Figure 17.12Radiation may damage DNA in several ways.It may cause breaks in the sugar-phosphatebackbone and the base crosslinks. It mayalso result in incorrect repair. (Rae Déjur)


their penetrating power. Beta particles can bestopped by a few millimeters of aluminumor a centimeter or two of tissue. Their energyis deposited over a distance, and hence is lessdamaging than alpha particles. Gamma par-ticles have a thousand times the penetratingpower of alpha particles. Several centimetersof lead are needed to stop gamma radiation.Gamma particles travel significant distancesthrough human tissue, and therefore, deposittheir energy over a relatively wide area.

It is important to consider the type ofradiation with respect to whether a radiationsource is external or internal to the body.External alpha, and to some extent beta radi-ation, are not as hazardous as externalgamma radiation. Clothing or the outer layerof the skin will stop alpha radiation. Beta

radiation requires heavy protective clothing,because it can penetrate the skin and burntissue. The body cannot be protected fromexternal gamma radiation using protectiveclothing. While external alpha particles arerelatively harmless, they cause significantdamage internally. When inhaled andingested, alpha emitters may be depositeddirectly into tissue, where they can damageDNA directly.

As long as radiation dose equivalentexposures are low, radiation damage is non-detectable. General effects of short-termradiation exposure are summarized in Table17.5. The dose equivalents in Table 17.5 arelisted in rems. These values are several ordersof magnitude greater than what humansreceived in a year. Annual human exposure

Table 17.5Biological Effects of Short-Term, Whole Body Radiation Exposure


is normally around several hundred mrems.Human exposure to radiation occurs thatfrom both natural and human sources. Natu-ral radiation is referred to as backgroundradiation. It consists mainly of radon thatemanates from radioactive elements in theEarth’s crust and cosmic rays (Figure 17.13).The main human source of radiation expo-sure is medical x-rays. Humans carry a cer-tain amount of radioactive material insidetheir bodies. Potassium, as the isotope K-40,is responsible for nearly half of the exposurehumans receive from consumed food. Othersources of human exposure include C-14 andradium, which acts like calcium in the body.Radioactive isotopes are incorporated intobone and other tissue along with stable isotopes.

Radiation exposure from both naturaland human sources varies widely. Back-ground radiation depends on the local geol-ogy and elevation. Areas where radioactiverocks are located close to the surface orwhere mining has exposed mineral depositshave higher background levels. Higher

background levels also exist at higher ele-vations because there is less atmosphere toscreen out cosmic radiation. Patients beingtreated with radiation therapy may haveexposures of several hundred thousandmrems. Exposures of this magnitude involveradiation focused on the diseased tissue,rather than whole-body exposure.

SummaryChemical change involves the interac-

tion of the valence electrons in atoms asbonds are broken and formed. Even thoughchemical change creates new substances,atoms of elements are conserved. Nuclearchanges result in the transformation of oneelement into another. These changes alsoinvolve a number of elementary particles andenergy. Although a sea of radiation sur-rounds humans, the word “radiation” is oftenassociated with negative images. Humanshave used natural and synthesized radioac-tive isotopes for national defense, medicine,basic research, and energy production, and

Figure 17.13Annual human exposure to radiation in millirems (mrems). Note that although radon is radia-tion from the Earth, it is displayed as a separate category.


in numerous commercial products. Forexample, smoke detectors contain a sourceof alpha-emitting americium and gas lanternmantles contain radioactive thorium. Lumi-nescent exit signs found in many schools,movie theaters, and airlines contain tubesfilled with tritium. As with any technology,the benefits of nuclear applications must beweighed against their costs. This is particu-larly true with radioactive materials. As theuse of these materials continues, scientistswill seek to control their deleterious effects.

In the future, nuclear power may betaken for granted, and our generation’sburning of fossil fuels may be viewed as

primitive. For now, the use of nuclear tech-nology is in its infancy. In Star Trek, thestarship Enterprise relies on two forms ofnuclear power to power it through the uni-verse. Fusion power is used for the impulseengines to achieve speeds below the speedof light. Warp speeds, which are faster thanthe speed of light, are obtained by usingenergy obtained from a reaction betweenmatter and antimatter. Protons and antipro-tons are used to power the Enterprise’swarp drive. While today fusion power useas a practical energy source is science fic-tion, it may one day be a viable source ofenergy.

IntroductionThroughout human history, homo sapienshave had an impact on our planet. As theplanet’s human population has grown,accompanied by advances in technology,human impacts have also increased. In thelast century, the world population has grownfrom 3 billion in the year 1900 to over 6 bil-lion today. Each day 2 million more peopleare added to the planet. The industrial revo-lution and advances in science throughoutthe twentieth century have extended the lifespans of humans. The Earth’s increasingpopulation and use of its resources to sup-port this population have natural conse-quences for the environment. Humanactivities continually affect the Earth’s litho-sphere (land), hydrosphere (water), atmo-sphere (air), and biosphere (organisms).The Earth’s natural resources supply thechemicals used by humans for the energyand the raw materials used to sustain life.Throughout history, most of the substancesused by humans have consisted of naturalproducts, but in the last fifty years chemistshave learned how to synthesize a plethora ofnew products. While synthetic substanceshave provided many benefits to society, they

have also created additional impacts on theplanet.

Environmental chemistry deals with thechemical aspects of the interaction betweenhumans and their environment. Several ofthe issues of concern to environmentalchemists are global in nature and have beenwell publicized. These include stratosphericozone destruction, acid precipitation, andthe greenhouse effect. Although these issuesgarner the bulk of the popular media’s atten-tion, many common applications of envi-ronmental chemistry have a far greaterimpact on everyday life. For example, dis-infecting drinking water with chlorine haseliminated many diseases transmittedthrough water in many regions. Wastewateris also subjected to chlorination and mayundergo further chemical treatments beforebeing discharged into receiving waters.Whenever a vehicle is started, we immedi-ately initiate a number of chemical reactionsthat transform fossil fuels into other chemi-cals that are eventually released into theatmosphere. Modern agriculture depends ona variety of synthetic chemicals used as fer-tilizers, pesticides, and herbicides. Agricul-tural chemicals have an impact on soil,

18

Environmental Chemistry

262 Environmental Chemistry

groundwater, surface water, and the atmo-sphere. Use of synthetic chemicals is preva-lent in many industries. Most of these arepresumed safe, although some have beendiscovered to be harmful after theirwidespread use. Examples include PCBs inelectrical transformers, dioxins in the paperindustry, and CFCs used in refrigerants andpropellants.

In this chapter, we examine the envi-ronmental chemistry associated with bothglobal and local issues. One goal of thischapter is to provide a balanced perspectiveon the subject. This in turn should help thereader use chemistry to make informeddecisions about environmental issues. Anaxiom of the environmental movement is to“think globally and act locally.” While bothglobal and local issues will be treated sepa-rately, the nature of having an environmen-tal perspective means knowing that localdecisions made by many individuals cumu-latively have global consequences.

Stratospheric Ozone Depletion

Ozone, O3, is a relatively unstableallotrope of oxygen consisting of three oxy-gen atoms. Although it exists in only minuteconcentrations in the atmosphere, ozone’spresence is critical for life on Earth. Ozoneacts as a protective shield by absorbingharmful ultraviolet radiation in the strato-sphere. The stratosphere is the region of theatmosphere lying between approximately 13and 20 kilometers (8 to 12 miles) above theEarth’s surface. Ozone is constantly createdand destroyed in the stratosphere. Thisoccurs when ultraviolet light with sufficientenergy splits the double bond of the ordi-nary molecular oxygen, O2, producing twooxygen atoms. The reaction can be repre-sented as

O2 � UV 2O

Ultraviolet light, like other forms ofelectromagnetic radiation, can be character-ized by its wavelength and frequency. Wave-length and frequency are inversely related,as wavelength gets larger the frequency getssmaller. Figure 18.1 shows the electromag-netic spectrum for common forms of radia-tion.

The energy of a photon of electromag-netic radiation is given by Planck’s equation:

In this equation, E is energy of a photon injoules, h is Planck constant � 6.63 � 10�34

J-s, c is the speed of light � 3.0 � 108 m/s,and � is the wavelength of electromagneticradiation in meters. Planck’s equation showsthat the energy of electromagnetic radiationis inversely proportional to its wavelength.The maximum wavelength of light neces-sary to split a molecule of oxygen is 242nm. Ultraviolet light includes radiationbetween about 50 and 400 nm.

Oxygen atoms produced by the ultravi-olet dissociation of O2 can take part in sev-eral reactions that create and destroy ozone.One set of reactions involving atomic andmolecular oxygen is known as the Chap-man reactions. The reactions are namedafter Sydney Chapman (1888–1970) whofirst proposed them in 1930. In one reaction,an oxygen atom combines with O2 to formozone. Alternately, it can recombine withanother oxygen atom to produce an oxygenmolecule or react with ozone to produce twooxygen molecules. These reactions are sum-marized as

O � O2 O3

O � O O2

O3 � O 2O2

Ozone formed in the stratosphere isunstable. It can be broken down by ultravi-olet light shorter than 320 nm (UV-C andUV-B) according to the reaction:

hcE �

�

Environmental Chemistry 263

O3 � UV O2 � O*

where O* represents an excited oxygenatom. The form of oxygen, which exists inthe stratosphere, is related to the altitude.The top of the stratosphere is highly ener-getic with an abundance of UV-C radiation.In this region oxygen exists predominantlyas atomic oxygen. Collisions between oxy-gen atoms can form molecular oxygen, butthe presence of UV-C radiation keepsmolecular oxygen concentrations low. Thisin turn precludes any appreciable accumu-lation of ozone. In the lower stratosphere,ozone concentrations are low due to theabsence of UV-C radiation. Because UV-Chas been absorbed in the upper and middlestratosphere by O2 and O3, it is not availableto produce atomic oxygen. Because atomicoxygen is not present in any appreciableamount in the lower stratosphere, ozone isnot formed here. The presence of adequateconcentrations of both atomic and molecu-lar oxygen in the middle stratosphere meansthat stratospheric ozone concentrations peak

at about 35 kilometers altitude. Even thoughozone reaches its highest concentrations atthis level, its concentration is still less than10 ppm. Thus, the term “ozone layer” ismisleading. Ozone does not exist as a pro-tective blanket, but more as a diffuse screenthat partially protects the Earth from harm-ful ultraviolet radiation.

The Chapman reactions show thatozone is continually being created anddestroyed in the stratosphere. Natural pro-cesses such as volcanoes, climate change,and solar activity change ozone concentra-tions over long time periods. Seasonaleffects in ozone concentrations also existdue to changes in solar radiation. Ozonealso reacts with many chemicals in theatmosphere such as nitrogen, chlorine,bromine, and hydrogen. In this process, asubstance captures an oxygen from ozone.If X symbolizes a chemical substance thatremoves oxygen, then a general reaction forthe destruction of O3 is represented as:

X � O3 XO � O2

Figure 18.1The E-M Spectrum (Rae Déjur)


The XO compound formed in this reactionthen reacts with oxygen to regenerate X andmolecular oxygen:

XO � O X � O2

The regenerated S can destroy another O3

molecule, and the process can continueindefinitely. Thus a single atom or moleculecan destroy many ozone molecules.

Even though stratospheric ozone con-centrations are the result of dynamic pro-cesses, a natural balance of creation anddestruction can be expected to produce asteady-state concentration. Steady-stateoccurs when the concentration of ozoneremains constant and exists when the rate ofozone formation equals the rate of ozonedestruction. Steady-state is similar to lettingwater run into a bathtub while it drains out,but the water level in the tub remains thesame. Humans have disrupted this equilib-rium in recent decades by introducing chem-icals into the atmosphere that have increasedthe rate of destruction. Reduction of ozonelevels varies with location, but during thelast three decades, there has been about a5% decrease worldwide. Reduction in ozonemeans that a greater percentage of UV-Bradiation can penetrate the atmosphere. UV-B is ultraviolet radiation with wavelengthsbetween 280 and 320 nm, UV-A includeswavelengths between 320 and 400 nm, andUV-C wavelengths are between 220 and 280nm. Higher UV-B levels have been associ-ated with increased rates of skin cancer andcataracts. Another consequence of increasedUV-B radiation is its ability to suppress theimmune system in organisms. UV-B mayaffect entire ecosystems. Increased UV-Bradiation has been mentioned as a con-tributing factor in the destruction of coralreefs throughout the tropics; it may alsodestroy or disrupt phytoplankton and reduceproductivity in the oceans.

Although ozone depletion is global, thepolar regions, especially Antarctica, haveexperienced the greatest impacts and havereceived the most attention with respect tothis problem. The “ozone hole” appearsabove Antarctica around September or Octo-ber. It is called an ozone hole because at thistime of the year ozone concentrations dropprecipitously by as much as 95%. The rea-son for this condition is a combination ofmeteorological conditions and chemistry.During the Southern Hemisphere’s winter, aweather condition known as the polar vortexis established over Antarctica. The polar vor-tex is a wind pattern created as cool airdescends over the continent. The descendingair is deflected by the Coriolis force into awesterly direction around Antarctica. Thevortex effectively isolates the atmosphereover Antarctica and creates the region whereozone depletion occurs. The strength andduration of the polar vortex varies annually,but in general it can be thought of as aboundary that defines a giant reaction vesselwhere ozone-destroying reactions take place.

During the dark, polar winter the tem-perature drops to extremely low values, onthe order of –80°C. At these temperatures,water and nitric acid form polar strato-spheric clouds. Polar stratospheric cloudsare important because chemical reactions inthe stratosphere are catalyzed on the surfaceof the crystals forming these clouds. Thechemical primarily responsible for ozonedepletion is chlorine. Most of the chlorinein the stratosphere is contained in the com-pounds hydrogen chloride, HCl, or chlorinenitrate, ClONO2. Hydrogen chloride andchlorine nitrate undergo a number of reac-tions on the surface of the crystals of polarstratospheric clouds. Two important reac-tions are:

ClONO2(g) � HCl(s) Cl2(g) � HNO3(s)

ClONO2(g) � H2O(s) HOCl(g) � HNO3(s)


During the long Antarctic night, appreciableamounts of molecular chlorine, Cl2, andhypochlorous acid, HOCl, accumulatewithin the polar vortex. When the sunreturns during the spring (in September inAntarctica), ultraviolet radiation decom-poses the accumulated molecular chlorineand hypochlorous acid to produce atomicchlorine, Cl. Atomic chlorine is a free rad-ical. Free radicals are atoms or moleculesthat contain an unpaired or free electron.The Lewis structures of free radicals containan odd number of electrons. The unpairedelectron in free radicals makes them veryreactive. The free radical Cl produced fromthe decomposition of Cl2 and HOCl cat-alyzes the destruction of ozone as repre-sented by the reaction:

Cl � O3 ClO � O2

ClO � O Cl � O2

net reaction O3 � O 2O2

This reaction shows that the original chlo-rine atom acts like “X” in the general ozone-destroying reaction presented previously. Achlorine atom captures an oxygen atomfrom ozone in the first reaction. A chlorineatom is regenerated in the second reaction.Thus, a single chlorine atom can lead to thedestruction of numerous ozone molecules.

There are numerous natural contributorsof chlorine to the stratosphere, for example,volcanic eruptions. The main concernregarding ozone destruction in recent yearsis associated with human activities that haveincreased chlorine and other syntheticchemical input into the stratosphere. At thetop of the list of such chemicals are chlo-rofluorocarbons, or CFCs. CFCs are com-pounds that contain carbon, chlorine, andfluorine; they were first developed in 1928.Common CFCs are called Freons, a tradename coined by the DuPont chemical com-pany. CFC compounds are nonreactive, non-toxic, inflammable gases. Because of their

relative safety, CFCs found widespread useas refrigerants, aerosol propellants, solvents,cleaning agents, and foam blowing agents.Their use grew progressively from the1930s until the 1990s, when their annualproduction was over a billion pounds.

Chlorofluorocarbons are named using acode that can be used to determine theCFC’s formula. Adding 90 to the CFC num-ber gives a three-digit number that indicatesthe number of carbon, hydrogen, and fluo-rine atoms in the molecule. Any remainingatoms in the molecule are assumed to bechlorine. For example, the formula for CFC-12 is found by adding 90 to 12 to give 102.Therefore, the CFC-12 molecule has onecarbon atom, no hydrogen atoms, and twofluorine atoms. Because carbon is bonded tofour other atoms, the remaining two atomsin the molecule are chlorines. The chemicalformula for CFC-12 is CF2Cl2.

F. Sherwood Rowland (1927–) andMario Molina (1943–) predicted thedestruction of stratospheric ozone in 1974.Rowland and Molina theorized that inertCFCs could drift into the stratosphere,where they would be broken down by ultra-violet radiation. Once in the stratosphere,the CFCs would become a source of ozone-depleting chlorine. The destruction of ozoneby CFCs can be represented by the follow-ing series of reactions:

CF2Cl2 � UV Cl � CF2Cl

Cl � O3 ClO � O2

ClO � O Cl � O2

net reaction CF2Cl2 � O3 � O Cl

� CF2Cl � 2O2

This set of reactions shows that ultravioletradiation strikes a CFC molecule removinga chlorine atom. The chlorine atom collideswith an ozone molecule and bonds with oneof ozone’s oxygen atoms. The result is theformation of chlorine monoxide, ClO, andmolecular oxygen. Chlorine monoxide is a


free radical and reacts with atomic oxygen,producing another atom of chlorine that isreleased and available to destroy moreozone.

The actual destruction of ozone in thestratosphere actually involves hundreds ofdifferent reactions. Besides the Chapmanreactions and destruction by CFCs, manyother chemical species can destroy ozone. In1970, Paul Crutzen (1933–) showed thatnitrogen oxides could destroy ozone. Nitricoxide can remove an oxygen atom fromozone and be regenerated according to thefollowing reactions:

NO � O3 NO2 � O2

NO2 � O NO � O2

Crutzen’s work provided arguments againstthe development of high-flying super sonicplanes, such as the Concorde, that woulddeposit nitrogen oxides directly into thestratosphere.

The pioneering work of Rowland,Molina, and Crutzen raised questions con-cerning human activity at least a decadebefore serious depletion of the ozone layerwas confirmed. These three scientists sharedthe 1995 Nobel Prize in chemistry for theirwork. Ozone measurements taken in themid-1980s demonstrated that humans weredoing serious damage to the ozone layer.This prompted international action andadoption of the Montréal Protocol in 1987.The original agreement called for cuttingthe use of CFCs in half by the year 2000.Subsequent strengthening of the agreementbanned the production of CFCs and halonsin developed countries after 1995. Halonsare compounds similar to CFCs that containbromine, for example, CF2BrCl. Productionof other ozone chemicals such as methylbromide, CH3Br, is also banned. Methylbromide later was widely used as a soilfumigant to treat agricultural land.

The phasing out of CFCs and otherozone-depleting chemicals has given rise to

new compounds to take their place.Hydrofluorochlorocarbons, or HCFCs, arebeing used as a short-term replacement forhard CFCs. Hard CFCs are those that do notcontain hydrogen, while soft CFCs containhydrogen. The presence of hydrogen atomsin HCFCs allows these compounds to reactwith hydroxyl radicals, OH, in the tropo-sphere. Some HCFCs can eventually beconverted in the troposphere to compoundssuch as HF, CO2, and H2O. Other HCFCscan be converted into water-soluble com-pounds that can be washed out by rain in thelower atmosphere. HCFCs provide a short-term solution to reversing the destruction ofstratospheric ozone. Their use signaled thefirst step in repairing the damage that hasoccurred over the last fifty years. Signs arepositive that the ozone layer can repair itself,but it may take another fifty years to bringozone levels back to mid-twentieth centurylevels.

Acid RainIf the pH of natural rain were measured,

you might expect a pH of around 7.0.Because a pH of 7.0 indicates neutral con-ditions, many people assume this to be thepH of rain. The theoretical pH of pure rain-water is actually about 5.6. Pure rain isacidic due to the equilibrium establishedbetween water and carbon dioxide in theatmosphere. Carbon dioxide and water com-bine to give carbonic acid:

CO2(g) � H2O(l) H2CO3(aq) H�(aq)

� HCO�3(aq)

Acid rain or precipitation refers to rain,snow, fog, or gaseous particles that have apH significantly below 5.6. There is noabsolute pH that defines acid rain, but a gen-eral guideline that can be used is that pre-cipitation below 5.0 can be consideredacidic. Although the term “acid rain” is used


in our discussion, consider rain to includeall forms of precipitation and moisture. Acidrain occurring as rain or snow is referred toas wet deposition. Dry deposition occurswhen gaseous acidic particles are depositeddirectly on surfaces.

Robert Angus Smith (1817–1884), aScottish chemist, was the first to identifyacid rain and study it in depth in 1852.Smith studied the differences in rain chem-istry in proximity to Manchester, England.He described how rain changed with respectto geography, fuel combustion, and localconditions. In 1872, he coined the term“acid rain” in his book Air and Rain: TheBeginning of a Chemical Climatology. Theexamination of acid rain accelerated around1920 when Scandinavian scientists startedto notice detrimental impacts on lakes intheir countries. As more data were collected,it became apparent that acid rain was aglobal phenomenon that required interna-tional attention. It has been a prominentenvironmental issue in this country since the1960s.

The principal cause of acid rain is thecombustion of fossil fuels that produce sul-fur and nitrogen emissions. The primarysources are electrical power plants, automo-biles, and smelters. Power plants producemost of the sulfur emissions and automo-biles most of the nitrogen emissions. Othersources of acid rain include nitrogen fertil-izers, jet aircraft, and industrial emissions.Just as in our discussion of ozone, numer-ous reactions are involved in the formationof acid rain. The process can be understoodby considering the transformation of sulfurand nitrogen oxides into their respectiveacidic forms: sulfuric acid and nitric acid.Sulfur, present up to a few percent in fuelssuch as coal, is converted to sulfur dioxidewhen the fuel is burned. The sulfur dioxidereacts with water to produce sulfurous acid,H2SO3(aq), that is then oxidized to sulfuricacid. The reactions are

S8(s) � 8O2 8SO2

SO2(g) � H2O(l) H2SO3(aq)

2H2SO3(aq) � O2(g) 2H2SO4(aq)

Nitric acid, HNO3, and nitrous acid, HNO2,form when nitrogen dioxide reacts withwater:

NO2(g) � H2O(l) HNO3(aq) � HNO2(aq)

Many reactions contribute to the overall for-mation of acid rain. These reactions ofteninvolve a number of complicated steps thatdepend on the atmospheric conditions. Thereactions just shown represent general reac-tions that form sulfuric and nitric acids, anddo not show the numerous reactions thatactually occur.

Acid rain has a number of negativeeffects on humans, plants, animals, buildingmaterials, and entire ecosystems. Acid raindirectly affects human health by irritatingthe respiratory passages. Elderly and asth-matic individuals are particularly suscepti-ble to acidic gases. Conditions includepersistent coughing, headaches, wateryeyes, and difficulty in breathing. Acid raincan have an indirect effect on human healthwhen drinking water is obtained from acid-ified lakes because acidic water can dissolvetoxic metals, such as lead and copper inpipes, introducing these into drinking water.

Acid rain has caused significant damageto forests throughout much of Europe. Insome areas of eastern Europe, as much ashalf of the forests have been damaged byacid rain. Acid rain has both direct and indi-rect effects on plants. Acid rain dissolves thewaxy protective coating of leaves interfer-ing with plant respiration. In areas whereacid rain is a problem, plant leaves oftenbecome discolored. For example, a yellow-ing discoloration of conifer needles hasbeen observed at high elevations. Trees mayalso shed their leaves or needles when sub-jected to acid rain. Such conditions weaken


with hydrogen ions to produce bicarbonate,HCO3

�. Bicarbonate reacts with hydrogenions to produce carbonic acid, H2CO3. Thereactions are

CaCO3(s) � H�(aq) Ca2�

(aq) � HCO3�

(aq)

HCO3�

(aq) � H�(aq) H2CO3(aq) CO2(g) � H2O

In areas where the geology is dominated bygranite, lakes have less buffering capacityand are much more susceptible to theimpacts of acid rain. Fish and aquatic organ-isms differ in their ability to adapt to acidicconditions. The natural pH of lakes isapproximately 8.0. The pH of poorly buf-fered lakes is between 6.5 and 7.0. Theeffects of pH on different aquatic organismsare summarized in Table 18.1.

Another detrimental effect of acid rainis the damage it does to buildings and otherstructures. Limestone and marble used asbuilding material are readily attacked byacid rain. The calcium carbonate in lime-stone and marble dissolves in acid. Thesematerials can also react with sulfur dioxideto produce gypsum, CaSO4 • 2H2O, whichis much weaker and more soluble than cal-cium carbonate. Many cultural treasureshave been severely damaged by acid rain,including well-known buildings such as theTaj Mahal, Westminster Abby, and theParthenon. European cathedrals and statuesthroughout the world are slowly dissolvingas they are subjected to acid rain. In Egyptand Mexico, it is estimated that more dam-age to ancient ruins has occurred in the last50 years than in the previous 2,500 years.Acid rain also attacks iron structures, cat-alyzing the formation of rust. In areas ofacid rain, the oxidation of metals is greatlyaccelerated, greatly increasing maintenancecosts.

To control acid rain requires interna-tional cooperation and much effort has goneinto crafting agreements between countries

plants and make them less resistant to dis-ease, insects, and extreme environmentalconditions, such as frost and wind. Inextreme cases, entire forest areas die. Agri-cultural plants are subjected to the sameimpacts as natural plants. Crop damage anddisease cause significant economic losseseach year throughout the world.

Besides affecting plants directly, acidrain affects plants indirectly by altering thepH of the soil, although some soils are nat-urally acidic. The pH of some Europeansoils subjected to acid rain has been loweredby one pH unit over the last 60 years.Because the pH scale is a logarithmic scale,this means that the soil is ten times moreacidic now than sixty years ago. Lower soilpH interferes with the ability of plants toabsorb nutrients. This is why it is importantto have the correct pH before fertilizing alawn. Soil acidity is typically neutralized byapplying lime (CaCO3) to the soil. In addi-tion to affecting the nutrient balance inplants, acid rain mobilizes and leaches outsoil cations such as Al3�, Mg2�, Ca2�, Na�,and K�. These cations dissolve in acidic soilwater. This results in the release of toxicmetals into the environment that can affect anumber of other organisms. This conditionis especially pronounced in the spring, whenrapid melting of winter snow causes a pulseof dissolved substances in runoff that floodsinto water bodies. Toxic metals in this runoffcan have a pronounced effect on embryosand eggs of amphibians, fish, and otheraquatic organisms.

Acid rain has resulted in the loss of lifein a number of lakes. The ability of a lake towithstand the impacts of acid rain is relatedto the geology of the lake’s basin. In areaswith limestone (calcium carbonate)deposits, a lake has a natural bufferingcapacity. The buffering capacity refers to theability to resist changes in pH. In well-buffered lakes, calcium carbonate reacts


fur dioxide and nitrogen oxides by the years2010 and 2000, respectively. The Title IVamendments call for an annual reduction insulfur dioxide emissions from power plantsof 10 million tons from a 1980 level of 19million tons. Significant progress has beenmade on reaching the 2010 goal. In 2002,the total emissions from power plants wasapproximately 11 million tons per year.Progress on reduction of nitrogen oxides haslagged behind that of other pollutants.While power plant emissions have fallenfrom 6.1 million tons per year in 1980 toapproximately 5.0 million tons per year in2002, total nitrogen oxides emissions haveincreased by 9% over the past twenty yearsdue to emissions from engines (used forpurposes other than transportation) anddiesel vehicles.

Industry has adopted both input andoutput approaches to reducing emissions.An input approach involves using fuels with

to reduce sulfur and nitrogen emissions. Theproblem with acid rain is that much of theemissions responsible for this problem areintroduced into the atmosphere through tallsmokestacks to alleviate local air qualityproblems. Reactions producing acid raintake place over several days as air massescarrying sulfur and nitrogen emissions movewith prevailing winds. When the moisture inthe air mass is released as precipitation, itmay be several hundred to several thousandmiles away from the source of emissions. Itis estimated that half of Canada’s acid rain isproduced in the United States. Similarly,emissions in England and Germany canimpact Sweden and Finland. In the UnitedStates, the Clean Air Act addresses the prob-lem of acid rain by requiring reductions insulfur and nitrogen emissions. Much of thelegislation is targeted at coal burning elec-trical power plants. Title IV of Clean Air Actenacted in 1990 mandates reductions in sul-

Table 18.1Effect of pH on Aquatic Organisms


and how much it reradiates back into space.Any object naturally radiates electromag-netic energy over a wide range of wave-lengths. The principal wavelength ofradiated energy is determined by the body’ssurface temperature. Wien’s Law states thatthe peak radiation is inversely proportionalto the absolute temperature:

In this equation, the peak wavelength is innanometers (1 nanometer, abbreviated nm, isequal to 10�9 meter) and T is the absolutetemperature in Kelvin. Wien’s equation showsthat as the surface temperature increases, thepeak wavelength decreases. Assuming theSun’s surface temperature is 6,000 K, the peakwavelength of radiation from the Sun isapproximately 483 nm. The peak wavelengthfor Earth with an average surface temperatureof 288 K is about 10,000 nm. The peak elec-tromagnetic energy radiated from the Sun isin the visible range, while that from Earth isprimarily infrared radiation. Radiation fromthe Sun is often termed shortwave radiationand radiation from Earth longwave radiation(Figure 18.2).

Incoming radiation from the Sun andbackradiation emitted by Earth interactswith the atmosphere. Although about half ofthe Sun’s radiation passes directly to Earth’ssurface, a portion is reflected back directlyinto space, while another portion isabsorbed by atmospheric gases and reradi-ated. Figure 18.3 shows the fate of radiationintercepting Earth. About half of the incom-ing solar radiation actually reaches the sur-face of Earth. The rest is reflected orabsorbed by the atmosphere or clouds.Infrared radiation reflected from Earth’s sur-face is partially absorbed and reflected bythe atmosphere and clouds. Some of thisradiation is reradiated back toward Earth’s

lower sulfur content. Soft bituminous coalsulfur’s content may be as low as a fewtenths of a percent compared to hard east-ern anthracite coal, which has sulfur con-tents of several percent. One problem,though, is the transportation of western coaloften makes its use economically unfeasi-ble. An alternative input approach is to usea process to remove sulfur from coal beforeburning it. Coal cleaning involves using aphysical process to separate pyrite, FeS,from coal by dispersing coal in oil andallowing the heavier pyrite to settle out.Another input approach includes actuallyswitching from coal to natural gas.Although natural gas emits less sulfur, thehigher combustion temperatures involvedproduces more nitrogen oxides. An outputapproach involves removing the sulfur fromthe flue gas after combustion. Emissionscan be removed using scrubbers that reduceharmful emissions. In scrubbers, chemicalsconvert a gas pollutant to less harmful prod-ucts. For example, a solution of calciumcarbonate can be sprayed down stacks emit-ting sulfur dioxide to produce solid calciumsulfate, CaSO4:

CaCO3(s) � SO2(g) CaSO3(s) � CO2(g)

2CaSO3(s) � O2(g) 2CaSO4(s)

The Greenhouse EffectThe greenhouse effect refers to an ele-

vation of the Earth’s temperature as a resultof changes in the Earth’s atmosphere due tohuman activities. The name comes from thefact that the atmosphere acts like glass in agreenhouse trapping radiation and causinga temperature increase. To understand thegreenhouse effect requires knowing howradiation interacts with gases to regulate theEarth’s temperature. The Earth’s surfacetemperature is primarily determined by theamount of radiation it receives from the Sun

62.9 � 10� �peak T


The ability of certain molecules in theatmosphere to absorb infrared radiation andreradiate back to Earth depends on themolecule’s structure. Molecules must pos-sess a dipole moment and have bonds capa-ble of vibrating at a frequency in theinfrared range. Because the major atmo-spheric gases N2, O2, and Ar do not possessa dipole moment, they do not absorbinfrared radiation. Gases capable of absorb-ing infrared radiation are referred to asgreenhouse gases. The most common green-house gas is water. Water’s ability to absorbinfrared radiation is demonstrated by thewarming effect of clouds.

Human activities have changed the con-centration of a number of greenhouse gases.Unlike water, these gases exist in traceamounts and the relative human impact on their concentrations is significant. Theprimary greenhouse gases include carbondioxide, methane (CH4), CFCs, and nitrous

surface, causing a raise in temperature. Theatmosphere is responsible for Earth’s tem-perature being about 30°C higher than itwould be if the atmosphere was not present.In fact, the atmosphere is responsible for anatural greenhouse effect. Many scientistsdefine the natural temperature increase asthe greenhouse effect and the temperatureincrease due to human activities that mod-ify the atmosphere as the enhanced green-house effect. We will define the greenhouseeffect as excess warming due to humanactivities, but keep in mind that a naturalgreenhouse effect preceded humans on ourplanet. In summary, Earth’s surface temper-ature is due to the balance between incom-ing and outgoing radiation. The fact that theatmosphere is more transparent to short-wave solar radiation compared to longwaveradiation from Earth leads to a naturalgreenhouse effect that produces a surfacetemperature of 288 K (15°C or 59°F).

Figure 18.2Peaks of Solar and Earth Radiation (Rae Déjur)


tury begins. The controversy has shiftedfrom uncertainty concerning whetherhumans affect Earth’s climate to how muchhumans are affecting the climate and whataction, if any, should be taken. A major dif-ficulty surrounding the greenhouse effect isseparating human influences from the cli-mate’s natural variability. While there arestill skeptics, the consensus seems to be thathumans are affecting the climate, butexactly how, and to what extent, continuesto be debated. The issue is complicated bynumerous factors. The oceans play a majorrole in the carbon cycle because they have agreat capacity to absorb carbon dioxide.Vegetation also affects global carbon con-centrations through photosynthesis. Somestudies have indicated that primary produc-tivity increases with elevated CO2, whileothers show no increase. A major humanfactor is changes in land use, for example,

oxide (N2O). Carbon dioxide is the mostimportant greenhouse gas, being responsiblefor roughly 60% of the warming. The atmo-sphere’s CO2 concentration has risen fromabout 280 ppmv (parts per million by vol-ume) to 360 ppmv in the last 150 years. Thisincrease is due primarily to the burning offossil fuels. Land clearing has also contrib-uted to increased CO2 levels. Methane is thenext most important greenhouse gas, con-tributing about 20% to global warming.Methane is produced by methanogenic bac-teria under anaerobic conditions. Decayingvegetation from wetlands and rice paddiesproduces large quantities of methane. It alsois produced in the digestive tract of livestock,in landfills, and by biomass combustion.Nitrous oxide contributes about 5% to thegreenhouse effect.

The greenhouse effect continues to be acontroversial topic as the twenty-first cen-

Figure 18.3Earth’s Radiation Budget (Rae Déjur)


1990s to reduce greenhouse gases. Anyworkable agreement must include theUnited States. The United States does notwant to jeopardize its economy and believesthe overall benefits of entering into an agree-ment must outweigh the costs. An agree-ment must also be enforceable. As theworld’s number one emitter of carbon (Table18.2), the United States holds the key toglobal reductions.

Water QualityWater is essential for human life.

Although Earth is often called the “waterplanet” with nearly three-fourths of Earth’ssurface covered by water, most of itstremendous supply of water is salty and notdirectly available for human use. Only 3%of the planet’s water is fresh. Approximately70% of the human body is composed ofwater, and we can survive only a few dayswithout it. Most of the water in the world,and the largest percentage in this country, isused for agriculture. Large quantities ofwater are also used in power generation, fordomestic use, and in industry. Table 18.3

deforestation and agriculture. Other issuesinclude variations in solar output and therole of clouds in Earth’s energy balance.

During the last decade national repre-sentatives have attempted to craft an agree-ment to reduce global greenhouse gasemissions. At the Earth Summit held in Riode Janeiro, Brazil, in 1992, industrializednations agreed to stabilize emissions at their1990 levels by the year 2010. This nonbind-ing agreement was seen as inadequate bymany nations, and several subsequent inter-national meetings have attempted to resolvedisputes between countries regarding reduc-tion in emissions. The European nationshave pushed for more aggressive reductionsthan the United States. The United States isalso dissatisfied that many developing coun-tries, including the two most populous coun-tries on the planet, China and India, areexempt from making immediate emissionreductions. The last significant negotiationstook place in December of 1997 in Kyoto,Japan. The nations at the Kyoto Summit hadtrouble agreeing how to curb greenhousegases. The European Union wanted reduc-tions of 15% below 1990 levels. Severalother countries led by the United Statesresisted such deep cuts, fearing they wouldcause too much harm to their economies.The last round of talks on the Kyoto Sum-mit took place in November 2001 inMorocco. The meeting in Morocco con-cluded four years of negotiations andrequired industrialized countries to reducegreenhouse emissions by 5% below their1990 levels by the year 2012. The U.S. rep-resentative at the meeting announced theBush administration’s dissatisfaction withthe Kyoto agreement. The administrationdoes not plan to sign the agreement and isgoing to put forth its own solution for reduc-ing emissions. This stance has provokedharsh criticism from European leaders andmay unravel much of the work done in the

Table 18.2Carbon Emissions by Country


Water pollution results from a numberof sources including agricultural runoff, ero-sion, industrial wastes, domestic wastes, androad runoff. Water pollution can be classi-fied as either point source or nonpointsource. Point source pollution is emittedfrom a specific, well-defined location suchas a pipe. Nonpoint sources refer to pollu-tants dispersed over a wide area from manydifferent areas. A sewer pipe would be anexample of a point source and fertilizerrunoff from a field would represent nonpointsource pollution.

A major water quality problem occurswhen waterways receive too many nutrientssuch as phosphate or nitrate. This occurs inareas subjected to increased agriculturalrunoff, sewage effluent, and erosion. Excessnutrients fertilize water bodies, and thisresults in phytoplankton blooms. Increasedalgae results in large zooplankton blooms.When the heavy plankton populations dieand undergo decomposition by bacteria,oxygen levels in the water drop. In extremecases, the oxygen levels in the lower levelsof the lake can drop to zero. This acceleratedinput of nutrients causes premature aging ofthe lake. The process is called culturaleutrophication. Up until 1980, detergentswere a major source of phosphate input towaterways and a primary cause of culturaleutrophication. Many detergents and soapproducts now carry a “no phosphate” labelto let the consumer know that the product isphosphate free.

One of the foremost global issuesregarding water quality is access to safedrinking water. Currently, one-sixth of theworld’s 6.1 billion people do not have accessto a source of clean water, and 40% do nothave adequate sanitation facilities. Pollutedwater is responsible for diseases such astyphoid, cholera, and dysentery. Three mil-lion people die annually due to the latter.The discovery that diseases were transmit-

gives the approximate use of water in ourdaily activities.

Water is often called the universal sol-vent. Natural water is far from pure andexists as a natural solution containing manydissolved minerals. Hard water contains anabundance of divalent metals, primarilyCa2� and Mg2�. Water hardness is measuredin parts per million of CaCO3. Soft waterhas less than 60 ppm CaCO3, moderatelyhard water has concentrations between 61and 120 ppm CaCO3, and hard water hashardness values greater than 120 ppmCaCO3. Water hardness does not present ahealth concern until it reaches several hun-dred ppm. When the hardness of water isdue to dissolved calcium and magnesiumbicarbonates, Ca(HCO3)2 and Mg(HCO3)2,it is termed carbonate hardness or tempo-rary hardness. It is temporary because whenwater containing carbonate hardness isheated, the carbonates precipitate out. Thiscan lead to scales developing in pipes andboilers used in hot water heating systems.Permanent hardness consists of sulfatesand chlorides that remain in solution whenheated.

Table 18.3Daily Water Use


One drawback of chlorine is that it can reactwith organic materials present in naturalwaters to produce harmful substances. Tri-halomethanes are organic compounds withthe general formula CHX3, where x repre-sents chlorine or bromine. Whenhypochlorous acid reacts with organic mat-ter, chloroform, CHCl3, can be produced.Chloroform, which was one of the firstanesthetics used in 1847, is a suspected car-cinogenic. Because of this, the drinkingwater standard for trihalomethanes has beenset at 100 parts per billion. Although chlo-rine is still the most widely used disinfec-tant in the United States, some communitieshave switched to other disinfectants. Chlo-rine dioxide and ozone are two commonalternatives. Ozone is used widely through-out Europe for treating drinking water. It isa much more effective disinfectant and hasthe ability to kill viruses. Its drawback isthat it cannot be stored so it must be gener-ated on site. It also does not provide theresidual protection that chlorine provides.Residual chlorine is the chlorine dissolvedin the water once it leaves the treatmentplant. This residual chlorine protects thewater from recontamination. Ozone-treatedwater does not have the residual protectiononce it leaves the treatment plant.

To protect public health, governmentagencies have established guidelines fornumerous chemicals in drinking water. In1974, the United States passed the SafeDrinking Water Act. As part of this act, theEPA determined safe levels for chemicals indrinking water. The maximum contami-nant level (mcl) is the highest acceptableconcentration for a chemical in drinkingwater. Although not an absolute guaranteeto ensure safety, the mcl is a guideline thatestablishes a maximum acceptable value forchemicals in water. Table 18.4 shows themaximum contaminant level for a numberof substances in drinking water.

ted by bacteria through water in the mid-nineteenth century led to advances in sani-tation and water treatment. Severalchemicals are used as disinfectants to killpathogenic organisms in water. The mostcommon disinfectant in this country is chlo-rine. The first public water supply to betreated with chlorine was in England in1904 and in the United States in 1908.Throughout the twentieth century, as thepractice of water chlorination spread, theincidence of water-borne diseases hasdecreased.

As mentioned, chlorine, in the form ofCl2 or the hypochlorite ion, OCl�, is used asa disinfectant. Free chlorine, Cl2, reacts withwater to produce hypochlorous acid, HOCl,and hydrochloric acid, HCl:

Cl2(g) � H2O(l) HOCl(aq) � HCl(aq)

The hypochlorous acid oxidizes the cellwalls and kills bacteria. Solid calciumhypochlorite, Ca(OCl)2, and liquid solutionsof sodium hypochlorite, NaOCl, can be usedto generate hypochlorous acid in place ofchlorine gas, for example, in chlorinatingswimming pools. The hypochlorite ion gen-erated from Ca(OCl)2 and NaOCl forms anequilibrium with water represented by theequation:

OCl� � H2O HOCl � OH�

The use of hypochlorites tends to raise thepH of water because hydroxide ions are pro-duced, as opposed to chlorine gas, whichtends to lower the pH due to the productionof hydrochloric acid. Bleach is a dilute solu-tion (approximately 5%) of sodiumhypochlorite that is often used as a house-hold disinfectant.

Chlorine has been criticized as a waterdisinfectant because of undesirable reac-tions that produce carcinogenic substances.


der, and kidney. Initial symptoms includeskin discoloration. Arsenic contamination isa serious problem in a number of countries,especially in rural areas that depend on wellwater as a drinking source. In Bangladesh,it is estimated that millions of peopledepend on water sources contaminated witharsenic. In the United States, the EPA low-ered the arsenic mcl from 50 parts per bil-lion to 10 parts per billion during the waningdays of the Clinton administration in 2000.The standard was lowered due to evidencethat arsenic in drinking water poses a highercancer risk than once thought. This rulingwas initially suspended by the Bush admin-istration pending further review, but the 10ppb was finally accepted in late 2001. Waterutility companies in areas where arsenic isnaturally high fear excessive costs in meet-ing the stricter standard.

Mercury, known in ancient times asquicksilver, is another metal that has beenstudied for years. Mercury is widely used inbatteries, thermostats, scientific instruments(thermometers and barometers), alloys,fungicides, and precious metal extraction.During the 1800s, mercury (II) nitrate,Hg(NO3)2, was used in the curing process tomanufacture felt hats. Individuals workingin the hat industry had a high rate of mer-cury poisoning, although at the time theeffects of mercury on humans wereunknown. Because mercury poisoningresults in neurological damage, the phrase“mad as a hatter” became associated withhat workers who used mercury compounds.Mercury poisoning symptoms includeslurred speech, tremors, and memory loss.Lewis Carroll’s character the Mad Hatterfrom Alice in Wonderland, written in 1865,was based on the perception of aging hatworkers going mad.

The greatest environmental concerninvolving mercury involves methylatedforms of mercury. These forms result when

Most of the substances listed in Table18.4, as well as a number of others, havebeen associated with serious environmentalproblems in this country and many others.Arsenic, mercury, and lead are metals thatseriously affect human health when found inexcessive amounts in drinking water.Arsenic is the quintessential poison, appear-ing in Shakespeare, Sherlock Holmes, andmodern murder scenes. It has been knownsince ancient times and was traditionallyused as an antibiotic; it was commonly usedto cure syphilis until the advent of penicillinand modern antibiotics. Arsenic’s extensiveuse in pesticides, primarily before theadvent of organic pesticides, and its unin-tentional release during mining are majorsources of water contamination. Arsenicpoisoning can lead to several forms of can-cer including cancer of the skin, lungs, blad-

Table 18.4Maximum Contaminant Level


restricted in recent years as the detrimentaleffects of lead on human health have beendiscovered. In the 1970s, leaded gasolinebegan to be phased out and in the 1980s leadwas reduced in paints. Drinking water isgenerally not a major source of lead inhumans, but it can be as high as 60% ininfants drinking water-based milk formula.Lead in drinking water is associated with theuse of lead pipes and lead solder used inplumbing systems. This is especially preva-lent in older water systems built before 1930where the water is soft. Hard water containscarbonates that form protective deposits onthe interior of pipes. In some systems wherethe possibility of lead contamination frompipes exists, phosphates are added to thewater to form protective deposits on theinterior of pipes. Ironically, new homes lessthan five years old may also suffer from leadproblems because their pipes have not beencoated with deposits. The 1986 amendmentsto the Safe Drinking Water Act prohibitedthe use of lead pipe and solder in all publicwater systems.

Water pollution from benzene, C6H6,results from its presence as an octanebooster in gasoline. Health effects of ben-zene include anemia, suppression of theimmune system, nerve damage, andleukemia. Groundwater contamination frombenzene, and its associated derivates such astoluene and xylene, continues to be a preva-lent problem due to leakage from under-ground storage tanks. Millions of thesetanks exist at gas stations and other localesthroughout the United States. Starting in the1990s, the EPA began an intensive effort toput into place a comprehensive series of reg-ulations regarding the monitoring andreplacement of underground storage tanks.

PCBs, or polychlorinated biphenyls, con-sist of a group of chlorinated organic chemi-cals that found widespread use startingaround 1960. The biphenyl molecule consists

bacteria convert Hg2� into methyl mercurycompounds during anaerobic respiration.Dimethyl mercury, CH3HgCH3, methylmercury, CH3Hg�, and other methylatedmercury compounds appear in the sedi-ments where anaerobic bacteria exist. In thismanner, mercury is incorporated into thefood chain and ends up in fish consumed byhumans. A number of lakes in the UnitedStates and Canada contain mercury-contaminated fish. Fish with the highestconcentrations are large, carnivorous fishthat reside high on the food chain, such asnorthern pike and largemouth bass. In areasknown to be affected, state agencies haveposted warnings, and people are advised toregulate their consumption of fish. Mercurypoisoning is especially detrimental to nativecommunities whose diet has traditionallyconsisted of large quantities of fish. Saltwa-ter species such as shark, mackerel, andswordfish are often associated with highmercury levels. The Food and Drug Admin-istration has issued warnings concerningconsumption of these species, and pregnantwomen are especially advised to monitortheir fish intake.

Lead is yet another metal that is associ-ated with environmental problems. It hasbeen used for thousands of years and is cur-rently used in lead-acid batteries, nuclearsheathing, pipes, paints, solder, and alloys.Leaden ware is often cited as a primarysource of poisoning in Roman times. Stud-ies on the bones of ancient Romans haveshown lead levels several orders of magni-tude higher than those of modern humans.Lead poisoning can result in anemia, paral-ysis of joints, fatigue, reproductive prob-lems, and colic. Severe cases of leadpoisoning result in damage to the brain, kid-ney, and nervous system. In recent years,lead has been implicated in the delayedmental development of children, resulting inlower IQ scores. The use of lead has been


During the last thirty years the UnitedStates has made significant progress inreducing pollution to the nation’s waters.The Clean Water Act passed by Congress in1972 has been the primary legislationresponsible for curbing water pollution dur-ing this period. This act mandated stricterwater standards and regulated those entitiesthat discharged effluent into the nation’swater. Federal permits were required formany discharge activities, for example,municipal treatment plant effluent. TheClean Water Act also provided grant moneyfor municipalities to improve their waste-water treatment facilities by building newplants and applying modern technology toupgrade existing facilities. Although themedia often highlights pollution incidentsthat tend to capture attention, the overallwater quality in the United States hasgreatly improved since passage of the CleanWater Act.

Air PollutionJust as water bodies such as rivers,

lakes, and the ocean can be considered enor-mous aqueous solutions, the atmosphere is agigantic gaseous solution. The atmosphereis composed primarily of nitrogen, 78%, andoxygen, 21%. The third most abundant gasin the atmosphere is argon, which makes upabout 0.9%. The remaining 0.1% consists of

of two benzene rings joined by a single bond.It is formed by heating benzene in the pres-ence of a catalyst:

PCBs are produced when biphenyl is reactedwith chlorine and the chlorine atoms replacethe hydrogen atoms. More than two hundreddifferent PCBs can be produced. BecausePCBs are nonflammable, inert, and poor con-ductors, they made ideal electrical insulatorsin transformers. PCBs also found use inhydraulic fluids, plasticizers, pesticides, copypaper, cutting oils, and fire retardants. Unfor-tunately, it was discovered that PCBs mightbe associated with liver, gastrointestinal, andreproductive problems. Some workers in thePCB industry developed a permanent, disfig-uring form of acne called chloracne. Animaltests showed cancers developed in test ani-mals. Another concern was that PCBs werefat soluble and bioaccumulated in the foodchain. Because the structure of PCBs is sim-ilar to DDT, it was feared that their continueduse could have severe consequences forwildlife (Figure 18.4). Although the acutetoxicity of PCBs is low, there was enoughevidence that PCBs could have long-rangehealth effects to prompt Monsanto, their soleproducer in North America, to discontinuetheir production.

Figure 18.4DDT and PCB. The Xs in PCB denote where chlorine atoms may attach.


the respiratory system. The NAAQS pollu-tants are called criteria pollutants. Criteriapollutants are generic pollutants that allstates and local governments need to con-sider. The Clean Air Act has specific regu-lations regarding emissions of many othersubstances such as toxic chemicals, butthese substances have unique sources oftenassociated with specific industries.

Chemicals can be labeled as either a pri-mary air pollutant or secondary air pollu-tant. Primary air pollutants are those such ascarbon monoxide and sulfur dioxide thatenter the atmosphere directly as a result ofhuman or natural events. Carbon monox-ide’s primary source in the atmosphere is theincomplete combustion of gasoline. Hun-dreds of different chemicals are present ingasoline. The combustion of octane, C8H18,can be used to represent the general reactionof hydrocarbons in an automobile engine toproduce energy:

2C8H18 � 25O2 18H2O � 16CO2 � energy

Carbon monoxide results when insufficientoxygen is present during the combustionprocess. Carbon monoxide’s toxic effectresults from the fact that hemoglobin’s affin-ity for CO is more than 200 times greaterthan it is for oxygen. Because carbonmonoxide displaces oxygen in the blood,breathing air with high carbon monoxidelevels results in a lack of oxygen. The effectsof carbon monoxide poisoning depend onthe concentration and length of exposure(Table 18.6). If the CO concentration is highenough, people can slowly suffocateunaware (CO is odorless and invisible) thatthey are suffering from CO poisoning. Sev-eral hundred people die each year in theUnited States due to CO poisoning, mainlydue to faulty furnaces or using gas heatersin enclosed areas without proper ventilation.While normal CO levels in areas with heavy

a variety of gases including carbon dioxide,water, xenon, helium, methane, and ozone.The previously mentioned problems ofozone depletion, acid rain, and globalwarming illustrate how human impacts onthe chemical composition of trace gases inthe atmosphere can produce serious prob-lems. On the local level, air quality is aproduct of many factors such as climate,weather, population density, land use, indus-try, and transportation.

The Clean Air Act of 1967 forms thelegal basis for air pollution control in theUnited States. This act and its subsequentamendments give the federal governmentthe right to establish national ambient airquality standards (NAAQS) for major pol-lutants. These standards are intended to pro-tect public health as well as to protect publicwelfare. The latter includes protecting visi-bility, nonhuman species, and structures.The NAAQS for the six major categories ofair pollutants are given in Table 18.5. Table18.5 shows that while several of the pollu-tants are gases, lead and particulate matterconsist of solid particles suspended in air.PM-10 refers to solid particles with diame-ters of less than 10 micrometers. Particles ofthis size are the most detrimental to humanhealth because they can be inhaled deep into

Table 18.5NAAQS


ular basis. Photochemical smog formsthrough a complex series of reactions in-volving nitrogen oxides, hydrocarbons, andsunlight in relatively warm climates (Figure18.5). Automobile exhaust contains nitrogenoxides (NO and NO2) and hydrocarbons thatare the two main ingredients necessary forthe formation of photochemical smog.Nitrogen dioxide (NO2) is a brown gas thatproduces an irritating odor. When the sunrises in areas with high concentrations ofnitrogen dioxide and hydrocarbons, numer-ous reactions occur to produce a variety ofphotochemical oxidants. For example,ozone is formed by the following reactions:

NO2(g) � sunlight NO(g) � O(g)

O(g) � O2(g) O3(g)

It is interesting to note that ozone formationin the stratosphere is desirable, but ozoneformation in the troposphere is not. PANresults from a complex series of reactions

traffic congestion are not lethal, they canresult in drowsiness, headaches, irritability,and impaired judgment.

Secondary pollutants result from chem-ical reactions that occur in air. Sulfuric acidand nitric acid responsible for acid rain areexamples of secondary pollutants. Anotherlarge class of secondary pollutants is pho-tochemical oxidants. Photochemical oxidants include ozone, PAN (peroxyacyl-nitrate), and formaldehyde. A common termfor photochemical oxidants is photochemi-cal smog. Smog, a term derived from com-bining “smoke” and “fog,” was originallycoined to described polluted air in London.Smog is particularly characteristic in citieswhere industrial pollution occurs. Photo-chemical smog is another type of smog thatoccurs mainly in dry climates where there isample sunlight and an abundance of cars.Los Angeles and Mexico City are citieswhere photochemical smog occurs on a reg-

Table 18.6Effects of CO Poisoning


primates grooming each other, we can imag-ine our ancient ancestors performing simi-lar activities to combat pests. Agriculture asa means of obtaining food started approxi-mately 10,000 years ago, and this increasedthe number of species that could be classi-fied as pests. Several methods can be usedto control pests. These include physical,genetic, biological, and chemical methods,but in this section only chemical controlswill be considered. The word “pesticide” is a collective term in that it includes nu-merous other “cides” such as herbicides,insecticides, rodenticides, fungicides, bac-tericides, and avicides.

The use of chemical pesticides can betraced back several thousand years. TheSumerians, occupying present-day Iraq,burned sulfur compounds to produce fumi-gants. The Greeks also employed this prac-tice as early as 1000 B.C. Evidence existsthat the Chinese extracted pesticides from

involving nitrogen oxides, oxygen, alde-hydes, and hydrocarbons. PAN is a power-ful eye irritant and irritates the respiratorysystem.

Just as the Clean Water Act hasimproved water quality in the United States,the Clean Air Act has greatly improved airquality since it was passed in 1967 andamended several times after. States andmunicipalities must meet the NAAQS orface the loss of federal dollars for road projects. Programs such as vehicle emissioninspections, reduction in speed limits, emis-sion permitting, alternative fuels, and use ofbest available technology have served toclean the nation’s air.

PesticidesAs long as humans have existed they

have been plagued by pests. Today when weobserve animal species rolling in mud or

Figure 18.5Smog Formation (Rae Déjur)


Müller (1899–1965) discovered that thecompound dichlorodiphenyltrichloroethane(DDT) was an effective insecticide (Figure18.6). DDT had first been synthesized in1873, but it was Müller who discovered itsefficacy as an insecticide. DDT was initiallymarketed in 1941, and found its firstwidespread use in World War II. DuringWorld War I several million deaths, includ-ing 150,000 soldiers, were attributed totyphus. There are several forms of typhus,but the most common form is due to bacte-ria carried by lice. During World War II,fearing a repeat of typhus outbreaks, theAllied forces used DDT to combat not onlytyphus but also malaria, yellow fever, andother diseases carried by insects. Soldiersliberally applied talcum powder with 10%DDT to clothes and bedding to kill lice.American and European allies were rela-tively free from typhus and other diseases,while the Germans who did not use DDThad many more noncombat deaths due toinfectious diseases. In liquid form, DDTwas used in the Pacific Theater to preventmalaria and yellow fever. In addition to itsuse in the war, DDT was employed in trop-ical areas as a generic insecticide to preventinfectious diseases, especially malaria. Oncethe war ended, its use to advance publichealth in tropical undeveloped countries wasexpanded for use in agriculture in developedcountries. Paul Müller was awarded the

plants as early as 1200 B.C. The first pesti-cides used widely in recent history were poi-sonous metals. In France in the earlyeighteenth century a mixture of copper sul-fate (CuSO4) and lime (CaCO3) was sprayedon grapes to make them unappealing tothieves. It was noted that this mixture alsoprotected the grapes from certain fungi, andthus a pesticide known as Bordeaux mixturecame into existence. Paris green, a com-pound of copper and arsenic, was a commonpesticide used throughout the eighteenthcentury. Lead arsenate, Pb3(AsO4)2, and sev-eral metal compounds containing antimonyand mercury started to replace Paris greenat the beginning of the twentieth century.Metal, inorganic pesticides were commonthroughout the first half of the twentiethcentury. Sodium fluoride and boric acid,B(OH)3, were commonly used as ant andcockroach poisons; hydrogen cyanide(HCN) was a fumigant. Nicotine solutionsmade from soaking tobacco leaves in waterhave been used for several hundred years asinsecticides and rodenticides.

The common inorganic compoundsused as pesticides have several problems.Although they kill the target pest, they tendto be very lethal to humans and other mam-mals. Pesticides that kill indiscriminatelyare classified as broad spectrum pesticides.This means they do not distinguish betweentarget and nontarget species, but kill manyharmless and helpful species. Another majorproblem with the inorganic pesticides is thatthey are persistent in the environment, last-ing years and even decades. A final majordetriment to the widespread use of inorganicmetal compounds is that they are expensive.To try to alleviate some of these problemsand produce large quantities of pesticidesthat were economically feasible, chemiststurned to synthetic organic pesticides in the1930s. In 1939, while working for the Geigychemical company, the Swiss chemist Paul

Figure 18.6DDT


Nobel Prize in medicine in 1948 for his dis-covery of the insecticide potential of DDT.By 1950 DDT and several related com-pounds were viewed as miracle insecticidesthat were inexpensive and that could be usedindiscriminately.

Even though DDT seemed to be a cheapand effective pesticide, enough was knownin its early development to raise concerns.DDT is a persistent chemical that lasts along time in the environment. DDT is fatsoluble and not readily metabolized byhigher organisms. This meant that DDTaccumulated in the fat tissues of higherorganisms. As organisms with longer lifespans residing higher on the food chain con-tinually fed on organisms lower on the foodchain, DDT would accumulate in their tis-sue. For example, the concentration of DDTin a lake might be measured in parts per tril-lion, plankton in the lake may contain DDTin parts per billion, fish a few parts per mil-lion, and birds feeding on fish from the lakeseveral hundred parts per million. The accu-

mulation of a chemical moving up the foodchain is a process known as biological mag-nification (Figure 18.7). Another concernraised was that certain pests seemed todevelop an immunity to DDT and the appli-cation rate had to be increased to combatinsects. This was because natural selectionfavored insects that had the genetic charac-teristics to survive DDT and passed thisability on to their offspring. Direct deaths ofbird and fish populations had also beenobserved in areas with heavy DDT use. Theproblems associated with DDT and otherpost-WWII organic pesticides became anational concern with Rachael Carson’s(1906–1964) publication of Silent Spring in1962. Carson’s book alerted the public to thehazards of insecticides, and while not call-ing for a ban, challenged the chemical andagricultural industry to curtail itswidespread use of chemical pesticides. Mostdeveloped countries started to ban the use ofDDT and related compounds in the late1960s. DDT was banned in the United

Figure 18.7Biological Magnification (Rae Déjur)


States in 1973. Although banned in devel-oped countries, its use to improve publichealth in undeveloped countries continues.The World Health Organization estimatesthat DDT has saved 25 million lives frommalaria and hundreds of millions of otherlives from other diseases.

DDT belongs to a group of chemicalinsecticides known as organochlorides.These contain hydrogen, carbon, and chlo-rine and kill by interfering with nerve trans-mission; hence, they are neurotoxins.Organochlorines were the dominant type ofchemical insecticide used from 1940 to1970. Some common organochlorinesbesides DDT are chlordane, heptachlor,aldrin, and dieldrin. Because of their prob-lems, and subsequent ban in many regions,numerous other classes of insecticides havebeen synthesized to replace organochlo-rines. Several major groups includeorganophosphates, carbamates, andpyrethroids.

Organophosphates were originally dis-covered in Germany during World War IIwhile researchers were searching for a sub-stitute for nicotine. Nicotine extractsobtained by soaking tobacco in water werecommonly used as an insecticide as early as1700. Organophosphates are derived fromphosphoric acid, H3PO4, and, likeorganochlorides, act on the nervous system

(Figure 18.8). They work by interfering withspecific enzymes in the nervous system.Organophosphates are more lethal tohumans and other mammals thanorganochlorides. Because of this, chemicalweapons collectively known as nerve gas areorganophosphates. The chief advantage oforganophosphates is that they are unstableand break down in days to weeks in the envi-ronment. Although this reduces the toxicityto humans on food due to pesticide residues,organophosphates pose a hazard to agricul-tural workers and nontarget species.Because of this, several countries and dif-ferent states in the United States ban specificorganophosphates. Some commonorganophosphates are parathion, malathion,and diazinon (Figure 18.8).

Carbamates act similar to organophos-phates by interfering with nerve systemenzymes. They are derived from carbamicacid, H2NCOOH, and share many of theproperties of organophosphates (Figure 18.9).The carbamate known as carbaryl (commer-cial trade name is Sevin) was synthesized in1956 and has been used extensively since thattime. Carbaryl’s advantages are that its toxic-ity to mammals is low and it kills a broadspectrum of insects. It is widely used in anumber of household lawn and garden prod-ucts. Other common carbamates includealdicarb (Temik) and carbofuran (Furadan).

Figure 18.8Organophosphates such as Malathion are derived from phosphoric acid.


Pyrethroids are based on mimicking thestructure of the natural insecticide pyrethrin.Pyrethins are found in the flowers ofchrysanthemums. Ground flowers were tra-ditionally used to obtain pyrethin insecti-cides and used to kill lice in the early 1800s.Synthetic pyrethins were first produced inthe early 1970s. The exact nature of howpyrethroids work is unknown, but becausethey paralyze insects it is speculated thatthey affect the nervous or muscular system.Pyrethroids are effective in low dosages andare nonpersistent.

Our discussion of pesticides has focusedprimarily on insecticides. In the United Statesthe primary use of pesticides is in the form ofherbicides, those pesticides used to controlweeds. Approximately 70% of the pesticidesused in the United States are herbicides and20% are insecticides. The use and develop-ment of herbicides parallels that of insecti-cides. The first herbicides were inorganicmetal compounds and salts. During WorldWar II organic herbicides were synthesizedand their use increased dramatically. One ofthe first major classes of herbicides synthe-sized in the mid-1940s was phenoxyaliphaticacids. As this name implies, the phe-noxyaliphatic acids contain the benzene ring,oxygen, and an aliphatic acid. The two mostcommon phenoxyaliphatic acids are 2,4dichlorophenoxyacetic acid, called 2,4-D and2,4,5 trichlorophenoxyacetic acid, known as2,4,5-T (Figure 18.10). The numbers in these

compounds refer to the carbon to which thechlorine atom is attached. Dichlorophenoxy-acetic acid works by producing excessgrowth hormones in plants. Plants die due totheir inability to acquire sufficient nutrients.The herbicide 2,4,5-T acts as a defoliant thatcauses plants to shed leaves. During the Viet-nam War, 2,4-D and 2,4,5-T were combinedinto a formulation called Agent Orange.Agent Orange was applied extensively toforests in Vietnam to expose enemy posi-tions. Subsequent birth defects in children ofboth Vietnamese and American soldiersexposed to Agent Orange became a contro-versial topic in postwar years. Research hasshown that in the manufacture of 2,4,5-T sidereactions occur that produce small quantitiesof dioxins. Dioxins are compounds that arecharacterized by the dioxin structure:

Dioxins have been shown to cause birthdefects in tests with laboratory animals.Long-term effects include certain cancers.Because of the health concerns associatedwith 2,4,5-T, this herbicide was banned inthe United States in 1983.

Another class of herbicides is called tri-azines. Triazine compounds consist of abenzene-like structure with alternatingnitrogen and carbon atoms in a hexagonal

Figure 18.9Carbamates such as carbaryl are derivedfrom carbamic acid.

Figure 18.10The Two Most Common PhenoxyaliphaticAcids


ring (Figure 18.11). Amino groups areattached to two of the carbons in the ringand chlorine is attached to the third. Themost common triazine is called atrazine.Atrazine disrupts photosynthesis in plants.Atrazine is widely used on corn plantsbecause corn has the ability to deactivateatrazine by removing the chlorine atom.

Another common herbicide is paraquat.Paraquat belongs to a group of herbicidescalled bipyridyliums. The name comesfrom the combination of two pyridine rings(Figure 18.12). Paraquat works rapidly bybreaking down the cells responsible for pho-tosynthesis. It is used as a preemergent her-bicide, which means it is applied to soilbefore plants emerge. Paraquat has beenwidely used to destroy marijuana plants. Itsuse to combat marijuana crops has led tospeculation that some marijuana users maybe susceptible to lung damage from mari-juana contaminated with paraquat.

Since their development sixty years ago,synthetic pesticides have increased thirty-fold. Worldwide use of pesticides is approx-

imately 3 million tons, with half of this inthe United States and Europe. The use ofsynthetic chemical pesticides continues tobe a controversial topic. There are strongarguments both for and against the use ofpesticides. Positive aspects of pesticide useinclude their ability to save lives and reducehuman suffering, increase food supplies,and lower food costs because they are rela-tively cheap compared to other forms of pestcontrol. On the negative side, pesticidespose some health risks to humans—bothshort term and long term—kill nontargetspecies, have uncertain ecological impacts,and produce hardier pests over time. Sincethe creation of the EPA in 1970, the federalgovernment has heavily regulated the use ofsynthetic chemical pesticides. Over fiftypesticides have been banned during the lastthirty years. Certain pesticides are availableonly to commercial users who have the cre-dentials to apply them. New pesticides mustundergo a rigorous screening that includesvarious laboratory tests that seek to ascer-tain the potential human and environmentalrisks of the pesticide. Currently, approxi-mately 600 different active ingredients areapproved by the EPA for pesticide use. Still,even with extensive federal regulations,many people feel oversight is lax and poorlyenforced. Most of the chemicals approvedby the EPA have never undergone extensivetesting. Recent international trade agree-ments, such as NAFTA, mean that pesti-cides banned in the United States could becrossing our borders indiscreetly.

Figure 18.11Triazine Structure

Pyridine Bipyridine Paraquat

Figure 18.12Paraquat is a combination of two pyridine rings.


Although a decrease in the use of syn-thetic chemical pesticides is not foreseen,scientists continue to explore the productionof safer pesticides and other methods toaugment chemical pesticides. Integratedpest management (IPM) is an approach forcontrolling pests that utilize multiple tech-niques in an ecological systematic fashion.In this multifaceted approach, techniquessuch as cultivation methods, genetic engi-neering, natural biological controls, andinsect sterilization are combined with chem-ical techniques to control pests. As IPMmethods develop during the twenty-firstcentury, the role of synthetic chemical pes-ticides in modern society will also change.

SummaryAs our planet’s population continues to

grow and undeveloped countries such asChina industrialize, an increasing emphasiswill be placed on using our nonrenewableresources wisely. The planet’s land, water,and atmosphere can be considered giantreaction vessels where humans conductlarge-scale chemical reactions. In manycases, we know what the results of the reac-tions will be, but in other instances theresults of human activities are unclear. Howwill an increasing CO2 concentration in theatmosphere affect global climate? What arethe long-term effects of synthetic organicchemicals introduced in the last half cen-tury? Can nuclear energy supply our energy

needs safely? These are just a few of themany questions in which knowledge ofenvironmental chemistry is critical fordetermining the answer. Even more impor-tant are those areas for which we do noteven know a question exists. While we tryto anticipate problems and consequenceswith new discoveries and products, we cannever fully anticipate how nature willrespond to our actions. Furthermore, theresponse from nature is not always imme-diate, but delayed. It is the role of environ-mental scientists to provide policy makersand the public with their best scientific per-spective on issues. In addition to informingthe public on what is known, it is alsoimportant to tell what is not known. The sci-ence of most environmental issues is“fuzzy,” and environmental scientists mustcontinually strive to bring the science ofthese issues into focus. Human-inducedchanges to the environment must be sepa-rated from natural changes, a task that isoften difficult. Even when the cause of anenvironmental problem is known, thecourse of action is often unclear. In additionto science, the nature of environmentalproblems involves economics, politics, andculture. While environmental scientistsmust provide their expertise, it is the role ofthe individuals to try to understand the fun-damental science behind environmentalissues and act accordingly. This chapterattempts to help you take a step in thatdirection.

IntroductionAll one has to do is take a quick look aroundto appreciate modern society’s dependenceon the chemical industry. The paper thesewords are printed on and the ink used toprint them are both products of modernchemical technology. In fact, we would behard pressed to find materials that are notrelated in some fashion to the chemicalindustry. Plastics, medicines, paints, textiles,explosives, fertilizers, cosmetics, fuels,detergents, glass, electroplating, metallurgy,and photography are examples of productsand processes that illustrate the impact ofthe chemical industry on our daily lives. Thegrowth of the modern chemical industrycorresponds to the development of chem-istry, commencing in the late eighteenthcentury. The growth of chemical industriesduring the 1800s was part of the industrialrevolution. While Great Britain and theUnited States led the industrial revolution inareas of agriculture and textiles, it was inGermany where the chemical industryachieved prominence. This situation per-sisted until World War I, when by necessity,the United States, Great Britain, and other

countries at war with Germany were forcedto develop their own chemical industries.

Traditionally, chemical technology wasmore art than science. Processes to producechemicals and products were crafts thatoften involved techniques passed betweenfamily members and generations. Physi-cians, alchemists, and apothecaries mixedsmall quantities of chemicals available formedicines and used as ingredients for prod-ucts such as soaps, candles, perfumes,foods, and beverages. Regions located closeto specific raw materials developed indus-tries that capitalized on their geographiclocation; for example, coal areas gave raiseto the steel industries around Pittsburgh andthe Ruhr Valley in Germany. As raw materi-als became depleted and demand for prod-ucts increased, chemical technologies wereapplied to meet demand. Chemical indus-tries developed to produce substances inlarge batch quantities that were then mar-keted internationally.

The development of chemistry as a mod-ern science provided a theoretical basis forchemical technology, and knowledge gainedin basic chemistry could be applied to meetsocietal needs. This tradition continues today

19

The Chemical Industry

290 The Chemical Industry

with the modern chemical industry providinga primary source of research funds. All of thelarge chemical companies have departmentsof research and development dedicated toimproving current technology and bringingnew discoveries to market. In this chapter, wetrace the history of the modern chemicalindustry. Additionally, we look at the basicchemistry used by industry to produce a fewof the materials that are part of daily life.

Acids and AlkalisThe first large chemical industries that

developed in modern times involved theproduction of acids and alkalis. The mostimportant industrial chemical used through-out history is sulfuric acid. Each year sulfu-ric acid tops the list of chemicals used byindustry, and it is often said that a country’seconomic status can be gauged by theamount of sulfuric acid it consumes in ayear. In ancient times sulfuric acid was pro-duced by heating the ore green vitiriol,FeSO4 •7H2O:

FeSO4•7H2O H2SO4(l) � FeO(s) � 6H2O(g)

Sulfuric acid produced in this manner wascalled oil of vitriol. During the early 1700s,sulfuric acid was produced in glass jars ofseveral liters using sulfur and potassiumnitrate, KNO3. In 1746, John Roebuck(1718–1794) substituted large lead-linedchambers for glass jars and was able toincrease the amount of sulfuric acid pro-duced from a few liters or pounds at a timeto tons. The lead-chamber process was usedthroughout the 1800s to produce sulfuricacid. The reactions representing this processare

While the lead-chamber process increasedthe amount of sulfuric acid that could beproduced, it relied on a source of nitrate thatusually had to be imported. The process alsoproduced nitric oxide gas, NO, which oxi-dized to brown nitrogen dioxide in the atmo-sphere. To reduce the supply of nitraterequired and the amount of nitric oxide pro-duced, Gay-Lussac proposed that the nitricoxide be captured in a tower and recycledinto the lead chamber. Although Gay-Lussac first proposed this modification tothe lead-chamber method around 1830, itwas not until the 1860s that John Glover(1801–1872) actually implemented Gay-Lussac’s idea with the Glover tower.

The lead-chamber process supplied theworld’s need for sulfuric acid for a centuryand a half. In the late nineteenth century, thecontact process replaced the lead-chamberprocess. The contact process utilized sulfurdioxide, SO2, which was produced as a by-product when sulfur-bearing ores weresmelted. The contact process was namedbecause the conversion of sulfur dioxide tosulfur trioxide, SO3, takes place on “con-tact” with a vanadium or platinum catalystduring the series of reactions:

The production of sulfuric acid was suf-ficient to meet world demand in the mid-eighteenth century. Sulfuric acid was usedin producing dyes, bleaching wools and tex-tiles, and refining metals. Its demand greatlyincreased at the end of the eighteenth cen-tury when a method was discovered forpreparing sodium carbonate, also known assoda ash or soda, using H2SO4.

Sodium carbonate, Na2CO3, and potas-sium carbonate (potash), K2CO3, were two

heat

6KNO � 7S rÆ 3K S � 6NO � 4SO3(s) (s) 2 (g) 3(g)

SO � H O r H SO3(g) 2 (l) 2 4(aq)

heat

S O r SO(s) 2(g) 2(g)

heat, catalyst

2SO O 2SO2(g) 2(g) 3(g)

SO H O H SO3(g) 2 (l) 2 4(l)r

r——

The Chemical Industry 291

common alkalis used for making soap,glass, and gunpowder. Throughout historythese alkalis were obtained from naturalsources. Natron imported from Egyptianlakes was one source of soda ash. It was alsoproduced by burning wood and leaching theashes with water to obtain a solution thatyielded soda ash when the water was boiledoff. In a similar manner, potash (the namecomes from the process used to produce a“pot of ashes”) was produced. Potash wasused in the manufacture of saltpeter (potas-sium nitrate), which was a principal compo-nent of gunpowder; potash was viewed as acritical chemical by national governments.Another source of soda ash was the barillaplant. The scientific name of this plant isSalsola soda, but it goes by the commonnames of sodawort or glasswort because thesoda produced from it was used in makingglass. Barilla is a common plant found insaline waters along the Mediterranean Seain Spain and Italy. Brown kelp from thegenus Fucus found off Scotland’s North Seacoast was also a source of potash. Barillaand brown kelp were dried and burned toproduce alkali salts. The depletion of Euro-pean forests and international disputes madethe availability of alkali salts increasinglyuncertain during the latter part of the eigh-teenth century. This prompted the FrenchAcademy of Science to offer a reward toanyone who could find a method to producesoda ash from common salt, NaCl. Some ofthe leading chemists of the day includingScheele and Guyton sought a solution to thesoda ash problem, but it was NicholasLeBlanc (1743–1806) who was creditedwith solving the problem. LeBlanc proposeda procedure in 1783, and a plant based onLeBlanc’s method was opened in 1791.Unfortunately, LeBlanc’s association withFrench royalty led to the confiscation of theplant at the time of the French Revolution.Furthermore, conflicting claims for

LeBlanc’s method were made by severalother chemists, and he never received thereward. LeBlanc, disheartened and destitute,committed suicide in 1806.

LeBlanc’s method uses sulfuric acid andcommon salt to produce sodium sulfate,Na2SO4. Sodium sulfate is then reacted withcharcoal and lime to produce sodium car-bonate and calcium sulfide:

H2SO4(l) � 2NaCl(aq) Na2SO4(s) � 2HCl(g)

Na2SO4(s) � 2C(s) � CaCO3(s) Na2CO3(s)

+ CaS(s)

The sodium carbonate and calcium sulfidewere separated by mixing with water.Because sodium carbonate was soluble inthe water and calcium sulfide insoluble, theformer would be suspended in solution.

LeBlanc’s process increased thedemand for sulfuric acid, and the alkali andacid industries were the first large-scalechemical industries. Plants using theLeBlanc process were situated in areas asso-ciated with salt mines, and this naturally cre-ated locales for other industries thatdepended on soda ash. The alkali industryusing the LeBlanc process created environ-mental problems near the alkali plants. Thehydrogen chloride gas killed vegetation inthe immediate vicinity of the plants. Todecrease air pollution, the gas was dissolvedin water, creating hydrochloric acid that wasthen discharged to streams, but this justturned the air pollution problem into a waterpollution problem. Another problem wascreated by the solid calcium sulfide product.Calcium sulfide tailings stored around alkaliplants reacted with air and water, creatingnoxious substances such as sulfur dioxideand hydrogen sulfide. Landowners adjacentto alkali plants sought relief from the envi-ronmental damage resulting from soda pro-duction. The situation got so bad in Englandthat the Parliament’s House of Lords


enacted the first of the “Alkali Acts” in1863. These laws regulated the productionof soda ash and required producers toreduce their environmental impacts on thesurrounding countryside.

The LeBlanc process was the principalmethod of producing soda ash until 1860when the Belgian Ernest Solvay(1838–1922) developed the process thatbears his name. The Solvay process, some-times called the ammonia method of sodaproduction, utilized ammonia, NH3, carbondioxide, and salt to produce sodium bicar-bonate (baking soda), NaHCO3. Sodiumbicarbonate was then heated to give sodaash. The series of reactions representing theSolvay process are

ExplosivesDuring the beginning of the nineteenth

century, the alkali and acid industries pro-vided the model for other chemical indus-tries. One characteristic of the chemicalindustry is that development in one areaoften stimulates development in anotherarea. For example, the lead-chamber methodproduced enough sulfuric acid to make theacid practical for use in the LeBlanc pro-cess. Similarly, the Solvay process usedammonia produced when coke was madefor steel production. Certain chemicalindustries were perceived by royalty andnational leaders as critical to their nation’swelfare. One of these was the manufactureof gunpowder, also known as blackpowder.Gunpowder is a mixture of approximately

75% potassium nitrate (saltpeter), 15%charcoal, and 10% sulfur. Gunpowder wasinvented by the Chinese in the ninth centuryand introduced into Europe around the thir-teenth century. Gunpowder mills wereestablished in Europe soon after its intro-duction, and they were heavily controlled byEuropean governments. Lavoisier madeextensive studies of the saltpeter (potassiumnitrate, KNO3) used in French gunpowdersat the Royal Arsenal in Paris; as a result,French gunpowder improved to the highestquality in Europe.

One of Lavoisier’s apprentices at thearsenal was a sixteen-year-old namedEleuthère Irénée DuPont de Nemours(1772–1834). DuPont came from an influen-tial family associated with the French roy-alty. DuPont’s father, Pierre Samuel DuPont(1739–1817), who attached the “deNemours” to the DuPont family name to dis-tinguish them from numerous otherDuPonts, barely escaped being guillotinedwith Lavoisier. Pierre and his sons wereforced to flee to the United States in 1800.The gunpowder industry in the United Stateshad traditionally been a cottage industry; thefirst gunpowder mill did not appear in thecolonies until 1775. Eleuthère Irénée DuPontpurchased land along the Brandywine Rivernear Wilmington, Delaware, to establish anAmerican gunpowder mill. Using knowl-edge gained from working at the FrenchRoyal Arsenal, DuPont started his plant in1802. DuPont began selling gunpowder in1804, and his business quickly expanded. Bythe War of 1812, DuPont was the largestmanufacturer of gunpowder in the UnitedStates producing more than 200,000 poundsannually. One hundred years later the com-pany sold over 1.5 billion pounds of explo-sives to the Allied forces for World War I.DuPont remained a family business well intothe twentieth century. Eleuthère’s son HenryDuPont (1812–1889), a graduate of West

2NH � CO � H O r (NH ) CO3(g) 2(g) 2 (l) 4 2 3(aq)

(NH ) CO � CO � H O4 2 3(aq) 2(g) 2 (l)

r 2NH HCO3 3(aq)

(NH )HCO � NaCl4 3(aq) (aq)

r NaHCO � NH Cl3(s) 4 (aq)

heat

2NaHCO r Na CO � H O � CO3(g) 2 3(s) 2 (l) 2(g)


Point, ran the company from 1850–1889.Henry’s son Lamont DuPont (1831–1884)was a chemist who perfected a gunpowderformula using sodium nitrate in place ofpotassium nitrate to make an improved, moreeconomical gunpowder. Lamont was killedin a company explosion in 1884. By the endof the nineteenth century, the E.I. DuPont deNemours Company was experiencing finan-cial difficulties and several elder DuPontsmade a decision to sell the company. Theyoung Alfred DuPont and two cousins con-vinced the older DuPonts to let them run thecompany. The young DuPonts revitalized thecompany into the largest chemical companyin the United States. While explosives con-tinued to form the core of the business, thecompany increasingly expanded into otherareas during the twentieth century, produc-ing dyes, chemical coatings, synthetic fibers,pigments, and general chemicals. Addi-tionally, E.I. DuPont diversified into other businesses such as automotives (GeneralMotors), agriculture, and petroleum.

Gunpowder was the primary explosiveused for almost one thousand years. In1846, the Italian chemist Ascanio Sobrero(1812–1888) first prepared nitroglycerin,but it was twenty years before Alfred Nobel(1833–1896) developed its use commer-cially. Nobel was born in Stockholm, Swe-den, where his father, Immanuel Nobel(1801–1872), ran a heavy constructioncompany. When Alfred was four, hisfather’s company went bankrupt andImmanuel left for St. Petersburg, Russia, tostart over. Immanuel rebuilt a successfulbusiness in Russia, in part due to his abilityto develop and sell mines to the RussianNavy for use in the Crimean War. Alfredand the rest of his family joined his father inRussia when he was nine, and Alfredreceived an excellent education with privatetutors. He studied in the United States andParis where he met Sobrero. Nobel studied

chemistry, literature, and mechanical engi-neering as his father groomed Alfred to joinhim in the construction and defense indus-try. After another downturn in his fortune,Immanuel and his sons, Alfred and Emil,returned to Stockholm in 1859 to startanother business. It was at this time thatAlfred began to experiment with nitroglyc-erin, seeking a safe method for its use as anexplosive. Nobel mixed nitroglycerin withother substances searching for a safe way totransport it and make it less sensitive to heatand pressure. Several explosions at Nobel’slab, one of which killed Emil, prompted thecity of Stockholm to ban nitroglycerinresearch inside the city and forced Nobel tomove his studies to a barge in one of thecity’s lakes. By 1864, Nobel started to man-ufacture nitroglycerin, an explosive thatwas almost eight times as powerful as gun-powder on a weight basis. In 1864, Nobelmixed nitroglycerin with silica to producea paste that could be shaped into tube pro-ducing dynamite. He also perfected a blast-ing cap to detonate the nitroglycerin. Nobelreceived a patent for dynamite in 1867, andhis business expanded very rapidly as dyna-mite was increasingly used for constructionand defense purposes. While Nobel’s fameand fortune were based on his invention ofdynamite, he was a very able inventor andchemist. Over the years he received 355patents, including ones for synthetic rubberand artificial silk. His will requested thatthe bulk of his fortune, which approached10 million dollars, be used to fund annualprizes in the areas of chemistry, physiologyand medicine, physics, literature, andpeace.

Synthetic DyesBy the middle of the nineteenth century

the chemical industry involved using basicinorganic chemicals to supply basic needs in


bulk quantities, for example, soap, glass,and gunpowder. The birth of organic chem-istry with Wöhler’s synthesis of urea in 1828provided the promise that chemicals couldbe synthesized to replace those obtainedfrom natural sources. There were severalproblems associated with obtaining chemi-cals from natural sources to supply sub-stances such as medicines, dyes, flavorings,and cosmetics. One problem was that mostof these chemicals were available only inlimited supplies from specific geographiclocations. This meant there was no guaran-tee that chemicals would be available on aregular basis. Wars, weather conditions,trade barriers, and plant and animal diseasesdisrupted the supply of chemicals on a reg-ular basis. Even when available in plentifulsupplies, the costs of gathering, transport-ing, and refining chemicals made many ofthem expensive and unavailable to commoncitizens. Another problem was that the qual-ity of chemicals extracted from naturalsources varied depending on the source,genetic variability of species, local growingconditions, and other factors.

The first fine chemicals to be synthe-sized in large quantities were dyes. Through-out the ages, dyes were extracted fromplants, animals, and minerals. For example,the blue dye indigo was obtained from the

leaves of various species of the genusIndigofera found in India. Indigofera wasbrought to South Carolina, and indigo wasthe most widely used dye during the eigh-teenth and nineteenth centuries in associa-tion with the cotton industry. Madder wasanother tropical plant family that containedthe plant Rubia tinctorium used to obtain reddyes. The molecule responsible for the dye-ing characteristic is alizarin (Figure 19.1).The Madder family also contained coffeeand medicines such as quinine, the latter ofwhich played a key role in the developmentof synthetic dyes as seen shortly.

Wöhler’s discovery in 1828 that a natu-ral organic chemical could be synthesizedfrom inorganic chemicals motivated otherchemists to search for synthetic substitutesfor other natural products. In the years pre-ceding Wöhler’s discovery, Faraday had dis-covered benzene as a liquid residue from thedistillation of coal used to produce gas forlighting in 1825. Aniline, C6H5NH2 hadbeen prepared from distilling indigo plants(1826). Other developments in Europe alsoset the stage for the synthetic dye industry.As the eighteenth century progressed, Euro-pean forests were being depleted, and cokewas increasingly substituted for charcoal asa source for carbon in iron smelting. Char-coal and coke are prepared by heating wood

Figure 19.1Natural and Synthetic Dyes


or coal in a low oxygen environment. Whencoal is heated to make coke, the residue iscalled coal-tar. By 1845, it was found thatcoal-tar consisted of a mixture of a numer-ous organic molecules including benzene,toluene, naphthalene, and xylene. With aready supply of coal-tar provided by the pro-duction of coke, methods were devised toisolate different organic compounds fromeach other. A leader in early coal-tar workwas August Wilhelm von Hofmann(1818–1892) who, although German, wasdirector of the Royal College of Chemistryin London from 1845 to 1864.

William Henry Perkins (1838–1907)enrolled in the Royal College at age 15, andby the time he was 18, he was a laboratoryassistant to Hofmann. Hofmann assignedPerkins the problem of synthesizing quinine,C20H24N2O2 from aniline. Quinine is analkaloid that was commonly used to treatmalaria. It reduced the fever of those suffer-ing from malaria, but it did not provide acomplete cure. Quinine’s natural source isthe bark and roots of cinchonas trees foundin the Andean regions of Bolivia and Peru.Jesuits introduced quinine to Europeanswho called the drug “Jesuit’s powder.” Dur-ing the mid-nineteenth century, the BritishEmpire had colonies surrounding the globe.Many of these colonies were in tropicalareas where malaria was prevalent. Becausethe natural supply of quinine was insuffi-cient to meet its demand for medicinal use,a synthetic source would provide bothmedicinal and commercial benefits. Perkinsattempted to synthesize quinine by oxidiz-ing allyltoluidine, C10H13N:

2C10H13N � 3O C20H24N2O2 � H2O

Perkins did not obtain the white crystals that characterize quinine, but instead got areddish-black precipitate. Perkins decided torepeat the experiment replacing allyltolui-

dine with aniline. When he oxidized anilinewith potassium dichromate, K2Cr2O7 he gota brown precipitate. Upon rinsing out thecontainer containing this residue, he pro-duced a purple substance that turned out tobe the first synthetic dye. Instead of quinine,Perkins had produced aniline purple. Ratherthan continue his studies at the Royal Col-lege, to Hofmann’s displeasure, Perkinsdecided to leave school and commercializehis accidental discovery. Backed by hisfather and a brother, Perkins established adye plant in London and obtained a patentfor aniline purple or mauve. It took severalyears for Perkins to perfect the manufactureof mauve, and in the process, he had to con-vince the textile industry of the advantagesof his synthetic dye. Mauve first foundwidespread use in France, but by 1862Queen Victoria was wearing a dress dyedwith mauve. Perkins went on to producenumerous other synthetic dyes at his factory,including alizarin. The synthetic productionof alizarin, a red dye, ruined the farming ofmadder in the Provençe region of France.Other synthetic dyes ruined local economiesin areas that had produced natural dyes, forexample, indigo in parts of India. Perkinssold his dye factory when he was thirty-seven years old. By this time, he was inde-pendently wealthy and devoted the rest ofhis life to basic chemical research.

Perkins’ discovery of mauve and hiscommercial success should not hide the factthat Perkins worked under the German Hof-mann at the Royal College. As mentionedpreviously, during the last half of the nine-teenth century and up until World War I, Ger-many was the undisputed center for advancesin industrial chemistry. Several leadingchemists besides Hofmann conductedresearch on dyes and organic synthesisthroughout the country. Justus von Liebig(1803–1873) was the leader of the Germangroup that dominated the chemical industry


of this, the dye industry was directly respon-sible for the birth of the pharmaceuticalindustry. Louis Pasteur (1822–1895) estab-lished the science of microbiology, but hewas educated as a chemist at Paris’ ÉcoleNormale. Pasteur worked closely with Lau-rent and early in his career worked on crys-tals and chemicals associated with theFrench wine industry. His first significantdiscoveries dealt with optical isomers, theability of solutions to bend polarized lightin opposite directions. In 1854, Pasteur’sinterest in biological problems grew as hestudied the different organisms responsiblefor various chemicals produced in fermen-tation. Pasteur found that fermentation inwine, milk, butter, and other foods was theresult of microorganisms and applied thisidea to develop the germ theory of disease.Pasteur developed several vaccines includ-ing one for anthrax in livestock. A numberof prominent scientists used Pasteur’s ideasto advance the study of human disease usingdyes. Robert Koch (1843–1910) used ani-line dyes to selectively isolate bacterium forstudy. Paul Ehrlich (1854–1915), assistantto Koch, believed dyes could be used to killbacteria and prevent disease. For example,he found methylene blue was effective intreating malaria. Another discovery relatedto Pasteur’s work included Joseph Lister’s(1827–1912) use of phenol (carbolic acid,C6H5OH), a coal-tar derivative, as an anti-septic in 1865 to sterilize wounds and sur-gical instruments. The Bayer company,originally founded in 1861 to produce thedye fuchsine, acquired its reputation formanufacturing aspirin (see Chapter 13,“Aspirin”).

Germany’s dominance in the chemicalindustry ended with the start of World WarI in 1914. The war forced the Europeanallies to abandon many German sourcesand develop their own chemical industries.Furthermore, chemical companies in coun-

in the latter half of the nineteenth century.Liebig was professor of chemistry at Giessenwhere he established a research institutionthat provided a steady stream of well-educated chemists who advanced chemicaltheory and technology throughout the coun-try. Liebig’s educational philosophy wasbased on a heavy emphasis in experimenta-tion and chemical analysis. Hofmann was astudent at Giessen before being appointed tohead the newly created Royal Institute ofChemistry in London in 1845. Hofmannreturned to Berlin in 1865. Other illustriousGerman chemists educated at Giessenincluded Kekulé, Karl Fresenius (1818–1897), and Robert Bunsen (1811–1899).These individuals in turn established centersfor chemical research in other German cities,Fresenius in Wiesbaden, Bunsen in Heidel-berg, Hofmann in Berlin, and Kekulé inBonn. The German government and privateindustry saw research as the key to the devel-opment of the chemical industry. During thistime, several major chemical companies werefounded to produce synthetic dyes includingBASF (Badische Anilin and Soda Fabrik) in1861, Frederick Bayer & Co. (maker ofBayer aspirin) in 1863, and Hoechst in 1880.Across the German border in Basel, Switzer-land, Chemische Industrie Basel (CIBA) wasestablished in 1859. This company mergedwith the Geigy (founded in 1759) in 1970 toform Ciba-Geigy. While universities, privateindustries, and the German governmentworked cooperatively to develop their chem-ical industry, other countries did not see thebenefit of chemical research that did not haveimmediate financial benefits. Fortunately,many American and European chemists wereeducated in Germany and returned to theirnative countries bringing the German chem-ical philosophy with them.

Although synthetic dyes principallyfound use in the textile industry, dyes even-tually found use in treating disease. Because


were not truly synthetic. Charles Tophamwas searching for a suitable filament forlight bulbs in 1883 when he produced anitrocellulose fiber. Louis Marie HilaireBernigaut (1839–1924) applied Topham’snitrocellulose to make artificial silk in 1884.Bernigaut’s artificial silk was the first rayon.Rayon is a generic term that includes sev-eral cellulose-derived fibers produced bydifferent methods. Bernigaut’s rayon wasnitrocellulose in which cellulose from cot-ton was reacted with a mixture of sulfuricand nitric acid. A general reaction to repre-sent the nitration of cellulose is:

�C6H7O2(OH)3 n � HNO3

�C6H7O2(NO3)3 n � H2O

The sulfuric acid acts to take up the waterthat is formed in this reaction. After con-verting the cellulose into nitrocellulose, theliquid is removed from the solution, and thenitrocellulose can be forced through a spin-neret to produce fibers. Bernigaut’s rayonwas highly flammable, and he spent severalyears reducing the flammability of his nitro-cellulose rayon before starting commercialproduction in 1891. Rayon was first pro-duced commercially in the United States in1910. By this time, several other methods oftreating cellulose had been developed toreplace nitrocellulose rayon. One involveddissolving cellulose in a copper-ammoniumhydroxide solution and is called thecupraammonium process. The third method,known as the viscose process, involvedreacting cellulose that had been soaked in analkali solution with carbon disulfide, CS2, toproduce a cellulose xanthate solution calledviscose. The viscose process became themost widely accepted method for producingrayon and accounts for most of the currentrayon production.

Cellulose acetate is another form of cel-lulose that was produced in the late 1800s.

tries such as neutral Switzerland and theUnited States, which did not enter the waruntil 1917, were more than willing to meetthe demands of the Allied countries. Spe-cific chemical industries, notably thoseassociated with explosives, particularlybenefited from the war. World War I alsoforced Germany to develop specific tech-nologies such as the ammonia industry asnoted at the end of Chapter 12. Germany’sdefeat in WWI led to a loss of its monopolyof a number of chemicals. Before the war,Germany controlled over 90% of the syn-thetic dye industry, but by the end of thewar, it shared this industry with the UnitedStates and Switzerland. For war reparations,German patents, trademarks, and holdingsin foreign countries were relinquished tothe Allied powers. By the end of WWI, thechemical industry was still dominated by alimited number of industries: syntheticdyes, acid-alkali, fertilizer, and explosives.The war set the stage for a new period inindustrial chemistry. International allianceswere built between corporations to searchfor new products and markets using syn-thetic petrochemicals as the building blocksof the chemical industry.

Synthetic Fibers and PlasticsThroughout human history a limited

number of fibers provided the fabric usedfor clothing and other materials—wool,leather, cotton, flax, and silk. As early as1664, Robert Hooke speculated that pro-duction of artificial silk was possible, but ittook another two hundred years before syn-thetic fibers were produced. The productionof synthetic fibers took place in two stages.The first stage, started in the last decades ofthe nineteenth century, involved chemicalformulations employing cellulose as a rawmaterial. Because the cellulose used in thesefibers came from cotton or wood, the fibers


ting plastic hardens into its final shape uponheating, whereas thermoplastics can berepeatedly heated and reformed into variousshapes. Baekeland established the GeneralBakelite Company in 1910.

By the 1920s, many companies wereproducing cellulose-based synthetic materi-als, and the stage had been set for the pro-duction of truly synthetic materials. By thistime DuPont had diversified its interest fromgunpowder and munitions production, whichutilized large quantities of nitrocellulose, intoa comprehensive chemical company. In 1920,DuPont started to produce rayon, and at thistime DuPont started to invest heavily inresearch. In 1928, DuPont hired the Harvardorganic chemist Wallace Hume Carothers(1896–1937), whose specialty was polymers,to lead a team of highly trained chemists inbasic research. Carothers’ team quicklystarted to develop commercially viable prod-ucts such as the synthetic rubber neoprene in1931. During the early 1930s the DuPontteam produced a number of different fibers,but these had specific problems such as lack-ing heat resistance. In February 1935, a fiberknown in the lab as “fiber 66” was producedthat held promise for commercialization.Fiber 66, also called nylon 66, was the firstnylon produced. Like rayon, nylon is ageneric term used for a group of syntheticallyproduced polyamides. The “66” refers to thenumber of carbon atoms in the reactants usedto produce it. In the case of nylon 66, the sixrefers to six carbons in adipic acid,HOOC(CH2)4COOH and six carbons in hex-amethylenediamine, H2N(CH2)6NH2. Nylon66 is produced when these two reactants arecombined under the proper conditions:

The basic cellulose unit contains threehydroxyl groups. The triester cellulose triac-etate forms when cellulose is reacted withglacial acetic acid. Hydrolysis removes someof the acetate groups to form a secondaryester, which averages about 2.4 acetyl groupsper unit rather than three. The secondaryester is then dissolved in acetone and thesolution ejected through a spinneret to formfibers. Cellulose acetate processed in thismanner is referred to as acetate rayon, but itmay be more commonly known by its tradename Celanese.

The development of plastics accompa-nied synthetic fibers. The first syntheticplastic with the trade name Celluloid wasmade in 1870 from a form of nitrocellulosecalled pyroxylin, the same substance used toproduce the first rayon. Celluloid was devel-oped in part to meet the demand for expen-sive billiard balls, which at the end of the nineteenth century were produced fromivory obtained from elephant tusks. JohnWesley Hyatt (1837–1920) combined pyr-oxylin with ether and alcohol to produce ahard substance called collodion. Hyatt’s col-lodion, like Bernigaut’s original rayon, wasunstable and potentially explosive. Hesolved this problem by adding camphor tothe collodion to produce a stable hard plas-tic he called Celluloid.

Celluloid, like rayon, was derived fromcellulose, and therefore, was not a truly syn-thetic material. The first completely syn-thetic plastic was Bakelite. This materialwas produced in 1906 by the Belgian-born(he immigrated to the United States in 1889)chemist Leo Hendrik Baekeland (1863–1944). Baekeland had made a small fortuneselling photographic paper to George East-man (1854–1932). Using this money,Baekeland studied resins produced fromphenol and formaldehyde by placing thesematerials in an autoclave and subjectingthem to heat and pressure. Bakelite was athermosetting phenol plastic. A thermoset-


was Velcro. The Swiss inventor George deMestral (1907–1990) received a patent forVelcro in 1955 after several years of trial-and-error experimentation. Mestral’s ideafor Velcro came to him while picking burrsoff his clothing and his dog’s fur afterreturning from a hike. He came up with theword “Velcro” from the French terms“velour” (velvet) and “crochet” (hooks).

Synthetic RubberRubber is another natural substance that

chemists synthesized during the twentiethcentury. Spanish explorers discovered nativeCentral and South Americans using rubberin the sixteenth century for waterproofing,balls, and bindings. Natural rubber comesfrom latex obtained from numerous plants,but the most important of these is Heveabrasiliensis. Europeans were intrigued withrubber products brought back by explorers,and the substance remained a curiosity fortwo centuries. Priestley noted the substancecould be used as an eraser if writing wasrubbed with it and so the name “rubber”came into use. Rubber factories developedin France and England at the start of thenineteenth century, but it was not until theend of the century that an increased demandled to an expansion of the industry. At thistime, rubber tree seedlings (from seedssmuggled out of Brazil) were cultivated inEuropean herbariums and transplanted tothe tropical colonies of various countries.Rubber plantations developed in a numberof regions, especially in the countries cur-rently known as Malaysia, Indonesia, SriLanka, and Thailand. Concurrent with therise of European plantations in SoutheastAsia, the United States obtained rubberfrom natural sources in Central and SouthAmerica. The rise of the auto industry andan increased demand for pneumatic tiresaccelerated the demand for rubber duringthe twentieth century.

The adipic acid used to produce nylon is theproduct of several reactions starting withbenzene. The adipic acid in turn is used toproduce hexamethylenediamine.

During the years 1936 to 1939, DuPontdeveloped the production methods neces-sary to market nylon. Unfortunately,Carothers never lived to reap the rewards asthe inventor of nylon. He committed suicidein April of 1937. Nylon’s first widespreaduse was as a replacement for silk in women’shosiery, but by 1941 nylon was being usedin neckties, toothbrushes, thread, and somegarments. During World War II, the UnitedStates government requisitioned the pro-duction of nylon solely for the war effort.Nylon replaced silk in military items suchas parachutes, tents, rope, and tires. Afterthe war ended, nylon’s use in civilian prod-ucts resumed. The demand for nylon stock-ings, which DuPont had created before thewar, resumed. In addition to its use as ahosiery fabric, it was used in upholstery,carpet, and clothing.

In addition to introducing nylon to thepublic in 1939, DuPont introduced numer-ous other synthetic fabrics includingacrylics (Orlon, introduced in the year1950), Dacron (a polyester marketed start-ing in 1959), Spandex (1959), and Kevlar(1971). Other companies developed count-less other fabrics and blends consisting ofrayon (1910), polyester (1953), andpolypropylene (1961). Gore-Tex, a polyte-trafluoroethylene (PTFE) material, wasdeveloped by Wilbert Gore (1912–1986)and his wife Genevieve in their basement.Gore was a DuPont engineer who started hisbusiness in 1958 after failing to persuadeDuPont executives to pursue PTFE fabrics.Gore worked on developing Teflon, which isanother PTFE product, at DuPont. TheGores mortgaged their home to start theirbusiness, and Gore-Tex outdoor clothingbecame popular in the late 1980s. Anothermaterial invented by individual initiative


reactions with acetylene at Notre DameUniversity. In one series of experiments,Nieuwland passed acetylene into a solutionof copper (l) chloride and ammonium chlo-ride to produce several polymers, one ofthese was monovinylacetylene. It was thiscompound that the chemists at DuPont wereable to convert into chloroprene to makeDuprene:

The double bonds in isoprene and chloro-prene allow these compounds to be vulcan-ized. In vulcanization, sulfur attaches to thedoubly bonded carbon to produce cross-linked chains:

Other producers were successful in linkingsingle polymers or two polymers, calledcopolymers, to produce synthetic rubbers. InGermany, a copolymer of butadiene(CH2CHCHCH2) and acrylonitrile(CH2CHCN) was produced by using asodium catalyst. It was called Buna, whichcame from combining “bu” for butadieneand “Na” for sodium. During the years lead-ing up to World War II, the United States,realizing how dependent it had become onforeign sources of rubber, continuedresearch on synthetic rubbers. The war andJapan’s control of the rubber-producing areasin Southeast Asia cut the United States’ nat-

Natural rubber tends to coagulatequickly into a hard tacky substance. To pro-duce different rubbers requires that variousadditives be combined with the latexobtained from rubber plants. Adding ammo-nia to the latex prevents coagulation andallows the latex to be shipped and processedas a liquid. Charles Goodyear (1800–1860)was attempting to improve the quality ofrubber in 1839 when he accidentallydropped a rubber-sulfur mixture on a hotstove. He discovered the product had supe-rior qualities compared to other natural rub-bers of his day. The process Goodyeardiscovered was vulcanization. Vulcanizedrubber is more elastic, stronger, and moreresistant to light and chemical exposure.Goodyear never reaped the rewards of hisdiscovery and died in poverty. Half a cen-tury later the Goodyear Tire Company wasnamed after him and is currently the largesttire producer in the United States and thirdin the world behind Michelin (France) andBridgestone (Japan).

Synthetic rubbers include a variety ofcompounds that mimic the properties of nat-ural rubber. In 1860, Charles GrevilleWilliams (1829–1910) isolated themonomer isoprene from rubber, showingthat rubber was a polymer of isoprene:

Numerous attempts were made to synthe-size isoprene rubber, but the first successfulsynthetic rubber was produced byCarothers’ group at DuPont. Dupont pro-duced Neoprene (DuPont’s name wasDuprene) from chloroprene in 1930.DuPont’s success was a result of work ini-tially performed by Father Julius Nieuwland(1878–1936), who conducted research on

H C3

|H C—C�C�CH2 2

|S|

H C—C�C�CH2 2

|CH3

Vulcanization

Cl|

H C�CH—C�CH2 2

Chloroprene


formed by the incomplete combustion ofnatural gas. Carbon black consists of fineparticles of carbon and was used in the early1900s for ink, pigments, and tires. The firstoil company to produce petrochemicals ona commercial basis was Standard Oil ofNew Jersey (Exxon). Standard of New Jer-sey started to produce commercial quanti-ties of isopropyl alcohol from petroleum in1920. The petrochemical industry grewsteadily during the twentieth century as oilcompanies developed chemical processesfor separating and converting petrochemi-cals from crude oil. Following the lead ofStandard of New Jersey, other oil companiesdeveloped chemical divisions and startedresearch and design (R & D) departments toprocess petrochemicals. Additionally, oilcompanies acquired or merged with estab-lished chemical companies to boost theirprofitability. This trend has continued to themodern day as witnessed by such recentmergers as Exxon-Mobil, Phillips-Atlantic-Richfield (ARCO), and the acquisition ofConoco by DuPont. A number of the largestchemical companies in the world (based onsales) are names traditionally associatedwith gasoline: Exxon, Chevron, Amoco,Shell, Phillips, Ashland, and BritishPetroleum (Table 19.1). It should be notedthat the figures in Table 19.1 are based on astrict definition of chemical sales. Chemi-cals such as gasoline, agricultural products,and pharmaceuticals, although chemicals,are not defined as chemical sales.

A relatively small number of chemicalsform the basis of the petrochemical indus-try. These are methane, ethylene, propylene,butylenes, benzene, toluene, and xylenes.These chemicals are used to derive thou-sands of other chemicals that are used toproduce countless products. Figure 19.2 listssome of the principal chemicals and prod-ucts derived from these seven basic chemi-cals.

ural rubber supply by over 90%. In response,several major companies including StandardOil, Dow, U.S. Rubber, Goodrich, Goodyear,and Firestone signed an agreement to shareresearch to boost synthetic rubber produc-tion. The cooperative agreement resulted inimproved qualities and grades of rubber andaccelerated growth in the synthetic rubberindustry. The United States developed a rub-ber called GR-S (government rubber-styrene) that was a copolymer made frombutadiene and styrene. Once the war ended,the development of synthetic rubbers haddisplaced much of the United States’ depen-dence on natural rubber. Plants controlled bythe government to support the war effortretooled to serve civilian needs in trans-portation, electrical, and consumer products.

PetrochemicalsThe use of petrochemicals was intro-

duced in Chapter 15. The modern chemicalindustry is largely based on the use of petro-chemicals, that is, chemicals derived frompetroleum and natural gas. Oil, natural gas,and coal have traditionally been used asfuels to provide energy. Coal tar providedthe bulk of organic chemicals used forindustrial products during the nineteenthcentury. The production of commercialquantities of oil from wells drilled in west-ern Pennsylvania at the end of the centuryprovided an alternative source of hydrocar-bons for synthesizing organic chemicals.Initially, oil production provided kerosenefor illumination and oil for heating. Thedevelopment of the automobile in the twen-tieth century shifted the emphasis of oil pro-duction toward gasoline. Oil refined toproduce gasoline also produced a host ofother chemicals that could be used as feed-stocks for other products.

The first petrochemical produced inlarge quantities was carbon black, which is


bromides, chlorides, and sodium hydroxide.Bromides were used in medicines, asbleaching agents, and in photographicchemicals. One of Dow’s first customerswas Kodak. For several years Dow metvarying degrees of success both scientifi-cally and financially. He began MidlandChemical Company in 1890, returned to theCleveland area and formed Dow ProcessCompany in 1895, and finally returned toMidland for good in 1897 where he foundedDow Chemical Company. One obstacleDow had to overcome to grow his businesswas fierce competition from the Germanand English bromide producers. These pro-ducers resented Dow marketing his bromidein Europe. The Germans slashed prices inAmerica in an attempt to drive Dow out ofbusiness, but Dow shrewdly and secretlymanaged to purchase the German bromideand resell it at a profit in Europe whereprices were not cut. Dow won the bromide

Industry HistoriesThroughout this chapter and in previous

chapters, brief histories of several well-known companies started were presented—DuPont, Alcoa, Bayer, and BASF. In thisfinal section on the chemical industry, thehistories of several well-known companiesusing chemical technologies are presented.

DuPont, presented at the beginning ofthis chapter, and Dow are the largest chem-ical companies in the United States. DowChemical Company was started by HerbertHenry Dow (1866–1930) in Canton, Ohio.Dow was a student at Case Institute inCleveland who studied the characteristics ofsalt brines acquired from wells around theGreat Lakes. Dow was determined to dis-cover methods to extract chemicals from thesalt brine. Rather than use the standard dis-tillation method of his day to obtain chemi-cals, Dow employed electrolysis to separate

Table 19.1Top Ten World Chemical Producers Based on Sales for Year 2000

Source: Chemical and Engineering News, July 23, 2001


a comprehensive chemical company pro-ducing hundred of different products.

Another familiar name in the chemicalindustry is Monsanto. Monsanto wasfounded in St. Louis as a producer of foodadditives in 1902 by John Francis Queeny(1859–1933). Queeny named the companyafter his wife’s maiden name. Its first prod-uct was saccharin. Saccharin was not pro-duced in the United States at that time andhad to be imported from Germany. Queenyand his associates sold saccharin to softdrink producers; Coca-Cola was one of his

war; he agreed not to sell in Germany, whilethe Germans would not sell in the UnitedStates, but bromide sales elsewhere wereallowed. Dow’s next assault came fromDuPont’s attempt to acquire the company in1915. Dow and his staff threatened to quit ifDuPont purchased Dow, and DuPont with-drew its offer. Dow was both an outstandingchemist and businessman. During his careerhe acquired nearly 100 patents. Heexpanded his electrolytic methods for sepa-rating ions from brine to seawater to pro-duce magnesium and iodine. Today Dow is

Figure 19.2A small sampling of some of the chemicals and products obtained from petrochemical sources


Air Products, and Electro Metallurgical toform the Union Carbide & Carbon Com-pany. In 1957, the company’s name wasshortened to Union Carbide. In addition tobatteries, some of its major products includePrestone antifreeze, Glad wrap, and Cham-pion spark plugs. Union Carbide’s repu-tation was severely hurt in 1984 when itspesticide plant in Bhopal, India, had an acci-dent resulting in the death of 2,500 Indians.In February of 2001, a merger betweenUnion Carbide and Dow was completed,making Union Carbide a subsidiary of Dow.

In the section on rubbers it was notedthat the Goodyear Tire and Rubber Com-pany was not founded by Charles Goodyear.Frank A. Seiberling (1859–1955) and hisbrother Charles (1861–1946) foundedGoodyear in 1898, thirty-eight years afterGoodyear’s death. Originally, the Seiberlingbrothers produced carriage and bicycle tires,but in 1901 they started producing automo-bile tires for the infant automobile industry.Although Goodyear built its reputation onrubber products, it expanded interests intotextiles, chemicals, and a number of petro-chemical products related to the rubberindustry. The B.F. Goodrich rubber com-pany was started by Benjamin FranklinGoodrich (1841–1888). Goodrich, alongwith several partners, founded a rubbercompany on the Hudson River in 1869.Goodrich, as president of the company,believed a location further west would bemore suitable for his company and exploredAkron, Ohio, as a possible site for reloca-tion. With backing from Akron businessinterests, Goodrich moved his operation toAkron, incorporating as Goodrich, Tew &Company on December 31, 1871 (Tew wasGoodrich’s brother-in-law). Goodrich’scompany went through several reorganiza-tions and grew slowly during its first decade.It was not until after Goodrich’s death thatthe company established itself as a viable

primary customers. Monsanto expandedproduction into caffeine, vanilla, aspirin,phenolphthalein, and several other chemi-cals over the next decade. World War Iaccelerated Monsanto’s growth, and itsteadily expanded its product line after thewar. Its many products include syntheticfibers, rubber, pesticides, detergents, andpharmaceuticals. More recently attemptshave been made to merge Monsanto withDow, which would create the largest chem-ical company in the United States.

Union Carbide grew out of companiesmaking carbon electrodes and acetylene forcity and home lighting. The National Car-bon Company founded in 1886 producedcarbon electrodes for street lamps. It devel-oped the first commercial dry cell batteriesand distinguished its batteries with theEveready trademark. Union Carbide wasfounded in 1898 by Canadian Thomas L.Willson (1860–1915) and James T. More-head (1840–1908). Morehead and Willsonwere using an electric arc furnace in anattempt to prepare aluminum. In 1892,while searching for a process to make alu-minum, Morehead and Willson combinedlime and coal tar in their furnace and pro-duced calcium carbide, CaC2. Acetylene,C2H2, is produced when calcium carbidereacts with water according to the followingreaction:

CaC2(s) � 2H2O(l) C2H2(g) � Ca(OH)2(aq)

Morehead started promoting the use ofacetylene for lighting, and calcium carbideplants were established in Sault Ste. Marie,Michigan, and Niagara Falls, New York.Morehead developed a high carbon fer-rochrome in 1897, which was used forarmor plating in the Spanish-American War.In 1917, Union Carbide merged with theNational Carbon Company, Prest-O-Lite(another calcium carbide producer), Linde


producer of bicycle and carriage tires.Goodrich’s most significant innovation dur-ing this period was the development of apneumatic tire suitable for the speeds andload of automobiles. Goodrich expandedinto the chemical business after World WarI and developed a method for producingpolyvinyl chloride. This development con-tinued, and in 1943 B.F. Goodrich opened achemical division producing plastics, vinylflooring, and chemicals. A third majorAmerican rubber and tire company wasfounded by Harvey Samuel Firestone(1868–1938). Firestone started FirestoneRubber Company while working for hisuncle’s carriage shop in Chicago in 1896. In1900, he relocated to Akron, Ohio, andfounded Firestone Tire & Rubber Company.Firestone built his company by supplyingtires to Ford Motor Company. He was aclose friend of Henry Ford (1863–1947)throughout his life. In 1988, Firestone Tireand Rubber was acquired by the Japanesefirm Bridgestone Tires.

As noted previously, Exxon Corpora-tion is one of the leading chemical compa-nies in the world. It regularly ranks amongthe top five world chemical producers.Exxon was one of a number of oil compa-nies formed from the breakup of John D.Rockefeller’s (1839–1937) Standard OilTrust. Rockefeller started his oil businesswith Samuel Andrews in 1863 in Clevelandand quickly developed into the major refinerof crude coming out of the fields of Penn-sylvania. In 1870, Rockefeller and his asso-ciates formed the Standard Oil Company.Rockefeller built an oil monopoly by obtain-ing rebates from the railroads and buyingout competitors. His practices were chal-lenged in state and federal courts forcingRockefeller to relocate from Ohio to NewJersey in 1892, where state laws were morefavorable to his attempts to form a trust andpreserve his monopoly. At the time, Rocke-

feller controlled more than 90% of the refin-ing and distribution of oil in the UnitedStates. In 1911 following years of litigation,the Standard Oil Trust was finally dissolvedunder the Sherman Antitrust Act. Standardof New Jersey became Exxon and continuesas a major force in the oil and petrochemi-cal industry. Other major petroleum compa-nies that resulted from Standard Oil includeAmoco (Standard of Indiana), SOHIO(Standard of Ohio and subsequently takenover by British Petroleum), Chevron (Stan-dard of California), and Mobil (Standard ofNew York). It is ironic that Standard Oiltrust, which was dissolved by the courts atthe beginning of the twentieth century, hasbeen somewhat rebuilt in recent years withmergers such as Exxon and Mobil.

Procter & Gamble is a familiar companythat has supplied household products forover 150 years. William Procter (1801–1884), an immigrant candle maker fromEngland, and James Gamble (1803–1891),an Irish soapmaker, met in Cincinnati whenthey married sisters. Procter and Gamblewere encouraged by their father-in-law tostart a candle- and soap-making businesstogether. They started their business in 1837in Cincinnati. At the time, Cincinnati was acenter for hog butchering, and the slaughter-ing houses provided an abundance of animalfat that was used to produce lye (NaOH), akey ingredient in both soap and candles.Procter & Gamble’s business grew steadily.The Ohio River and the railroads providedeasily accessible transportation to ship goodsboth east and west. Procter & Gamble was amajor supplier to the Union Army during theCivil War. During this period, due to theshortage of raw material created by the war,Procter & Gamble developed new methodsfor candle and soap making. In the mid-1870s, Procter & Gamble, under the direc-tion of Gamble’s son William, sought todevelop a low-cost high-quality soap to rival


expensive castile soaps made from olive oil.One early batch of their soap had inadver-tently been mixed excessively when a workerfailed to shut down a machine during thelunch break. Excessive air was introducedinto the batch, causing the hardened cakes tofloat. Rather than scrap the batch, the soapwas sold and customers demanded more ofthe floating soap. The soap’s ability to floatwas advantageous to people who oftenbathed in rivers and lakes at the time. If thesoap was dropped, it was easier to find thanother soaps that sunk when dropped. HarleyProcter, son of William, named the soapIvory inspired by Psalm 45 in the Bible:“The perfume of myrrh and aloes is on yourclothes; musicians entertain you in palacesdecorated with ivory.” Harley Procter orga-nized a large advertising campaign to pro-mote Ivory soap. As part of the advertisingcampaign, Procter & Gamble, which hadalways stressed quality control in its prod-ucts and used its moon and stars trademarkto distinguish P & G products from others,used independent chemical analyses of Ivorysoap, which showed 0.56% impurities.Hence, the slogan 99 44/

100% pure was born.

Soaps continued to be Procter & Gamble’smain product throughout the nineteenth cen-tury. The company pioneered the process ofhydrogenation of unsaturated fats and intro-duced Crisco to the public in 1912. In 1957,Procter & Gamble acquired Clorox, thelargest producer of household bleach in thecountry. Clorox was founded in 1913 by fivebusinessmen in Oakland, California. Theoriginal Clorox bleach solution, 21% sodiumhypochlorite (NaOCl), was delivered in five-gallon returnable buckets to dairies, brew-eries, and laundries that used the solution fordisinfecting equipment. A 5.25% solutionwas sold door-to-door and in local groceriesstarting in 1916, and Clorox bleach remainedthe company’s most sold product up until itwas acquired by Procter & Gamble.

Procter & Gamble’s number one com-petitor is another well-known producer ofhousehold items: Colgate-Palmolive. Wil-liam Colgate (1783–1857) was a candlemaker from New York City who opened hisbusiness in 1806. Colgate sold candles,starch, and soap out of his Manhattan shop.These products formed the core of Colgate’sbusiness for its first fifty years. In 1873, Col-gate marketed toothpaste in jars, and itintroduced the familiar tube dispenser in1896. Today it is the world’s largest producerof toothpaste. Colgate merged with The Pal-molive Company in 1928, forming Colgate-Palmolive. The Palmolive Company wasoriginally founded as B.J. Johnson SoapCompany in Milwaukee in 1864. In 1898,Johnson introduced Palmolive soap. Pal-molive soap was made with palm and oliveoil, hence the name Palmolive, rather thanwith animal fat and became so popular thatthe company eventually changed the com-pany’s name to that of the soap.

This section has focused on a few well-known chemical companies started in theUnited States. Remember that Germany ledthe world in industrial chemistry for half acentury before World War I forced othercountries to develop their own chemicalinfrastructure. Germany continues to be aleader in the modern chemical industry.BASF (Badische Analin und Soda Fabrik)and Bayer (of aspirin fame), traditionallytwo of the largest chemical companies in theworld, have exerted a major influence on theworld chemical industry since the mid-nineteenth century. Before World War II,BASF partnered with the two other Germangiants: Bayer and Hoechst, and with hun-dreds of smaller German companies to formthe chemical cartel known as IG Farben. IGFarben was a major supporter of AdolfHitler and instrumental in rearming Ger-many in the period between World War I andWorld War II. IG Farben factories were


heavily damaged during WWII. After thewar, many of the directors of IG Farbenwere tried as war criminals, and alliesargued about the fate of the German chem-ical cartel. It was eventually decided tobreak IG Farben into BASF, Bayer, andHoechst and nine smaller companies. WhileBASF and Bayer continue to be majorchemical producers, in 1998 Hoechstmerged with the French company Rhône-Poulenc to form the pharmaceutical andagrochemical company Aventis.

SummaryThis chapter has given a brief overview

of the chemical industry, with emphasis onchemical production in the United States.The importance of the chemical industry inthe United States cannot be overstated. TheUnited States is the largest producer ofchemicals in the world, accounting for 27%of the world’s production in the year 2000.The total value of chemicals produced inthe United States during 2000 was $460 bil-lion. Of this total, $80 billion was exportedand the United States’ trade surplus inchemicals was $6.3 billion. Thirty-one bil-lion dollars was spent on research anddesign in the year 2000. There are over12,000 chemical plants in the United States.One hundred and seventy U.S. chemicalcompanies operate an additional 3,000chemical plants in foreign countries. Thechemical industry employs over 1,000,000workers and accounts for 1.2% of our grossdomestic product. The top chemical pro-ducing states are Texas, New Jersey, Loui-siana, North Carolina, and Illinois.

In recent decades, much of the chemi-cal industry has been scrutinized withrespect to its products and practices.National stories on hazardous wastes gener-ated by the industry (Love Canal, TimesBeach) paint a picture of a toxic cornucopia.

The popular movie Erin Brockovich,although focusing on the power industry, isan example of how the popular media cansimplify the public’s perception of chemicalexposure. In Erin Brockovich the chemicalof concern was hexavalent chromium, Cr6�.Hexavalent chromium is listed as a carcino-gen by EPA, but complicated questionsdealing with exposure (drinking, inhalation,absorption through skin), toxicity levels,and specific health effects were lost in Hol-lywood’s version.

The modern chemical industry is ahighly regulated industry. National legisla-tion passed in the last few decades such asthe Resource Conservation and RecoveryAct (RCRA), Toxic Substances Control Act(TSCA), and Comprehensive EnvironmentResponse, Compensation, and Liability Act(Superfund) have required chemical com-panies to take responsibility for theirwastes. Furthermore, chemical industrieshave discovered that minimizing waste andbeing energy-efficient are sound businesspractices that increase profits. While acci-dents in the chemical industry will continueto happen and a small proportion of com-panies will try to “cut corners,” overall thechemical industry is safe and responsible.Society has become dependent on many ofthe products marketed by the chemicalindustry. Everyday life would be signifi-cantly changed without many of theseproducts.

This chapter has touched upon only afew of the core chemical industries:petroleum refining, rubber, plastics, andtextiles. A comprehensive review of thechemical industry would include agro-chemicals, glass, pharmaceuticals, paper,metals, photography, brewing, cement,ceramics, printing, and paints. As thetwenty-first century unfolds, the chemicalindustry will continue to play a dominantrole in dictating our quality of life. The core


tries. Additionally, the old guard will notstand still. The large chemical companies,which have developed over the last century,will continue to develop new productswhile at the same time will search the globefor acquisitions, partnerships, and mergersto increase their value.

industries such as petroleum, chlor-alkali,metal smelting, and rubber will continue toprovide the products of daily life. However,new industries based on biotechnology,newly created materials, and exotic pro-cesses (products created in space) willattract entrepreneurs to form new indus-

IntroductionScience begins with careful observationusing all of our senses. Most people associ-ate observation with the sense of vision, butthe senses of smell, taste, hearing, and touchare also used to observe. All five of thesesenses are important in working with chem-icals. Observations accompanying chemicalreactions are often made by smelling anodor. For example, a distinctive odor is pro-duced in reactions where ammonia isformed. Inadvertently coming into contactwith a strong acid or base causes a burningsensation. One way to characterize acids andbases is by taste; acids taste sour and basesbitter. Nearly all of our observations madeevery day are accepted at face value, but fre-quently there are observations that do not fitour expectations. These observations gainour attention. Many are quickly dismissed,but some tend to peak our curiosity. Wemight question whether we can believe oursenses and then attempt to verify an initialstrange observation. If we trust our senses,we might begin to ask questions to explainour observations. Many paths can be takenwhen trying to explain observations, but theone used by scientists is termed the scien-

tific method. In this chapter, we take a brieflook at the scientific method, discuss someelements of experimentation, and concludewith a few chemistry activities and experi-ments.

The Scientific MethodAs stated in the previous section, ques-

tions naturally arise from observations madein daily activities. Questions may also arisefrom something read, a discussion, or anexperience. There are various methods usedto seek the answers to questions. Askinganother person, consulting a book, orsearching the Internet is often sufficient.Scientific questions are questions to whichthe scientific method can be employed. Thescientific method is a systematic procedureused to answer questions in order to developlogical explanations for the world (universe)around us. The general scientific methodfollows a common pattern for answering aquestion. The procedure consists of devel-oping a question into a testable hypothesis,developing an experiment to test the hypoth-esis, performing the experiment, and usingthe results to come to some conclusionabout the original question. As described in

20

Chemistry Experiments

310 Chemistry Experiments

this fashion, the scientific method seemslike a straightforward procedure for tacklinga question. In reality, the scientific methodoften follows a circuitous path. Applying thescientific method is like driving in new loca-tions. In familiar locations, there is littleproblem of getting from one place toanother. Several routes may be used, butsteady progress is made in arriving at a des-tination. Conversely, driving in an unfamil-iar area, such as a large city, often consistsof wrong turns, backtracking, and runninginto dead ends. Even with a road map anddirections, the path taken is less than direct.In the rest of this section, some of the keyelements of the scientific method are brieflydiscussed.

When using the scientific method it isimportant to formulate a scientific questionabout a subject and that this question pointsthe researcher toward further knowledge onthe subject. Posing a question may seemtrivial, but great scientists are often charac-terized by having asked the right question.This often means looking past the obviousand focusing on questions that have notbeen asked. Alternatively, perhaps everyoneelse has ignored the obvious and overlookeda very basic question. Scientific questionsmust be clearly defined. Students startingprojects often use terms loosely in posingquestions. For example, a question such as“which battery is better?” is unclear. Ques-tions arise concerning what is meant by bet-ter (cheaper, longer life, more powerful),under what conditions, and the type of bat-tery. An alternate question might be “howdoes the life of rechargeable batteries on asingle charge compare to nonrechargeablealkaline batteries in audio equipment?” For-mulating an appropriate question leads theresearcher to developing a testable hypoth-esis.

A hypothesis is a possible explanationthat answers the question. It is often pre-

sented as a concise statement of what theoutcome of an experiment will be. Ahypothesis is often defined as being an edu-cated guess. Before a hypothesis is stated, asmuch background information about thesubject should be gathered. In fact, it maybe possible that an answer to the questionwill come from the preliminary research.Books, journal articles, the Internet, andexperts on the subjects can be used to guidethe researcher in refining the question anddeveloping a hypothesis. Researching thesubject allows the researcher to draw tenta-tive conclusions based on the work of oth-ers. This is why the hypothesis is only apossible explanation. Researching the sub-ject can help identify important variables informulating the hypothesis. Variables areconditions that change and may affect theresults of an experiment. For example, in astudy comparing the life of batteries it maybe important to know how batteries operateat different temperatures. If the purposewere to compare battery use of flashlightsused both indoors and outdoors, then tem-perature would be an important variable toconsider. This would affect how the hypoth-esis was formulated, and ultimately, how theexperiment was designed.

Once a hypothesis is formulated, anexperiment is conducted to test the hypothe-sis. Experimentation is what distinguisheschemistry, and other experimental science,from other disciplines. In simple experi-ments, the researcher designs the experimentto examine one independent variable andattempts to hold all other variables constant.Again, using the battery example, if theexperiment was designed to test battery lifein flashlights, then “battery life” would be thedependent variable. It would be im-portant todefine in exact terms the variable “batterylife” and exactly how it would be measured.The researcher would then attempt to controlall other variables, except for one indepen-

Chemistry Experiments 311

dent variable that is allowed to change,thereby conducting a controlled experiment.Temperature might be controlled by con-ducting the test in a controlled environment,and the type of flashlight could be controlledby using the same or identical flashlights.Although the researcher makes every attemptto control for every variable except the inde-pendent variable, it is impossible to controlfor all variables. In the flashlight example,even using the same flashlight means that oneset of batteries would be placed in a flashlightthat was slightly more used. On the otherhand, the researcher might use two identicalflashlights and conduct the experiment usingrechargeable batteries in one and alkaline bat-teries in the other during a first trial. Then asecond trial would be conducted with a newset of batteries and switching flashlights. Theresults of the two trials could then be aver-aged to control for order of batteries placedin flashlights. Although this controls fororder, the fact that two identical but differentflashlights were used means that the variable“not exactly the same flashlight” is intro-duced. There are methods that could be usedto control for order using a single flashlight,but the point of this example is to illustratehow difficult it is to totally control for vari-ables. Another variable important to controlin this experiment is “age of the batteries.” Ifthis was a student experiment and the studentrandomly picked batteries off the store shelfwithout considering the expiration date, thiscould seriously bring into question any con-clusions reached. Even when the expirationdate is used, this is no guarantee that the bat-teries are the same age.

This example also demonstrates why itis important to identify the important vari-ables when formulating the hypothesis anddesigning the experiment. A few variableswill generally be identified as important tocontrol; most variables will be ignored. Forexample, variables such as air pressure, rel-

ative humidity, position of flashlight duringthe experiment, color of batteries, and so on,would seemingly have no effect on batterylife. It is important to remember that anexperiment without proper controls is gen-erally useless in reaching any valid conclu-sions concerning the hypothesis.

In conducting an experiment, it isimportant to keep careful and accuraterecords. The standard method for this isusing a laboratory notebook. Official labo-ratory notebooks are sold at bookstores or anotebook of graph paper may be used. Asimple laboratory notebook can be made byplacing sheets of graph paper in any three-ring binder. The first several pages of thenotebook should be left blank to allow for atable of contents. Your name and phonenumber should be on the inside cover alongwith any emergency information. The latteris important in situations where an accidentmight occur and authorities may have to becontacted. The laboratory notebook is usedto write a comprehensive record of theexperiment. Entries should be dated andimportant relevant information on the exper-iment should be recorded. Information suchas reactions, notes, times, observations,questions, calculations, and so on should bewritten in the laboratory notebook. Concen-trations of solutions, forms of chemicals(examples include hydrated form or com-pany that produced), and quantities usedshould be recorded. By going back throughthe laboratory notebook, it should be possi-ble to retrace the steps taken during theexperiment. This is important in analyzingwhy things worked or did not work asexpected. It is also important to use the lab-oratory notebook to identify areas where theexperiment can be improved or modified.Results recorded in the laboratory notebookshould provide the basis for future experi-ments. The notebook should be writtenneatly in ink. Changes should be crossed out


with a single line, enabling the crossed-outmaterial to be read.

Once the experiment has been con-ducted and data collected and analyzed, aconclusion can be drawn concerning theoriginal hypothesis. The experiment doesnot establish absolute truth concerning thehypothesis. A researcher should resist thetemptation to state that the hypothesis wasproven or not proven. In science, absolutetruth can never be established. Results of anexperiment may either support or not sup-port the hypothesis. Ambiguous results cancause the researcher to modify the hypothe-sis and redesign the experiment. Even whenthe results are not ambiguous, an experi-ment is often repeated. One of the bench-marks of science is reproducibility. When anexperiment is conducted and certain con-clusions reached, another researcher shouldbe able to repeat the same experiment andobtain similar results.

When analyzing the results of an exper-iment, it is important to consider sources oferrors and how these sources affected theresults. Two primary types of errors occur inexperiments. These are random error andsystematic error. The researcher has no con-trol over random error. Random errorinvolves the variability inherent in the natu-ral world and in making any measurement.As its name implies, random error varies ina random manner. An attempt is made tocontrol for random error by taking multiplemeasurements. For example, say we wantedto know the mass of an average penny. Nat-urally, older pennies might be expected tohave smaller masses due to wear and newerpennies more mass. (We will ignore the factthat penny composition changed in 1982.) Ifa large random sample of pennies was used,it would be expected that there would be agood mixture of pennies. A large sample ofpennies should include pennies of manyages. When measured on a scale, some pen-

nies would have more than the true mass ofa penny, others less, but it is assumed thatthese would tend to balance out and give afairly accurate measure of a penny’s mass.

Systematic error is error that occurs inthe same direction each time. For example,if a scale was used that had a dirty pan, eachmeasurement would be increased by themass of dirt on the pan. Another example ofintroducing systematic error in an experi-ment might be reading the meniscus level ofa graduated cylinder at an angle. Lookingdown when measuring a liquid’s volumewould systematically underestimate the vol-ume. Systematic error can be eliminatedusing proper techniques, calibrating equip-ment, and employing standards. An exam-ple of accounting for systematic error ischecking a thermometer to see if it measures0°C in a freshwater-ice mixture and 100°Cin boiling freshwater at 1 atmosphere pres-sure. If both measurements were 2°C high,then it can be assumed that a systematicerror of �2°C exists in the temperaturemeasurements made with this thermometer.

Chemistry ExperimentationThe most important aspect of conduct-

ing chemistry activities and experiments issafety. The excitement of chemistry is oftenportrayed to the public through fire, explo-sions, and large noises. Both Jay Leno andDave Letterman regularly have scientistsperform attention-grabbing demonstrationson their shows. The caveat “don’t try this athome” is often spoken in jest, but should betaken seriously. Individuals performing pub-lic demonstrations have practiced them andhoned their skills in science classrooms andmuseums. They know what conditions(amount of reagents, safety precautions) areneeded to conduct the demonstration safely.Chemical demonstrations, like magic tricks,are often “hyped” for the audience. Never-


theless, if you are not thoroughly familiarwith a chemistry activity and its possiblehazards, do not attempt it.

Even when doing simple experiments itis important to practice safety. It is a goodhabit to wear safety glasses, avoid loose-fitting clothing, wash your hands after handling chemicals, dispose of chemicals properly, have extinguishers and other safetyequipment available, and when in doubt askquestions. The simple activities listed in thefollowing section can be performed safely athome. Before doing any activity or experi-ment written by someone else, it is importantto read and think through it before actuallyconducting it. Safety begins and ends withcommon sense. If a procedure is unclear,advice should be sought from an adult.

One of the difficulties in doing chem-istry at home is obtaining materials. Whilemany chemicals available in even a modestchemistry lab are not available at home, asurprising number of basic chemicals can beobtained at local groceries and pharmacies(see Table 5.6 for some common chemicals).Building supply stores, aquarium shops,garden supply stores, and hardware storescarry a number of chemicals. Substitutionsof commercial products can often be made,for example, nail polish remover for ace-tone. A small collection of inexpensiveinstruments may be assembled by searchingthe shelves of various stores, secondhandshops, and garage sales. Eyedroppers, tongs,tweezers, syringes, funnels, and measuringcups can be used without modification. Rub-ber or plastic tubing can be obtained fromhardware or automotive stores. Alcoholthermometers can be found for less than $2;they are usually encased in a housing fromwhich they can be removed. Mercury ther-mometers should not be used. A postagescale can substitute for a balance, a smallcamping stove for a Bunsen burner, Pyrexcookware for beakers, and a warming plate

or iron can be used to make a hot plate. If itis safe, gas or electric ranges, microwaveand regular ovens, and refrigerators may beused to conduct experiments.

Chemical ActivitiesThis section contains a few chemical

investigations that may be conducted athome. They are not full experiments, butthey should provide useful activities thatcould be developed into a full-fledgedexperiment. Each illustrates basic conceptsexplained in the previous chapters. Mostinvolve a few simple steps, while a few aremore complicated. Again, the most impor-tant aspects of these activities are to be safe,have fun, and learn. Before actually tryingto do the activity, read and understand theactivity. Assemble all materials before start-ing and do the activity in a safe place. Ifchemicals are stored, label and mark thecontents properly. Pay close attention tosafety and dispose of used chemicals andmaterials safely.

Quantities in this section are givenusing cooking measures. Although metricunits are preferred in science, most mea-suring units found in the home consist oftraditional cooking units such as cup andteaspoon. For this reason, these units areused. Approximate equivalent units aregiven in Table 20.1.

Table 20.1Approximate Metric and EnglishEquivalents


Activity 1 Dissolving an EggshellMaterials: egg, vinegar, jar or containerIn this activity, an eggshell is dissolved

with vinegar, illustrating a simple chemicalreaction. To dissolve the eggshell, fill a jaror other suitable container with enoughvinegar to cover an egg. Place an egg in thevinegar. Vinegar is diluted acetic acid,HC2H3O2. In two or three days, the shell willbe gone. As the egg dissolves, bubbles willbe observed. These bubbles are carbon diox-ide, CO2. Handle the egg carefully after theshell is gone. The egg’s membrane remainswith the albumen and yolk inside. Aceticacid reacts with the calcium carbonate in theeggshell. The reaction can be represented as

CaCO3(s) � 2HC2H3O2(aq)

Ca(C2H3O2)2(aq) � H2O(l) � CO2(g)

Dissolving an eggshell in vinegardemonstrates a modern environmental prob-lem. Calcium carbonate is present in Earth’scrust as marble, limestone, and chalk. Manymodern buildings, statues, and stone struc-tures contain calcium. Acid rain is slowlydissolving these structures.

Do not throw away the egg; it can beused for the next activity.

Activity 2 OsmosisMaterials: egg with shell removed from

Activity 1, corn syrup, water, food coloring,drinking glass or other container to holdegg, spoon

Osmosis takes place when a solventmoves across a semipermeable membrane.The solvent moves from an area of high sol-vent concentration to an area of low solventconcentration. Osmosis can be demon-strated using an egg with its shell removed.Remember to handle the egg carefully so themembrane does not break. Place the egg ina container with enough corn syrup to sur-round the egg. Observe what happens to theegg over the next day. The egg should shrivel

up. This occurs because water passes frominside the egg (area of higher concentrationof water) across the egg’s membrane into thecorn syrup (lower concentration of water).The membrane allows the water moleculesto pass through; the sugar molecules are toolarge to fit through the openings in the mem-brane. The water molecules move in bothdirections across the membrane, but overallmore molecules move from the water insidethe egg into the corn syrup than vice versa.

To reverse the process, use a spoon totake the egg out of the corn syrup. Carefullyblot off the syrup. Now place the egg inwater to which several drops of food colorhave been added. Again, observe what hap-pens over the next day or two. The egg willexpand and the inside of it should be tintedby the food color.

Activity 3 Surface TensionMaterials: small drinking glass such as

juice glass, water, small metal objects suchas pennies, paper clips, small nails,medicine dropper (optional)

Surface tension is the energy needed toincrease the surface area of a liquid per unitarea. It is a measure of the intermolecularattractions of a liquid. Because of hydrogenbonding, water has a high surface tension. Todemonstrate this, take a small glass and fill itto the very top with water. Once you believethe glass is completely full, start to place thesmall metal objects into the glass one at atime. Count how many of the objects you canplace in the glass before water spills out.


Another variation of this activity is to guessthe number of water drops that can be placedon a penny using a medicine dropper.

In both variations of this activity, amuch larger number of objects or drops thanexpected will probably be found. The reasonfor this is that the intermolecular forces andsurface tension of water cause the watermolecules to form into a spherical, tightlybound shape.

Activity 4 ElectrolysisMaterials: water, 9 V battery, vinegar,

distilled water, salt, clear drinking glassThe electrolysis of water can be seen by

taking a 9 V battery and placing it in enoughdistilled water to cover the entire battery.Make sure the electrodes are several cen-timeters below the water’s surface. Afterplacing the battery in the distilled water,note any evidence of a reaction. There arenot enough free ions in distilled water toconduct electricity and no evidence of areaction should be observed. Now add a tea-spoon of vinegar to the water and note whathappens at the battery terminals. Bubblesform around the terminals and then a steadystream of tiny bubbles emerge from bothterminals of the cell.

When vinegar is added to the water, itpartially dissociates into free hydrogen andacetate ions:

HC2H3O2(aq) H� � CH3COO�

The ions enable a current to flow throughthe water-vinegar solution. At the anode(negative terminal), oxidation of wateroccurs:

2H2O(l) O2(g) � �H�(aq) � 4e�

While at the cathode (� terminal), reduc-tion of hydrogen ions takes place:

2H3O� � 2e� H2(g) � H2O(l)

Try this activity with tap water. Evidence ofbubbles would indicate the presence of freeions from dissolved minerals in the water.

Instead of vinegar, add salt to the dis-tilled water. Bubbles will again appear ateach terminal. If the concentration of salt ishigh enough, chlorine gas, Cl2, is producedat the anode from the oxidation of chlorineions:

2Cl� � 2e� Cl2(g)

Hydrogen is produced at the cathode by thereaction:

2H2O � 2e� 2H2 � 2OH�

Activity 5 Solubility of Packing Materials

Materials: Styrofoam packing peanuts,starch packing peanuts, acetone, water, glass

Polystyrene is a common polymer usedto make Styrofoam containers, packingpeanuts, and insulation. Another kind ofpacking peanut used is made of starch. Placeabout a cup of acetone in a glass container.Acetone is available wherever paint is sold.Do this in a well-ventilated area. Be carefulwith the acetone, it can remove paint andruin finishes. Add a starch packing peanutto the acetone and then a Styrofoam pack-ing peanut. The starch peanut will not dis-solve, but the polystyrene peanut will.Dilute the acetone with ample amounts ofwater before disposing down the sink.

Repeat the activity using warm waterinstead of acetone. When water is used, thestarch peanut dissolves and the polystyrenepeanut does not. The acetone and polystyrenedo not exhibit hydrogen bonding, while starchpeanuts and water do exhibit hydrogen bond-ing. This illustrates the general principle of“like dissolves like.” Not all the polystyrenedissolves because of the presence of cross-linked units that form the polystyrene.


Activity 6 Cabbage pH IndicatorMaterials: red cabbage, pot, strainer,

household liquids to test (ammonia, vinegar,tap water, lemon juice, etc.)

Acid-base indicators can be extractedfrom different plant materials. An easy wayto prepare an indicator is to cut about halfof a small red cabbage into small pieces.Place these pieces in a pot, cover with water,and boil until the water turns a deep purple.It should take about 5 to 10 minutes of boil-ing. Let the solution cool and then eitherpour off the liquid or strain to separate theliquid from the cabbage leaves.

The cabbage can now be used as a pHindicator to test different substances. Anapproximate scale is shown below.

Different household liquids can be testedsuch as ammonia, vinegar, tap water, lemonjuice, and Sprite. Items such as orange juicethat possess a characteristic color should bediluted and then tested. Diluted solutions ofsolids like dishwasher soap, cream of tartar,and baking soda can also be made. The indi-cator can be frozen or mixed with alcohol topreserve for later use.

The chemical compounds responsible forthe indicator properties of red cabbage arecalled anthocyanins. These compounds losehydroxide ions at low pH, creating a colorchange. Other plants can be tried as indica-tors including beets and red onions beets.

Activity 7 PuttyMaterials: small glass jars, measuring

cup, Elmer’s glue, Borax, food coloringA simple putty can be made with

Elmer’s glue by mixing 2 tablespoons (30mL) of glue with 4 teaspoons (20 mL) ofwater in a jar. In another jar, mix about 1⁄

2cup (120 mL) of water, 1⁄

2teaspoon of

Borax, and 2 or 3 drops of food color. Takea tablespoon (15 mL) of the Borax solution

and add it to the glue solution. Mix the twosolutions until a viscous putty forms.Remove the putty from the jar and note itsability to stretch, bounce, and so on.

Activity 8 Copper to GoldCaution: This activity should be done

with adult supervision in a well-ventilatedarea. By using only small amounts of chem-icals, it can be done safely. Wear safetyglasses.

Materials: safety glasses, several shinynew pennies, metal file, hot-dipped galva-nized zinc nails 8 cm (3 inches) or longerare preferred, paper or cardboard, a hotplate, Red Devil lye, a small ceramic pot(capacity of around 1⁄

2cup), tweezers, sev-

eral plastic containers such as yogurt cupsor butter tubs

In this activity, galvanized nails are usedas a source of zinc. Galvanization is a pro-cess by which metals such as steel aredipped in zinc to protect them from rusting.One type of nail is dipped in a hot bath ofmolten zinc to form a protective coating.Zinc is more easily corroded than iron, so itoxidizes rather than the steel.

Before preparing the zinc, mix the lyesolution. Be careful with lye! Read theprecautions on the package! Do this partoutside or in a well-ventilated area. Placeabout 1⁄

2cup of cold water in a yogurt or

other plastic container. Slowly add a table-spoon of lye and mix gently. The solutionwill warm considerably. Place the solutionin a safe place where it won’t spill. It shouldsit for at least 30 minutes. To prepare thezinc, file the outside of the nails on a pieceof paper or cardboard. Your goal is to obtaina small pile of minute zinc chips. Enoughzinc can be obtained using as little as fournails. Only a little zinc is needed, perhaps apile large enough to cover this circle:


Turn on the hot plate and fill a plasticcontainer with about a cup of tap water. Setthe container aside for now. Place 3 table-spoons of the previously prepared lye solu-tion in your small pot. Add the zinc shavings.The zinc should sink to the bottom. A littleswirl of the pot will cause the zinc to form asmall pile on the bottom of the pot. Place apenny on top of the zinc pile with the tweez-ers. Now put the pot on the hot plate. Bringthe pot just to its boiling point; if you seebubbles that indicate boiling, turn down theheat. As the solution heats, zinc will plate onthe penny. The penny will obtain a shiny sil-ver luster. Use the tweezers to remove thepenny and rinse it in the tapwater container.Blot the penny dry. Now place the penny onthe hot plate. The penny will turn a dull goldcolor. It may only take a few seconds or ahalf a minute or so depending on the hotplate’s temperature. Use the tweezers toremove the penny and drop it in the water.The penny will turn into a shiny “gold” coin.If a gold color change is not observed on thehot plate, the hot plate’s temperature is toolow. In this case, the silver coin can be heldwith tongs and slowly waved through a gasstove flame until it turns color. Immerse thehot coin in water immediately after flaming.

Once the cooking process is perfected,several pennies can be prepared at a time.The lye-zinc solution may have to be swirledduring the coating process to get good cov-erage. You may also need to add more lyesolution. Add small amounts of solutioncarefully.

Silver-colored coins do not have to betransformed into gold coins immediately.The coins can be plated with zinc and savedfor heating at a later time. Excess lye solu-tion can be poured down the drain afterdiluting with plenty of cold water. Plateenough pennies until the zinc is used. Oldpennies can be plated, but will produce aduller, dirty finish.

The silver color is due to the plating ofzinc on the penny. The final heating processcauses the copper and zinc to fuse, produc-ing a brass alloy. Brass is an alloy of zincand copper. Typical brass alloys containbetween 3% and 30% zinc.

Activity 9 Polar Water MoleculeMaterials: combWater is a polar molecule. The oxygen

atom in the water molecule carries a partialnegative charge and the hydrogen atomscarry a partial positive charge. This polaritycan be observed by using a charged comb.Turn on a faucet so that a very small con-tinuous stream of water is coming out of thefaucet. Open the faucet the minimumamount to get a continuous stream. Take acomb and run it several times through yourhair. Your hair should be dry, and you shouldbe able to hear little discharge sparks as youcomb your hair. Alternately, the comb canbe rubbed on a piece of fur or even a pet’sfur. Bring the charged comb up to the streamof water. Notice how the water is attractedto the comb. Move the comb around andobserve how the water “dances” in responseto the comb’s movement. The positivehydrogen end of the polar water moleculesis attracted to the negatively charged comb,while the negative oxygen is repelled.

Activity 10 Immiscible LiquidsMaterials: old fruit juice glass bottle

(8–16 oz.) with screw-on lid, ethyl alcohol,charcoal lighter fluid, olive oil, water, dish-washing detergent

Immiscible liquids can be combined togive a colorful display in a bottle. Take aclear glass bottle and remove the label. Abottle used for fruit juices works well. Soak-ing the bottle in soapy water for severalhours makes it easier to remove the label.Some scraping may be necessary. Make surethe bottle has a screw-on cap. Once the bot-tle is clean fill it halfway with water, then


add ethyl alcohol until the bottle is about95% full. Ethyl alcohol is available in paintand hardware stores. Finally, add severaldrops of charcoal lighter fluid until the totallevel of the fluids in the bottle is just belowthe lip. Tightly screw the cap on. Turning thebottle on the side produces a rolling displayof lighter fluid suspended in the alcohol.Adding a drop or two of food coloringmakes a more colorful display.

As another example of immiscible liq-uids, take a jar with a lid and place waterand olive oil in the jar. The water and oil willform two distinct layers. Shake the jar tomix the oil and water, and they will separate.You can take an identical jar with the sameamount of water and olive oil and add a lit-tle dishwashing detergent. Shake both jarsand observe how long it takes for the waterand oil to separate in the jars. The jar withdetergent should keep the olive oil sus-pended in water longer. Detergents weakenthe intermolecular forces between watermolecules and the detergent molecules sur-round the oil molecules with watermolecules. Detergents’ action on oil in watergives them their cleaning ability.

Activity 11 Creation of CarbonDioxide and Oxygen

Materials: baking soda, vinegar, 3%hydrogen peroxide, yeast, wood splint, plas-tic containers such as butter or cottagecheese tubs, lids, matches

Gases are produced in many commonreactions. The reaction of baking soda,NaHCO3, and vinegar, HC2H3O2, producescarbon dioxide according to the reactions:

NaHCO3(s) � HC2H3O2(aq) NaC2H3O2(aq)

� H2CO3(aq)

H2CO3(aq) H2O(l) � CO2(g)

Place about a tablespoon of baking soda ina 16 oz. plastic container. To the container,add about a 1⁄

2cup vinegar and immediately

cover the container with a loose-fitting lid.Do not snap the lid on tight, but simply laythe lid across the top of the container. Apiece of cardboard or any suitable coveringcan be used for a lid. Let the reaction pro-ceed for 15–30 seconds, light a match, andthen lower it into the tub. The match willimmediately be extinguished by the CO2.Carbon dioxide is used in certain types offire extinguishers.

To produce oxygen, place about 1⁄2

cupof 3% hydrogen peroxide, H2O2, in the bot-tom of a tub. To the tub, add a teaspoon ofyeast and immediately cover the container.After the reaction has proceeded for a minuteor two, light one end of a wooden splint. Ashaved piece of wood or Popsicle stick can beused for a splint. Blow out the flame on thesplint and immediately lift the lid and placethe still glowing end of the splint inside thetub. The splint should reignite in the rich oxy-gen environment. Oxygen gas is producedfrom the decomposition of hydrogen perox-ide according to the reaction:

2H2O2 O2(g) � 2H2O(l)

The yeast acts as a catalyst to speed up thedecomposition reaction.

Activity 12 Air Pressure and Gas LawsCaution: This experiment should be

done with adult supervision. Wear safetyglasses or goggles.

Materials: safety glasses, hard-boiledegg, large glass juice bottle with a mouthslightly narrower than an egg, rubbing alco-hol, cotton, aluminum pie pan, tongs

On Earth partial vacuums are used forall kinds of applications. The most commonis the vacuum cleaner. Another one is drink-ing out of a straw. The word “suck” is usedwhen dealing with vacuums or straws, butthis term is misleading. What happens whena liquid moves up a straw is that a partialvacuum is created and atmospheric pressure


pushes liquid into this partial vacuum, forc-ing it up the straw.

To do this activity, first boil and peel anegg. The egg should be slightly wider thanthe mouth of the bottle. Wet the cotton ballwith a little alcohol. You don’t need a lot.Put the cap back on the rubbing alcoholcontainer and remove it from the area.Now place the cotton ball in the pie pan andlight it. Using the tongs, grab the flamingcotton ball and drop it in the bottle. Placethe egg, narrow end up, in the opening. In afew seconds the flame should go out andatmospheric pressure should push the egginto the bottle. To get the egg out, positionthe egg with its narrow end in the opening.Tilt the bottle back and blow into it as ifyou’re blowing a trumpet. The egg shouldslide out.

When asked why the egg is forced intothe bottle, most people respond that a par-tial vacuum is created because the oxygen isremoved during combustion. While this istrue, it also must be remembered that car-bon dioxide is created and replaces the oxy-gen. The egg is forced into the bottle

because of the relationship between temper-ature and pressure. The flame heats up theair in the bottle. Because the bottle is ini-tially open to the atmosphere, the pressureinside the bottle is the same as the atmo-spheric pressure. Once the opening of thebottle is blocked with the egg, the flamegoes out due to lack of oxygen. The gasmolecules comprising the air inside the bot-tle are no longer heated and are less ener-getic. This translates into a drop in pressureinside the bottle and in pops the egg.

To get the egg out, it wouldn’t makesense to suck on the bottle’s mouth. Thepressure inside the bottle is the same asatmospheric pressure. Extracting more airwould just lower the pressure inside the bot-tle even further. To get the egg out, you haveto increase the pressure inside the bottle sothat it is higher than the atmospheric pres-sure. This is accomplished by blowing intothe bottle. Give the bottle a good blow andout slides the egg. It may take a little prac-tice, but you should be able to get the eggout with a good blow.

Activity 13 Balloon in the BottleMaterials: safety glasses, gloves, baby

bottle with screw-on lid, 15� balloon,microwave oven, water, potholder

Wear safety glasses for this activity. Asan alternative or addition to Activity 11, thisactivity can be performed. Prepare a babybottle for this activity by removing the nip-ple from the top lid. The screw on lid shouldnow have a hole in place of the nipple. Takea balloon and inflate several times to stretchit. Stretch the balloon over the opening ofthe lid. Stretch the balloon evenly over thehole in the lid so that the balloon’s openingis centered. Put the lid close to themicrowave oven. In the bottle, place severaltablespoons of water. About 3 cm (1 inch)of water on the bottom of the bottle is suffi-cient. Put on gloves. Place the bottle in aSource: Rae Déjur


microwave for a minute or however long ittakes to boil the water in the bottle for atleast 30 seconds. After the water has boiledfor at least 30 seconds, using a potholder ormitten, quickly remove the bottle from themicrowave oven and screw on the lid. Oncethe lid is screwed on, make sure the balloonis upright and not folded over to one side.You may have to move it slightly to theupright position.

In a short period, usually less than aminute, the balloon will be pushed into thebottle similar to the egg being pushed intothe bottle. In this case, the baby bottle fillswith water vapor when it is boiling. Whenthe bottle is removed from the oven, it isfilled with steam and it immediately beginsto cool. Quickly placing the lid on the bot-tle traps the steam. As the bottle continuesto cool, the steam condenses. This decreasesthe pressure inside the bottle, and the bal-loon is forced into it. The same principle isused to seal jars when making jelly or can-ning.

Activity 14 Heat Capacity of WaterThis experiment should be done with

adult supervision. Do this activity out-doors or in a well-ventilated area.

Materials: old cotton socks, container,70% rubbing alcohol, measuring cup, alu-minum pie pan, tongs, bucket of water, longfireplace matches

Water has a very high heat capacity.Substances with high heat capacities expe-rience smaller temperature changes whenheat is exchanged between them and theirsurroundings as compared to substanceswith smaller heat capacities. In this activity,the high heat capacity of water enables apiece of fabric to stay below its kindlingtemperature and in the process the fabricdries. Prepare a mixture of 1⁄

3cup water and

2⁄3

cup 70% rubbing alcohol. Once you aredone with the alcohol, remove it from thearea. Cut a small, approximately 10 cm (3inch), square out of an old cotton sock oranother piece of clothing. Soak the fabric inthe water-alcohol mixture. Ring out the fab-ric, dry your hands, and place the fabric inthe pie pan. Light the fabric with a long fire-place match. Alternately, hold the fabricwith tongs and light it, then place it in thepie pan. The alcohol in the fabric burns, butthe fabric does not ignite. The water absorbsmuch of the heat and keeps the temperaturebelow the ignition temperature of the fabric.

This activity can be done with ethylalcohol or 95% rubbing alcohol. In thiscase, use 1⁄

2cup each of water and alcohol.

This activity can also be done with papersuch as a dollar bill.

Activity 15 Paper ChromatographyMaterials: white construction paper,

alcohol, nail polish remover, felt tip pensand markers, food coloring, small dish,small juice glasses or jar, toothpicks

Chromatography is a method used toseparate mixtures of substances into theirconstituents. It is based on the principle ofdifferential attraction of the components ofSource: Rae Déjur


a mixture to mobile and stationary phases asthe mobile phase moves through the sta-tionary phase. In this activity, acetone willbe the mobile phase and the constructionpaper will be the stationary phase. The com-ponents in inks will be separated.

Cut a sheet of white construction paperinto several strips approximately 15 cm (6�)long and 2 cm (1�) wide. About 2 cm fromthe bottom of the strip, make several smallmarks using the different felt pens.

Pour enough nail polisher remover (theprimary ingredient of nail polish remover isacetone) into a jar or small glass to cover thebottom. Only a small amount of acetone isneeded. The depth of the acetone in the jardoes not have to be more than 1 cm (1⁄

2inch). Place the prepared construction paperin the acetone. The dots should be above theacetone. The acetone will move up the con-struction paper and the ink mixture will beseparated. Compare several different col-ored inks and note the components present.Individual food colors or mixtures of foodcolors can be made and separated. To applyfood colors, place a drop of food color on adish or piece of paper, use a toothpick, anddab the color onto the paper. When finished,the acetone can be diluted with water andflushed down the sink.

Chemists use a number of very sophis-ticated chromatographic methods for analy-sis of drugs, DNA, petroleum products, and

air samples. Chromatography is used to sep-arate and identify different substances.

Activity 16 Solution SaturationMaterials: safety glasses, quart jar with

lid, tall drinking glass, slaked lime,�Ca(OH)2 , straw, permanent marker

Solvents have the ability to hold only acertain amount of solute, and then theybecome saturated. Once saturation isreached, a solute generally precipitates outof solution. A solution can become saturatedin a number of ways. Some of these includea change in the solution’s temperature orpressure, a chemical reaction occurring inthe solution, and adding more solute to thesolution. In this activity, the addition of CO2

to limewater causes a precipitate of CaCO3

to form.Limewater is a saturated aqueous cal-

cium hydroxide, Ca(OH)2, solution. Tomake limewater, a small amount of calciumhydroxide is needed. Calcium hydroxide ismarketed commercially as slaked lime orhydrated lime. It is used for cement, increas-ing the pH in soils, and water treatment.Lime may be obtained from building mate-rial stores in the cement section and in agri-cultural stores. The smallest quantities soldare generally 5- or 10-pound bags, whichcost a few dollars. Because only a teaspoonof lime is needed (the solubility of calciumhydroxide in water is 0.1g per 100 mL), askthe sales clerk if there are any broken bagsfrom which you can take a tablespoon oflime. Often there will be enough lime dustwhere it is stored to obtain an ample amountfor this activity.

Take a jar and using a permanentmarker label the jar “Limewater DO NOTDRINK!” Fill the labeled jar almost fullwith water and add 1⁄

4teaspoon of lime. Cap

the jar with the lid and shake the solution.The water will become cloudy. Set the jar ina safe place and let the Ca(OH)2 settle. Itwill probably take several hours. A good


idea is to mix the solution and let it settleovernight. After the settling period, thelimewater should be clear. Solid calciumhydroxide will be present on the bottom ofthe jar. Handling the jar gently so as not todisturb the solid Ca(OH)2 carefully pourabout 1⁄

2cup of the clear solution into a tall

drinking glass or jar. Put on safety glassesand using the straw blow gently into thelimewater. Continue to blow bubbles intothe limewater and notice how it changes. Becareful not to ingest any limewater. Thesafety glasses protect your eyes in case anylimewater splashes out of the jar.

The limewater should eventually turncloudy as you blow into it. It may take acouple of minutes of blowing, so be patient.Blowing into the limewater introduces CO2

from your breath into the solution. Carbondioxide reacts with the Ca(OH)2 in thelimewater to produce calcium carbonate,CaCO3:

CO2(g) � Ca(OH)2(aq) CaCO3(s) � H2O(l)

Calcium carbonate is insoluble in water, andthis produces the cloudy appearance of thelimewater.

Make sure to dispose of the used lime-water immediately upon completion of thisactivity and wash out the drinking glass.Water can be added to the unused limewa-ter in the labeled jar and stored until usedagain. A pH indictor can be added to thelimewater before blowing into it to see if achange in pH can be detected. Blowing intothe limewater causes carbonic acid, H2CO3,to form as CO2 dissolves in water. This low-ers the pH of the limewater:

CO2 � H2O H2CO3

Activity 17 Self-Sealing PolymerMaterials: 9–15� round balloon, bam-

boo skewer, Vaseline or baby oil

Inflate a round balloon to a moderatesize and tie off the end. Take a bambooskewer and rub some Vaseline or baby oil onit. Using the skewer, pierce the balloon nearthe knot. You may need to use a twistingmotion to break through the latex. The bal-loon should not pop when it is pierced. Con-tinue to thread the skewer through theballoon and pierce the end opposite theknot. The balloon will now have a skewerthrough it, and it may slowly deflate. Theskewer can be taken out of the balloon andused to pop the balloon by piercing throughthe side.

Latex balloons are made of polymers.The latex near the knot and top of the bal-loon are not stretched as tightly as latex onthe side. This can be seen by observing thetransparency of the different parts of the bal-loon. The knot and top areas of the balloonhave a greater polymer density. Therefore,the balloon in these areas has a greater abil-ity to stretch and partially seal itself aroundthe skewer. On the sides of the balloon, thetightly stretched latex cannot seal around theneedle and the balloon pops.

Activity 18 Cathodic ProtectionMaterials: nongalvanized steel nails,

penny, metal file, small dishCathodic protection involves connect-

ing a metal to be protected to another metalthat is more easily oxidized. The more eas-ily oxidized metal serves as the anode andthe metal to be protected is the cathode inan electrochemical cell. The metal that isoxidized is called the sacrificial anodebecause it is sacrificed to protect anothermetal. Metals such as zinc and magnesiumare often used to protect iron. Cathodic pro-tection is demonstrated in this activity byusing two steel nails. The nails are placed ona shallow dish. Using a white or light-colored dish displays the oxidation better�iron (III) oxide, Fe2O3, referred to com-monly as rust is more visible on a light-


colored surface . To obtain zinc, take apenny and file off the thin copper coating.Pennies dated before 1982 are 95% copperand 5% zinc. In 1982, the composition ofpennies was changed to 2.5% copper and97.5% zinc. Place the filed zinc penny onone side of the dish. Now place the two nailsin the dish. Place one nail so its tip is in con-tact with the penny and place the other nailby itself. Make sure the nails are separatedan appreciable distance. Pour just enoughwater in the dish to cover the sides of thenails, keeping the top exposed to air.

Over the next couple of days observethe pennies. The nail in contact with the zincshould accumulate much less rust comparedto the unprotected nail. The nail not in con-tact with zinc will show signs of oxidationin just a few hours. Appreciable rust willaccumulate on this nail. The zinc oxidizeson the other nail, but the oxidation results ina protective coating of zinc oxide formingon the zinc.

Activity 19 Water PurificationMaterials: activated charcoal, coffee fil-

ter, funnel, jars, food coloring, dirt, waterActivated charcoal is a pure form of

charcoal made by heating wood to a hightemperature in the absence of oxygen. Acti-vated charcoal is used in granular form toremove color, odor, and impurities fromgases and liquids. The charcoal granulesprovide a large surface area to contact thefluids being purified. In this activity, a char-

coal filter is used to purify water. Take around coffee filter and flatten it out. Fold itin half and then fold it in half once more toform a quarter of a circle. Place the pointedportion of the folded filter down into thefunnel and open it up to form a cone. Fill theinside of the cone with activated charcoal.Activated charcoal can be found whereveraquarium supplies are sold. A small bag canbe purchased for a few dollars. Place thefunnel’s neck inside a jar.

Next, place some water in a jar and pro-duce your “polluted” water by adding a cou-ple of drops of food coloring, a little dirt,and other items to give the water an odor. Toproduce an odor, a little vinegar, ammonia,onion juice, garlic juice, or other pungentitems can be used. Filter the polluted waterthrough the charcoal and observe the filtrate.Its color should be clear and much of theodor should be gone.

Activated charcoal filters are used as aform of treating wastewater. Small filterscan also be installed on home faucets toremove impurities. This improves theappearance and taste of drinking water.

Activity 20 Alka-Seltzer RocketMaterials: safety glasses, Alka-Seltzer

tablets, 35 mm film canisters (the white Fujibrand ones work better than the black ones),aluminum pie pan, water

Alka-Seltzer contains several ingredi-ents including citric acid, C6H8O7, andsodium bicarbonate, NaHCO3. When anAlka-Seltzer tablet is placed in water, thecitric acid and sodium bicarbonate react toproduce carbon dioxide. This reaction can


be used to propel a film canister into theair.

Caution: Wear safety glasses and dothis activity in an open area. Do not do itindoors with a low ceiling under light fix-tures.

Put on your safety glasses. To make arocket, fill a film canister about 1⁄

3full of

water. Place the canister in the center of thepie pan. Break an Alka-Seltzer tablet intothree piece of about equal size. Drop one ofthe pieces into the canister, snap on its top,and quickly invert the canister so its top isresting on the pie pan. Step back and wait.The pressure inside the canister increasesdue to the accumulation of carbon dioxidegas inside it. This pressure builds up untilthe canister body suddenly explodes fromthe pie pan launching pad. If the top of thecanister is quickly replaced and the canisterrepositioned, another blast can occur. In fact,several blast-offs can be obtained with oneAlka-Seltzer tablet.

Alka-Seltzer is a good substance to useto explore how temperature, surface area,and quantity affect the rate of a chemicalreaction. The degree of fizzing is a direct

indication of the amount of carbon dioxidebeing produced and how fast the reaction isproceeding. Bayer Company, the producerof Alka-Seltzer, has several experiments onits Web site: http://www.alka-seltzer.com/as/experiment/student.

Activity 21 Dry IceMaterials: film canisters, pie pan, cab-

bage indicator, plastic drinking cup, tongs,insulated gloves, dry ice. Caution: Dry iceexists at approximately –100°C. Use tongsand insulated gloves when handling dry ice.Because it sublimates directly from solid togas and is denser than air, always use andtransport dry ice with adequate ventilation.When transporting dry ice in a vehicle, keepwindows cracked open. Use an insulated con-tainer with a loose-fitting top to hold dry ice.Do not place dry ice in containers with a tightlid, it will explode. Never take dry ice in anelevator.

Dry ice is solid carbon dioxide. It is soldin some grocery stores, fish markets, andspecialty stores. The Alka-Seltzer rocket canbe repeated using dry ice in place of Alka-Seltzer. Do not use water, but place a smallpiece of dry ice into the canister.

Dry ice can also be used with cabbageleaf pH indicator. Place several drops ofcabbage leaf indicator into a plastic tub con-taining water. Drop in several cubes of dryice. The color should change indicating adrop in pH as the carbonic acid, H2CO3,concentration increases in the water.

Carbon dioxide is colorless. The vaporseen in the presence of dry ice is watervapor condensing as the dry ice cools thesurrounding air.

Activity 22 Cooling Curve of WaterMaterials: Styrofoam cups with lids,

thermometer, graph paperA cooling curve for water can be graphed

by taking heated water and placing it in a Sty-rofoam cup and taking temperature reading at

Source: Rae Déjur


regular time intervals. A graph can then beplotted with time on the x-axis and tempera-ture on the y-axis. This activity can be devel-oped into an interesting investigation. Takeone cup of room-temperature water and placeit in a Styrofoam cup with a lid. Punch a holein the lid with a sharp pencil. Place a ther-mometer through the hole in the lid. Put thecup assembly in the freezer and monitor itstemperature every five minutes for a givenperiod—30 minutes or one hour. Plot thecooling curve for this first trial. Now repeatthe activity, but this time use one cup of hotwater. Boil water in a pan and use this waterto repeat the experiment. Compare the cool-ing curves. Common sense says that theroom-temperature water should be colderthan the hot water after the same time period.Its temperature starts roughly 70°C lower thanthe hot water.

The results of this activity may be sur-prising. Try to determine some reasons foryour results.

The activities described in this sectionshould be used as a starting point for scien-tific investigation. There is no end to thenumber of scientific investigations that canbe performed in any given area. Many pro-fessional scientists spend their whole liferesearching one specific area. They contin-ually probe, ask questions, design experi-ments, collect data, and analyze the results.Throughout the process new questions andsurprises appear that require further investi-gation. The process of science is a never-ending attempt to understand nature; themore it is understood, the more it seems thatthere is to understand.

IntroductionThis book began by portraying a chemist asthe quintessential absent-minded professor.A bespectacled male concocting strangereactions in the lab is the stereotypical pic-ture of a chemist perpetuated in movies andcartoons. In reality, a relatively small pro-portion of chemists spend their days dressedin white lab coats conducting experiments,many are women, and they work in a vari-ety of settings. At the end of Chapter 19, itwas stated that over 1,000,000 workers areemployed in the chemical industry. Many ofthese workers are not chemists. Conversely,many chemists are employed in a myriad ofprofessions not directly related to the chem-ical industry, including banking, law, sales,teaching, and public service. Approximately100,000 chemists are employed in chemicalfields in the United States, with roughly60% of these involved in chemical manu-facturing. This final chapter presents infor-mation for exploring careers in chemistryand chemistry-related fields.

Chemist’s Job Descriptionand Training

Because chemistry is such a broad field,it is difficult to give a concise job descrip-tion of a chemist. In general, chemists studythe composition, structure, and characteris-tics of matter in an attempt to improve substances and products; discover new sub-stances and products; and improve anddevelop chemical processes. Chemists workin a multitude of industries including agri-culture, medicine, pharmaceuticals, textiles,petrochemicals, mining, plastics, hazardousmaterials, electronics, adhesives, paints,household products, and cosmetics.

Chemists are frequently grouped intosubcategories to distinguish their area ofexpertise. Five large chemical specialtiesinto which many chemists may be classifiedare inorganic chemist, organic chemist,physical chemist, biochemist, and analyticalchemist. Inorganic chemists work with noncarbon-based compounds and frequentlywork in areas such as solid-state electronics,

21

A Future in Chemistry

328 A Future in Chemistry

the metal industry, and mining. Organicchemists focus on carbon-based compoundsworking with petrochemicals, pharmaceuti-cals, plastics, synthetic fiber, and rubber.Many organic chemists use organic synthe-sis to create new compounds. Physicalchemists focus on the relationship betweenthe chemical and physical properties of mat-ter. Employment opportunities exist forphysical chemists in alternative energyfields, nuclear power, and material science.Biochemists work on compounds associatedwith life and are frequently employed by thedrug industry, in medical research, and ingenetic engineering. Analytical chemistsdetermine and quantify the composition ofsubstances. Analytical chemists are heavilyemployed in the pharmaceutical industry todetermine the nature and purity of drugs.They also are heavily employed in the envi-ronmental field where they sample the envi-ronment for pollutants. Analytical chemistsalso develop new methods for characteriz-ing matter.

A job in chemistry generally requires afour-year undergraduate degree in chem-istry, although chemical technician positionsare available with a two-year degree. Inareas such as industrial research or collegeteaching, an advanced degree, typically aPh.D., is required. Students preparing for acareer in chemistry should focus on the sci-ences and math in high school. In additionto having the proper academic backgroundto enter a college chemistry program, stu-dents should be naturally inquisitive, enjoylaboratory work, be able to work both indi-vidually and in groups, and be comfortableworking with computers and instruments.

Hundreds of colleges and universitiesoffer bachelor degrees in chemistry in theUnited States. Around 600 bachelor pro-grams, 300 master’s programs, and 200 doc-toral programs are approved by theAmerican Chemical Society. With so many

programs at all levels, it is expected that avariety of curricula and programs are inplace. Chemistry programs at large univer-sities offer a wide range of courses thatallow students to focus in one of several spe-cialty areas in chemistry. Small schools havemore focused programs, but may havedeveloped a unique niche in one specialtyarea. Although there is a wide diversity ofbachelor programs, an undergraduate stu-dent can expect to complete a core of basicchemistry, science, and math courses. Whilecolleges differ in how they organize andname their courses, there is consistency inthe content. For example, some schools mayoffer laboratory work as independentcourses, while others may include labora-tory work within the framework of a maincourse.

A model undergraduate curriculum fora bachelor degree in chemistry is presentedin Table 21.1. This model shows the coretechnical courses for a traditional four-yearundergraduate degree in chemistry.

The model program is based on asemester system and generic names are usedto identify courses. Most undergraduatechemistry programs reflect the courseslisted in the model program. Embeddedwithin the chemistry curriculum is contentthat may not be apparent in the course titles.Topics such as lab safety, chemical infor-mation retrieval, and record keeping fit intothis category. Depending on the school,actual coursework in chemistry may onlycomprise about one-third of the totalamount of credit hours to obtain the bache-lor degree. The remaining coursework isfilled with additional science courses andbasic requirements such as English, human-ities, and social sciences.

Elective courses can be used by the stu-dent to enhance their technical education.These courses are often neglected by stu-dents who assume that mastering the tech-

A Future in Chemistry 329

nical material is sufficient to become a suc-cessful chemist. For example, one of themost highly valued skills sought in gradu-ates by employers is the ability to commu-nicate clearly in writing and speech. Atechnical writing or speech course helps stu-dents develop these skills. Another benefitof elective courses is to help the studentdevelop career goals. For example, if a stu-dent wanted to work in chemical sales orwork in a non-English speaking country,courses in marketing or foreign languagewould be beneficial.

Another option that the student shouldconsider is completing an internship as partof the undergraduate program or summeremployment in the chemical industry. Somecolleges require an internship as part of theirprograms. There are many advantages to

completing an internship. An internship pro-vides a professional experience that exposesthe student to practical applications ofchemistry. The internship can be used to net-work with professionals to assist in obtain-ing employment upon graduation or beingadmitted into a graduate program. Aninternship exposes students to advancedequipment and processes. For instance,work as an intern in a refinery wouldacquaint the student with distillation towers,cracking columns, heat exchangers, reac-tors, and other specialized equipment. Stu-dents completing research internships oftenhave access to modern instrumentation notavailable at the college. Internships can befound through the college’s chemistry orcareer service department. Students mayalso contact the human resource department

Table 21.1Model Undergraduate Chemistry Curriculum


of large chemical companies, the AmericanChemical Society, or individual chemists orfaculty members to learn about internships.The Internet is an invaluable tool for explor-ing internship opportunities.

Careers in ChemistryIn addition to the five general special-

ties of inorganic, organic, analytical, physi-cal and biochemist, there are numerousother specialties that are named by placing adescriptive word in front of the word“chemist.” A small sampling includes envi-ronmental chemist, food chemist, cosmeticchemist, soil chemist, atmospheric chemist,medicinal chemist, polymer chemist,petroleum chemist, geochemist, and indus-trial chemist. Other specialties combine dif-ferent fields of study with chemistry and usethe term “chemical” as a descriptive term.Examples include chemical engineer, chem-ical oceanographer, and chemical techni-cian. There are other professions that requirea thorough knowledge of chemistry, forexample, pharmacy and hazardous materialmanagement. To fully describe the numer-ous jobs dependent on chemistry would beimpossible. In this section, a brief descrip-tion of selected chemical careers is given.Although only a limited number of careersare presented, the descriptions serve todemonstrate the importance of chemistry ina number of fields and the breadth of oppor-tunities available to people with back-grounds in chemistry.

Chemical engineers use engineeringprinciples to solve problems of a chemicalnature. Nearly three-fourths of the approxi-mately 33,000 chemical engineers in theUnited States work in chemical manufac-turing. They design chemical plants,develop chemical processes, and optimizeproduction methods. A primary concern ofchemical engineers working in industry is to

develop economic methods to lower pro-duction costs. Such measures might includereducing waste, improving energy effi-ciency, or modifying reaction conditions.Chemical engineers work in research anddevelopment departments to develop newproducts, improve existing products, anddetermine the best way to produce a prod-uct. Undergraduate programs in chemicalengineering include a basic core of chem-istry courses in inorganic, organic, andphysical chemistry along with a core of gen-eral engineering courses such as thermody-namics, computer-aided design, fluiddynamics, and economics. Specializedchemical engineering courses in an under-graduate curriculum include plant design,process control, unit operations, and trans-port processes.

Chemical technicians assist chemistsand chemical engineers. Much of their timeis spent in the laboratory, where they per-form chemical analyses, test products, cali-brate equipment, prepare solutions andreagents, and monitor quality control. Inaddition to laboratory work, chemical tech-nicians may collect samples from the field.For example, technicians in environmentalwork may collect water samples from lakesto analyze in the field or transport back tothe lab. Petroleum technicians monitor con-ditions at oil and gas wells. Chemical tech-nician positions do not require a four-yeardegree, although many individuals workingas technicians have or are in the process ofobtaining a bachelor’s degree in chemistry.A job as a chemical technician is an excel-lent entry-level position for chemists. Thereare a limited number of two-year chemicaltechnician programs offered. These are gen-erally found in community colleges andoffer a study in basic chemistry, science, andmath with a heavy emphasis on laboratorywork. Although not technically consideredas chemical technicians, many other techni-


cian jobs involve work in chemistry, forexample, medical technician, pharmacytechnician.

Hazardous materials managers overseethe removal, transportation, and disposal ofharmful materials and substances. The han-dling of hazardous materials requires anunderstanding of the chemistry of thesematerials. Especially important is knowl-edge on the compatibility of differentchemicals, how they move through theenvironment (air, soil, groundwater), andtheir effects on humans and other organ-isms. In addition to courses in chemistry,the academic preparation of hazardousmaterials normally includes courses inenvironmental regulations, toxicology, andindustrial hygiene. Because hazardouswaste management is a relatively new field,only a few colleges offer four-year or grad-uate degrees specifically in this area. Manyschools offer a certificate program requir-ing the completion of a limited number ofcourses in hazardous material management.These certificate programs are often com-pleted by individuals already possessing adegree in science or engineering who arealready working or desire to work as haz-ardous material managers.

Pharmacists provide patient care thatinvolves dispensing drugs, monitoring theiruse, and advising patient and health careworkers about their effects. Although themost obvious role of the pharmacist is as adispenser of drugs, this role is changing tothat of a general health care providerinvolved in education and advising patients.Actually less than half of all pharmacistswork at local pharmacies. The majoritywork at hospitals, universities or researchinstitutes, government agencies, or industry.Currently there are around 200,000 phar-macists in the United States. In the year2000, people in the United States spent $75billion on prescription and nonprescription

drugs. Because the population of the UnitedStates is aging, pharmacy is a growing pro-fession. The education required to becomea pharmacist has changed in recent years;the bachelor of pharmacy degree is beingphased out and students seeking to becomepharmacists will need to complete a doctor-ate in pharmacy program. This is generallya four-year competitive program thatrequires two years of prepharmacy course-work; therefore, a total of six years of post-secondary education will generally berequired to become a pharmacist. Prepara-tory courses in general chemistry, organicchemistry, math, physics, and the humani-ties are needed for entrance into a college ofpharmacy. The pharmacy program builds onthis education with courses in biochemistry,medicinal chemistry, pharmaceutical chem-istry, and pharmaceutics. Pharmacists mustbe licensed to practice in all of the fiftystates. To obtain a license, a prospectivepharmacist must pass a qualifying exam.

Many chemists find employment as sec-ondary or college teachers. Students canprepare themselves as secondary teachersby majoring in chemical education or com-bining a chemistry major with the appropri-ate education courses to become certified toteach. A number of colleges have master’sprograms that enable a person with anundergraduate degree in chemistry tobecome a certified teacher. Many areas ofthe country are experiencing a shortage ofqualified chemistry teachers. Large citiessuch as New York City currently recruitchemistry teachers from foreign countries tohelp fill their shortage. Teaching in a com-munity college requires at least a master’sdegree, but most college positions require adoctorate in chemistry.

Chemical sales and marketing isanother area where chemists are widelyemployed. Knowledge in chemistry allowsa chemical sales representative to better


communicate product information to cus-tomers. It also helps the sale representativeto provide information back to chemicalmanufacturers about customer needs,helping to shape product improvements anddevelopment. In addition to their chemicaltraining, chemical sales representatives ben-efit from courses in business, for example,marketing. A popular career path taken bymany chemists and chemical engineers is toobtain a master’s of business administration(MBA). This degree allows technicallytrained scientists and engineers to move intonumerous management positions withchemical manufacturers.

A field that combines chemistry withcrime investigation is forensic chemistry.Forensic chemists use the tools of analyticalchemistry to analyze samples collected fromcrime scenes. They often work with smallsamples such as blood stains, glass shards,paint chips, hairs, and fabric threads in anattempt to analyze the evidence of a case.Forensic chemists are employed primarilyby government crime labs at the federal,state, and local levels. There are a limitednumber of college programs offeringdegrees in forensic science. Many peopleenter the profession with degrees in chem-istry or other sciences and receive special-ized training from agencies such as the FBIor Drug Enforcement Agency (DEA).

A few of the many employment oppor-tunities available to those with an interest inchemistry have been summarized in thissection. The few examples should illustratethe breadth of possibilities available to indi-viduals who obtain degrees in chemistry andrelated fields. Chemistry majors also branchout into other diverse fields that make use oftheir technical training. Examples includelaw (patent law, environmental law), insur-ance inspectors, journalists, technical writ-ers, and librarians or chemical informationspecialists. Pay for chemical jobs vary

widely depending on geographic location,specialty field, and education required. Inthe year 2000, median salaries for newlygraduated chemistry majors was approxi-mately $33,000 and for chemical engineerswas close to $50,000. Chemical technicianstend to be at the low end of salaries startingin the low to mid-$20,000 range. At theother end of the salary range, pharmacists’starting salaries average about $65,000.

Professional Developmentand the American ChemicalSociety

All scientific disciplines have profes-sional organizations that promote their dis-ciplines and assist their members in theirprofessions. The American Chemical Soci-ety (ACS) is the predominant chemicalorganization in the United States. It isunique because it is the largest scientificsociety in the world with 150,000 members.ACS’s headquarters are in Washington,D.C., where elected officers and a large pro-fessional staff coordinate a myriad of activ-ities. These activities include publishingjournals, conducting meetings, offeringcourses for members, assisting in memberjob placement, providing outreach to thepublic, producing chemistry curriculummaterials, approving college programs, dis-tributing chemical information, and inter-acting with government and industryofficials. ACS has student memberships andapproximately 100 campuses have studentaffiliate organizations.

The American Chemical Society pub-lishes a weekly magazine called Chemical& Engineering News. This publicationinforms ACS members and others on cur-rent news and significant events in chem-istry and related areas. Topics include newadvances in research, awards in chemistry,book reviews, government legislation, per-


spectives on major global issues related tochemistry, and a classified section listingcurrent job openings in education, privateindustry, and government.

One of the roles of ACS is to promotechemical science among young people. Inaddition to ACS’s college program, it coor-dinates a number of activities targeted at ele-mentary and secondary students. Especiallypertinent to this chapter are two publicationswritten for high school students interestedin chemistry: Chemistry and Your Career:Question and Answers and I Know You’re aChemist, But What Do You Do. Informationon how to obtain them can be found under“Student Programs” on the “EducationPage” accessed through the ACS Web site(http://www.acs.org). In this same section ofthe ACS Web page, summaries of differentjob descriptions are given. Look for “Chem-ical Careers in Brief.”

As a professional organization, one ofthe roles of the American Chemical Societyis to present chemistry in a positive light. Topromote the highest standards among itsmembers, the ACS has adopted a code ofconduct. The code presents basic principlesto guide the action and responsibilities ofchemists.

Chemists acknowledge responsibilitiesto:

The PublicChemists have a professional responsibility

to serve the public interest and welfare and tofurther knowledge of science. Chemists shouldactively be concerned with the health and wel-fare of coworkers, consumers, and the commu-nity. Public comments on scientific mattersshould be made with care and precision, withoutunsubstantiated exaggerated, or premature state-ments.

The Science of ChemistryChemists should seek to advance chemical

science, understand the limitations of theirknowledge, and respect the truth. Chemists

should ensure that their scientific contribution,and those of their collaborators, are thorough,accurate, and unbiased in design, implementa-tion, and presentation.

The ProfessionChemists should remain current with devel-

opments in their field, share ideas and informa-tion, and keep accurate and complete laboratoryrecords, maintain integrity in all conduct andpublications, and give due credit to the contri-butions of others. Conflicts of interest and sci-entific misconduct, such as fabrication andplagarism, are incompatible with this Code.

The EmployerChemists should promote and protect legiti-

mate interests of their employers, perform workhonestly and completely, fulfill obligations, andsafeguard proprietary information.

EmployeesChemists as employers should treat subordi-

nates with respect for their professionalism andconcern for their well-being, and provide themwith a safe, congenial working environment, faircompensation, and proper acknowledgment oftheir scientific contribution.

StudentsChemists should regard the tutelage of stu-

dents as trust conferred by society for the pro-motion of the student’s learning and professionaldevelopment. Each student should be treatedrespectfully and without exploitation.

AssociatesChemists should treat associates with respect,

regardless of their level of formal education,learn with them, share ideas honestly, and givecredit for their contributions.

ClientsChemists should serve clients faithfully and

incorruptibly, respect confidentiality, advise hon-estly, and charge fairly.

The EnvironmentChemists should understand and anticipate

the environmental consequences of their work.Chemists have a responsibility to avoid pollutionand to protect the environment.

Modern chemistry was born in the lateeighteenth century. Building on three hun-dred years of tradition, chemists continue to


build models attempting to understand themysteries of matter. This knowledge is notcollected in secrecy and coveted like theancient alchemists. Guided by the principlesembodied in the Code of Conduct, chemistsoperate in the world community and collec-tively have a major influence on all humans.Chemists have developed the ability to mod-ify matter in an amazing number of ways.This has led to improved materials, newcompounds, and even new chemical ele-ments, the basic building blocks of matter.New knowledge is helping us to understandthe chemical basis of life itself, enabling sci-entists to modify life and even create newlife forms. This in turn has raised questionson the role of chemists in modern society.Many benefits have accrued from chemistryin the last three hundred years, and chem-

istry continues to improve the quality of life.As chemists continue to advance the fron-tiers of science, an educated public isrequired to understand these advances alongwith their actual and potential applicationand resultant consequences. Only in thismanner can informed discussion take placethat includes both scientists and nonscien-tists. Chemists, as good citizens, have anobligation to use their knowledge for thebenefit of humankind, and part of this obli-gation is educating the public on their disci-pline. In turn, all citizens have an obligationto understand basic science in order to makeinformed decisions. My hope is that thesepages have assisted you in your understand-ing of chemistry and provided a foundationfor you to use this knowledge to enhanceyour own life, as well as the lives of others.

Glossary

Absolute Zero 0 Kelvins or–273.15°C, thelowest possible temperature, the point on thetemperature scale where all molecularmotion ceasesAcceptable Daily Intake the maximumamount of a substance that can be consumedwithout damaging healthAcid a substance that yields hydrogen ionsin solution or donates protonsAcid Precipitation a collective term forrain, snow, or dry deposition with low pHvaluesActivated Complex momentary intermedi-ate arrangement of atoms when reactants areconverted into products in a chemical reac-tion, also called transition stateActivation Energy minimum energyneeded to initiate a chemical reactionActive how easily a metal is oxidizedActivity Series a ranking of elements inorder of their ability to reduce or oxidizeanother elementAcyclic an open-chained compoundAddition Polymerization bonding ofmonomers without the elimination of atoms,formation of polymer by the bonding ofunsaturated monomersAdhesion attraction between the surface oftwo different bodiesAlcohol organic molecules characterized bycontaining the �OH group

Aldehyde organic molecules characterizedby the �CHO groupAliphatic Hydrocarbon organic hydrocar-bons characterized by a straight chain ofcarbon atomsAlkali Metal group I elements in the peri-odic table: Li, Na, K, Rb, Cs, FrAlkaline Earth Metal group II elements inthe periodic table: Be, Mg, Ca, Sr, BaAlkaloid a nitrogen-containing compoundobtained from plants, for example, caffeine,nicotineAlkane an acyclic saturated hydrocarbonAlkene an acyclic hydrocarbon that con-tains at least one carbon-carbon double bondAlkyl Group the group left when a hydro-gen atom is removed from an alkane, forexample, methane gives the methyl groupAlkylation chemical process in which analkyl group is introduced into an organiccompoundAlkyne an acyclic hydrocarbon that con-tains at least one carbon-carbon triple bondAllotrope different forms of an elementcharacterized by different structuresAlloy a mixture of two or more metals, forexample, zinc � copper � brassAlpha Decay nuclear process in which analpha particle is emitted by the nucleus

336 Glossary

Alpha Particle a radioactive particle con-sisting of two protons and two neutrons,identical to the nucleus of a helium atomAmine organic compounds that result whenone or more hydrogen atoms in ammoniaare replaced by organic radicalsAmino Acid organic acids that contain bothan amino group �NH2 and a carboxylgroup �COOH, the building blocks of pro-teinsAmorphous Solid solids with a randomparticle arrangementAmphoteric a substance that exhibits bothacidic and basic propertiesAmylopectin branched glucose polymercomponent of starchAmylose straight-chain glucose polymercomponent of starchAnabolism metabolic reactions in whichsmaller molecules are built up into largermoleculesAnaerobic without oxygen, condition orprocess that exists when oxygen is absentAnalgesic a drug that relieves painAndrogen a male sex hormoneAnion a negatively charged ionAnneal to heat a metal or glass to a hightemperature and slowly cool reducing brit-tleness of materialAnode the positively charged electrode inan electrolytic cell where oxidation takesplaceAnoxic condition that exists when oxygenis absentAntimatter all particles having comple-mentary properties of matter, for example,positronsAntipyretic substance that reduces feverAromatic Hydrocarbon an unsaturatedcyclic hydrocarbon that does not readilyundergo additional reaction, benzene andrelated compoundsArteriosclerosis condition known as hard-ening of the arteries due to plaque depositsin arteries

Aryl Group an aromatic ring from which asingle hydrogen has been removedAssimilate to incorporate into the bodyAtmosphere envelope of gas surroundingEarth’s surface composed primarily of nitro-gen and oxygenAtomic Number the number of protonscontained in the nucleus of an atom of anelementAvogadro’s Law law stating that equal vol-umes of gases at the same temperature andpressure contain an equal number of parti-clesAvogadro’s Number 6.02 � 1023, the num-ber of particles defining one mole of a sub-stanceBackground Radiation amount of naturalradiation detected in the absence of nonnat-ural radioactive sourcesBase a substance that yields hydroxide ionsin solution or accepts protonsBecquerel SI unit for activity equal to one disintegration per second, abbrevi-ated BqBeta Decay nuclear process in which a betaparticle is emittedBeta Particle in nuclear processes, particleequivalent to an electronBinding Energy the amount of energyholding the nucleus of element together,equivalent to the mass defect � c2

Bioaccumulate the accumulation of achemical in an organismBiological Magnification an increase in theconcentration of a chemical moving up thefood chainBiomolecule molecules associated with life,for example, carbohydrates, fats, proteinsBiosphere total of living and decayingorganismsBipyridyliums class of herbicides thatincludes paraquatBoiling Point Elevation increase in normalboiling point of a pure liquid due to the pres-ence of solute added to the liquid

Glossary 337

Bond Energy the heat of formation of amolecule and its constituent atoms or theenergy necessary to form or break apart amoleculeBose-Einstein Condensate phase of matterthat is created just above absolute zero whenatoms lose their individual identityBoyle’s Law law that states volume of a gasis inversely related to its pressureBreeder Reactor type of nuclear reactorthat creates or “breeds” fissionable pluto-nium from nonfissionable U-238Buckministerfullerene C60, allotrope ofcarbon consisting of spherical arrangementof carbon, named after architect Buckmin-ister Fuller, Invertor of geodesic domeBuffer a solution that resists a change in pHCalcination slow heating of a substancewithout causing it to meltCalorie unit of energy equivalent to theenergy required to raise the temperature ofone gram of water from 14.5°C to 15.5°C,equivalent to 4.18 JCalorimetry experimental technique usedto determine the specific heat capacity of asubstance or various heats of chemical reac-tionsCalx oxide of metal formed during calcina-tion processCarbamates class of insecticides derivedfrom carbamic acidCarbohydrate large class of biomoleculesconsisting of carbon, hydrogen, and oxygenand having the general formula Cx(H2O)y,more specifically a polyhydroxyl ketone orpolyhydroxyl aldhydeCarbonate an ion with the formula CO3

2�

or a compound that contains this ionCarbonate Hardness condition of waterwhen magnesium and calcium are combinedwith carbonates and bicarbonates, alsocalled temporary hardness because it can beremoved by boiling, responsible for scale inpipes

Carbonyl a carbon atom and oxygen atomjoined by a double bond C�O, present inaldehydes and ketonesCarboxyl Functional Group a carbonylwith a hydroxyl (–OH) attached to the car-bon atomCarboxylic Acid group of organic com-pounds characterized by the presence of car-boxyl groupCarcinogen a substance that causes cancerCatabolism metabolic processes wheremolecules are broken downCatalysis chemical process in which a reac-tion rate is acceleratedCatalyst a substance used to accelerate achemical reaction without taking part in thereactionCathode the negatively charged electrode inan electrolytic cell where reduction takesplaceCathodic Protection method where a moreactive metal is connected to a metal struc-ture such as a tank or a ship protecting thestructure because the active metal is oxi-dized rather than the structureCation a positively charged ionCellulase a group of enzymes thathydrolyze celluloseCellulose common carbohydrate that is the primary constituent of cell walls inplantsChapman Reactions series of reactionsresponsible for the destruction of strato-spheric ozoneCharles’s Law law that states volume of agas is directly related to its absolute tem-peratureChemical Affinity the tendency of particu-lar atoms to bond to each otherChemical Change a transformation inwhich one substance changes into another,as opposed to a physical changeChemical Element a pure substance thatcannot be broken down into a simpler sub-stance by chemical means

338 Glossary

Chemical Family a group of elements thatshare similar chemical properties and sharethe same column in the periodic table, forexample, halogens, alkali earthChirality condition that describes the“handedness” of a molecule or whether amolecule exists in forms that can be super-imposed on each otherChlorofluorocarbons also called CFCs,compounds consisting of chorine, fluorine,and carbon that are responsible for strato-spheric ozone destructionCoagulation precipitation or separationfrom a dispersed stateCoefficient of Thermal Expansion mea-sure of the rate at which a substance willexpand when heatedCohesion intermolecular attractive forcebetween particles within a substanceColligative Property a property dependenton the number of particles in solution andnot on the type of particles, for example,boiling point elevation and freezing pointdepressionColloid particles having dimensions of1–1000 nm or a suspension containing par-ticles of this sizeCombination Reaction also called synthe-sis, reaction in which two or more sub-stances combine to form a compoundCombustion burning of a fuel to produceheat, oxidation of a fuel sourceComplete Ionic Equation equation used toexpress reactions in aqueous solutionswhere reactants and products are written asions rather than molecules or compoundsComplete Reaction a chemical reaction inwhich one of the reactants is completelyconsumedComplex Carbohydrate polysaccharidessuch as cellulose and starchCompound a substance consisting of atomsof two or more different elements chemi-cally bonded

Condensation transformation from the gasto liquid stateConjugate Acid in the Brønsted theory, thesubstance formed when a base accepts aprotonConjugate Acid-Base Pair in the Brønstedtheory, an acid and its conjugate base or abase and its conjugate acidConjugate Base in the Brønsted theory, thesubstance that remains when an aciddonates a protonConstructive Interference when wavescombine to reinforce each otherControl to remove or account for the effectof a variable in an experimentControl Rod rods used in nuclear reactor toabsorb neutrons and control fission ofradioactive fuelControlled Experiment method in whichall variables except one independent vari-able are held constant or accounted for insome fashionCoriolis Force pseudo-force introduced toaccount for motion that takes place on arotating Earth, to the right in the NorthernHemisphere and left in the Southern Hemi-sphereCosmic Radiation charged atomic particlesoriginating from spaceCovalent Bond chemical bond in whichelectrons are shared between atomsCovalent Crystal crystal in which atomsare held together by covalent bonds in arigid three-dimensional network, for exam-ple, diamondCracking process by which a compound isbroken down into simpler substances, typi-cally employed in petroleum industry tobreak carbon-carbon bondsCrenation condition that results when cellslose water and shrivel upCritical Mass quantity of fissionable mate-rial necessary to give a self-sustainingnuclear reaction

Glossary 339

Crystalline Solid solid in which atoms arearranged in definite three-dimensional pat-ternCultural Eutrophication accelerated agingof water body due to human influence andwater pollutionCurie measure of radiation activity equal to3.7 � 1010 disintegrations per secondCyclic arrangement of molecule in whichcarbons are bonded together in a ring-patternDalton’s Law of Partial Pressure law that states the total pressure in a mixture of gases is equal to the sum of pressureseach gas would exert in absence of othergasesDaughter the by-product in a radioactivedecayD-Block Element a transition element,valence electrons are in the d orbitalsDecay Activity the rate of decay of aradioactive substanceDecomposition Reaction reaction in whicha compound is broken down into its compo-nentsDelocalized not associated with a particularregion, in metallic bonding it refers to thefact that electrons are shared by the metal asa wholeDenatured to make unsuitable for humanconsumptionDephlogisticated Air term Priestley usedfor oxygen, air devoid of phlogistonDephlogisticated Nitrous Air term Priest-ley used for nitrous oxide, N2O, laughinggasDestructive Interference when waves com-bine to cancel each otherDeuterium isotope of hydrogen containing1 proton and 1 neutronDialysis the separation of particles from acolloid suspension by the passage of sus-pension solution through a semipermeablemembrane

Diffraction the spreading of light or otherwaves after passing through an opening orpassing by an obstacleDihydric an alcohol with two hydroxylgroupsDioxins chlorinated cyclic compoundsDipeptide two amino acids joined a peptidebondDipole a difference in the centers of positiveand negative charge in a moleculeDipole Moment condition when the centersof positive and negative charge in a mole-cule differDipole-Dipole Force type of intermolecu-lar force existing between molecules pos-sessing dipole momentsDisaccharide carbohydrates in which two monosaccharides are bonded to eachotherDispersing Medium in a colloid suspen-sion, the substance in which the particles aresuspendedDispersing Phase in a colloid suspension,the particles that are suspended in the dis-persing mediumDispersion Force intermolecular force thatresults from continuous temporary dipolesformed in molecules not possessing perma-nent dipole momentsDose the amount of substance administeredto an organism during a treatment or exper-imentDose Equivalent a means of comparing dif-ferent forms of radiation on different bodytissueDouble Replacement Reaction a reactionbetween two compounds in which elementsor ions replace each otherEffluent discharge from a pipe or another source, associated with water qual-ityElectrical Potential the electrical potentialenergy of charged body above ground mea-sured in volts

340 Glossary

Electrochemical Cell a cell in which elec-trical energy is used to cause a nonsponta-neous chemical reaction to occurElectrode a terminal that conducts currentin an electrochemical cellElectrolysis process involving the passageof current through an electrochemical cellforcing a nonspontaneous reaction to occurElectrolyte a substance that conducts cur-rent when it is dissolved in waterElectrolyze to bring about a chemicalchange by passing an electrical currentthrough an electrolyteElectromotive Force force due to a differ-ence in electrical potential causing currentto flow in a circuitElectron Configuration distribution ofelectrons into different shells and orbitalsfrom the lower to higher energy levelsElectronegativity measure of the attractionof an element for a bonding pair of electronsElectroplating process where a metal isreduced on to the surface of an object,which serves as the cathode in an electro-chemical cellElectrorefining purification of a metalusing electrolysisElectrostatic charge at restElement see Chemical ElementElementary Particle collectively the small-est units of matter, electrons, protons,quarks, and so on.Elementary Steps single reaction in aseries of reactions that produce a net chem-ical equationEmpirical experimental, to determineexperimentallyEmulsifier a substance that allows twophases to mix and form a solution or sus-pension, used in food processingEndothermic a chemical reaction in whichenergy is absorbed from the surroundingsEnergy the ability to do workEnzymes a biological catalyst

Equilibrium state of chemical reactionwhere the rates of forward and reverse reac-tions are equal, causing concentrations ofreactants and products to remain constantEquilibrium Constant a number equal tothe ratio of the concentration of products atequilibrium over the concentration of reac-tants at equilibrium all raised to a powerequal to the stoichiometric coefficient in thechemical equationEssential Amino Acid an amino acid thatcannot be produced by our bodies and mustbe ingestedEster class of organic compounds thatresults from the reaction of carboxylic acidand an alcoholEstrogen female sex hormonesEther class of organic compound contain-ing an oxygen atom singly bonded to twocarbon atoms C-O-CEutrophication the natural aging of a waterbodyExothermic a chemical reaction that liber-ates energy to the surroundingsFats biomolecules consisting of glycerolsand fatty acids (triglycerides) that are solidat room temperatureFatty Acid long unbranched carboxylic acidchainF-Block Element the lanthanides andactinides, valence electrons in the f orbitalsFeedstock a process chemical used to pro-duce other chemicals or productsFine Chemicals chemicals produced in rel-atively low volumes and at higher prices ascompared to bulk chemicals such as sulfuricacid, includes flavorings, perfumes, phar-maceuticals, and dyesFirst Law of Thermodynamics law thatstates energy in universe is constant, energycannot be created or destroyedFirst Order Reaction reaction in which therate is dependent on the concentration ofreactant to the first power

Glossary 341

Fission see Nuclear FissionFixed Air carbon dioxide, term used byJoseph Black for CO2

Fossil Fuels a fuel derived from the remainsof ancient organisms: oil, natural gas, andcoalFractional Distillation process in which amixture of chemicals is separated by theirdifference in boiling points, used exten-sively to refine petroleum into its differentcomponents or fractionsFree Radical a species possessing at leastone unpaired electronFreezing Point Depression decrease infreezing point of a pure liquid due to thepresence of solute added to the liquidFuel Cell type of electrochemical cell inwhich reactants are continually supplied tocreate electrical energyFullerene pure carbon molecule consistingof at least 60 carbon atoms arranged spher-ically like a soccer ballFunctional Group atom or group of atomsthat characterize a class of compoundsFusion see Nuclear FusionGalvanize a method used to protect metalsby plating them with another metal, forexample, coating iron with zincGamma Ray form of high energy electro-magnetic radiation with no mass and nochargeGamma Decay radioactive process inwhich gamma ray is emitted from thenucleusGay-Lussac’s Law law that states the pres-sure of a gas is directly proportional to thegas’s absolute temperatureGay-Lussac’s Principle law that stateswhen a chemical reaction takes placeinvolving gases, the volumes of reactantsand products are in small whole numberratiosGel colloid in which a liquid is dispersed ina solid

Gray SI unit to measure the amount of radi-ation dose equal to deposition of one jouleof energy into one kilogram of matterGreenhouse Effect warming of the Earthdue to absorption of infrared radiation byparticular atmospheric gases such as H2O,CO2, and CH4

Groups columns or families in the periodictableHaber Process method used to produceammonia from nitrogen and hydrogen athigh temperature and pressure under thepresence of a catalystHalf Reaction the oxidation or reductionreaction in an oxidation-reduction reactionHalf-Life time it takes for substance to bereduced to half its original amount or activ-ity in a chemical or nuclear reactionHalogen family consisting of group 17 ele-ments in the periodic table: F, Cl, Be, I, AtHalogenation chemical reaction involvingthe addition of a halogenHard Water water containing a high con-centration of divalent metal ions particularlycalcium and magnesiumHeat flow of energy from a region of highertemperature to a region of lower temperatureHeat of Fusion energy necessary to changeone kilogram of a substance from a liquid tosolid at the substance’s freezing pointHeat of Vaporization energy necessary tochange one kilogram of a substance from aliquid to gas at the substance’s boiling pointHeating Curve graph that shows how thetemperature of a substance changes as heatis added to the substanceHenry’s Law principle that states theamount of gas dissolved in a solution isdirectly proportional to the partial pressureof the gas above the solutionHeterocyclic Compound a cyclic com-pound in which at least one carbon atom inthe ring has been replaced by another atomor group of atoms

342 Glossary

Homogeneous a solution or mixture that iswell mixed with a constant compositionthroughoutHormone chemical messengers released byglandsHumidity the amount of moisture containedin the air expressed in mass of water vaporper unit volume of airHydration process by which solute parti-cles become surrounded by water moleculesin an aqueous solutionHydrocarbon organic compound consist-ing of only carbon and hydrogenHydrogen Bond intermolecular forceformed between hydrogen of one moleculeand highly electronegative atom such asnitrogen, oxygen, or fluorine on anothermoleculeHydrogenation reaction in which hydrogenis added to unsaturated organic compoundsin presence of a catalyst, used to turn liquidvegetable oils into solidsHydrolysis decomposition or change of asubstance by its reaction with waterHydronium the H3O

� ionHydrosphere the sum total of Earth’s waterincluding the oceans, polar ice caps,groundwater, atmospheric water, rivers, andlakesHyperbaric a condition of elevated pres-sureHypothesis a tentative statement used toexplain some aspect of nature and testedthrough an experimentIatrochemical medical school of thoughtthat attributed disease to chemical imbal-ances in the bodyIdeal Gas gas in which it is assumed gasmolecules occupy no appreciable portion ofthe total volume, do not interact with oneanother, and undergo elastic collision withone anotherIdeal Gas Law relationship describing thebehavior of an ideal gas, PV � nRT whereP is the pressure, V is volume, n is number

of moles, R is ideal gas law constant, and Tis absolute temperatureImmiscible when two liquids do not mix,for example, oil and waterIndependent Variable the factor which aresearcher varies in an experimentIndicator a substance that exhibits a colorchange dependent on pHIndustrial Hygiene dealing with the healthand safety of workers, a discipline that dealswith the maintenance of a healthy workenvironmentInhibitor a substance that slows down orstops a reactionInorganic Chemistry chemistry of com-pounds excluding hydrocarbons and othercompounds based on carbonIntermolecular between moleculesInternal Energy the total potential andkinetic energy of all particles comprising asystemIntramolecular within a moleculeInvert Sugar a sugar mixture consisting ofglucose and fructoseIodine Number measure of the amount of iodine absorbed by an unsaturated fat and used to measure the degree of satura-tionIon a charged species created when an atomor group of atoms gains or loses electronsIon-Dipole Force intermolecular forcebetween an ion and a dipoleIon Pair in a solution when a positive andnegative ion exist as a single particleIon Product Constant in an ionic reaction,the product of each ion’s concentration insolution raised to a power equal to the coef-ficient in the net ionic equation, for water itequals �H� �OH� Ionic Bond intramolecular force createdwhen electrons are transferred from oneatom to another creating ions that possesselectrostatic attraction for one anotherIonic Solid a solid composed of anions andcations such as NaCl

Glossary 343

Ionization process by which an atom ormolecule gains or loses electrons andacquires a charge becoming an ionIonizing Radiation alpha, beta, or gammaradiation possessing sufficient energy todislodge electrons from atoms in tissueforming ions that can cause cell damageIsomer different compounds that have thesame molecular formulaIsotope forms of the same element that havedifferent mass numbersJoule SI unit for energy and work equiva-lent to 0.24 calorieKetone class of organic compounds con-taining a carbonyl group bonded to twoother carbon atomsKilojoules unit of energy equal to 1,000joulesKinetic Energy energy associated withmotion equal to 1⁄

2mv2

Kinetic Molecular Theory model thatdefines an ideal gas and assumes the aver-age kinetic energy of gas molecules isdirectly proportional to the absolute tem-peratureKinetics that area of chemistry that dealswith how fast reactions take placeLatent Heat energy transferred to materialwithout a change in temperature, associatedwith phase changes such as melting andboilingLattice Points positions in a unit cell occu-pied by atom, molecules, or ionsLaw of Definite Proportion law that statesthat different samples of the same com-pound always contain elemental mass per-centages that are constantLaw of Mass Action mathematical expres-sion based on the ratio between productsand reactants at equilibrium, an equation todetermine the equilibrium rate constantLaw of Multiple Proportions law thatstates when two elements combine to formmore than one compound that the mass ofone element compared to the fixed mass of

the other element is in the ratio of smallwhole numbersLe Chatelier’s Principle principle that sayswhen a system is at equilibrium and achange is imposed on the system that thesystem will shift to reduce the changeLewis Electron Dot Formula a diagramshowing how the valence electrons are dis-tributed around an atom or distributed in amoleculeLipid a class of biomolecules including fatsand oils characterized by their insolubilityin waterLiquid Crystals materials that have prop-erties of both solids and liquids used exten-sively in digital displaysLithosphere outer surface of Earth includ-ing the crust and upper mantleLock-and-Key Model model to explainhow enzymes catalyze reactions with specific enzymes acting as locks that onlycertain substrates which act as keys can fitLondon Force intermolecular force thatarises when momentary temporary dipolesform in nonpolar moleculesLong-Range Order general term todescribe regular repeating structure found incrystalsLongwave Radiation infrared radiationemitted by EarthMain Group Element elements in groups1, 2, and 13–18, members in each grouphave the same general electron configurationMass Defect the change in mass when anucleus is formed from its constituentnucleonsMass Number sum of protons and neutronsin the nucleus of an atomMatter the material that comprises the uni-verse, anything that has mass and occupiesspaceMaximum Contaminant Level the maxi-mum concentration of a substance dissolvedin water deemed safe for consumption

344 Glossary

Meniscus the curved upper surface of a nar-row column of waterMetabolism the total of chemical reactionsand processes used by organisms to trans-form energy, reproduce, and maintain them-selvesMetal class of elements characterized bytheir ability to form positive ions, conductheat and electricity, malleability, and lusterMetallic Bond bond present in metals dueto delocalized electrons moving throughoutthe metal latticeMetallic Solid type of solid characterizedby delocalized electrons and metal atomsoccupying lattice pointsMetalloid elements have properties inter-mediate between metals and nonmetalsMixture combination of two or more sub-stances where the individual substancesmaintain their identityModerator a material such as graphite ordeuterium used to slow down neutrons innuclear reactorsMolality a way to express concentration ofa solution, moles of solute per kilogram ofsolventMolarity a way to express concentration ofa solution, moles of solute per liter of solu-tionMole number of C-12 atoms in exactly 12grams of C-12 equal to 6.022 � 1023, aquantity of substance equal to the molarmass expressed in gramsMolecular Equation a reaction in whichthe reactants and products are expressed as molecules or whole units rather than asionsMolecular Orbital Theory a model thatuses wave functions to describe the positionof electrons in a molecule, assuming elec-trons are delocalized within the moleculeMolecular Solid a solid that contains mol-ecules at the lattice pointsMolecule a group of atoms that exist as aunit and are held together by covalent bonds

Monomer discrete molecules that jointogether to form a polymerMonosaccharide a simple sugar moleculecontaining only a single aldehyde or ketonegroup, for example, glucoseNanometer one billionth of a meter, 10�9 mNebula cloud of dust and gases, principallyhydrogen, distributed throughout the uni-verseNet Ionic Equation a chemical equationthat shows only the ionic species that actu-ally take part in the reactionNeutralization process that occurs when anacid reacts with a base, a type of reactioninvolving an acid and baseNewton SI unit for force equal to 1 kg-m/s2

Nonelectrolyte a substance that does notconduct current when it is dissolved in waterNonionizing Radiation electromagneticradiation with insufficient energy to dis-lodge electrons and cause ionization inhuman tissue, for example, radio waves,microwave, visible lightNonmetal elements found on the right sideof the periodic table that conduct heat andelectricity poorlyNonpoint Source Pollution pollution thatdoes not originate from one specific loca-tion, such as a sewer pipe, but comes frommultiple locations spread over a wide area,for example, runoff from city streetsNonrenewable Resource a resource in fixedsupply or one that it is replenished at a rateso slow that it is exhausted before it isreplenished, for example, oilNormal Alkane an alkane in which all thecarbon atoms in the molecule are attachedin a continuous chainNormality a way to express the concentra-tion of a solution, the number of equivalentsper liter of solutionNuclear Fission the splitting of the nucleusof an atom of a heavier element into smallernuclei to produce energy in a nuclear reac-tion

Glossary 345

Nuclear Fusion the combination of atomicnuclei of lighter elements into heavier nucleito produce energyNucleic Acid complex biomoleculesresponsible for passing on genetic informa-tion and for protein synthesis, include DNAand RNANucleon a proton or a neutron, the numberof nucleons in an atom equals the sum ofprotons and neutrons in the nucleusOctet Rule general rule that states that themost stable electron configuration occurswhen an atom surrounds itself with eightvalence electronsOil biomolecules consisting of glycerolsand fatty acids (triglycerides) that are liquidat room temperatureOil of Vitriol concentrated sulfuric acidOptical Isomer isomers that differ in theirability to rotate light in opposite directionsOrganic Chemistry chemistry of carbon-based compoundsOrganochloride chlorine substituted hydro-carbon molecule widely used in industry,pesticides, CFCsOrganophosphate organic compoundscontaining phosphorus, widely used in cer-tain pesticidesOsmosis process in which a solvent flowsfrom a diluted solution to a more concen-trated solution across a semipermeablemembraneOsmotic Pressure pressure required to stopthe movement of the solvent across a semi-permeable membrane during osmosisOxidation process that describes the loss ofelectronsOxidation Number the amount of chargetransferred by or to an atom when it isassumed an ionic bond is formed as elec-trons are donated to the more electronega-tive atomOxidation Reaction a chemical reaction inwhich an atom, molecule, or ion loses elec-trons

Oxidizing Agent a substance that is reducedand hence causes oxidation to occur in aredox reactionParent in a nuclear reaction the originalsubstance that emits some form of radiationPartial Pressure pressure exerted by anindividual gas in a mixture of gasesParts Per Million way to express concen-tration by giving the number of units ofmass or volume out of one million units ofmass or volumePascal SI unit for pressure equal to 1 new-ton per square meterPeptide a sequence of 50 or fewer aminoacids joined by peptide bondsPeptide Bond bond formed between thecarboxyl group of one amino acid and theamino group of another amino acid in a pep-tidePercent by Mass proportion of the mass ofeach element in a compound expressed as apercent or proportion of solute mass as totalmass of solutionPerfect Gas see Ideal GasPeriod a horizontal row in the periodic tablePermanent Hardness condition of waterwhen magnesium and calcium are combinedwith chlorides and sulfates rather than car-bonates, cannot be removed by heatingPetrochemical chemicals derived from oilor natural gaspH the negative of the base 10 logarithm ofthe molar concentration of the hydrogen ionconcentration of a solution, –log�H� Phenol C6H5OH, the molecule formedwhen a hydroxyl group is substituted for ahydrogen in benzene or a compound thatcontains the phenol groupPhenyl C6H5, group formed when a hydro-gen is removed from benzenePheromones compounds produced in ani-mals, especially insects, to communicatemessages to members of the same speciesPhilosopher’s Stone imaginary substancethat alchemists believed could turn base

346 Glossary

metal into gold or was an elixir to cure dis-easesPhlogiston a material once thought to be anelement responsible for combustionPhospholipid lipid containing phosphorusderived from phosphoric acidPhotochemical Oxidants air pollutantsproduced when hydrocarbons, nitrogenoxides, and other chemicals react under theinfluence of sunlight, for example, ozone,peroxyacylnitrates (PAN)Physical Changes transformations in thenatural world, such as freezing and boiling,that do not involve a change in substancesor in which a chemical reaction does nottake placePhytoplankton marine and freshwaterplants that float in the surface waters, manyare microscopic and include different formsof algaePlasma state of matter in which atoms ormolecules have been ionized to form posi-tive ions and electronsPlasmolysis condition that results whencells absorb water and rupturePneumatic Chemistry study of gases andair, important in the development of modernchemistryPneumatic Trough device used to collectand measure the volume of gases inventedin 1700s by Stephen Hales, consisted ofglass container submerged in water in whichgas displaced waterpOH the negative of the base 10 logarithmof the molar concentration of the hydroxideion concentration of a solution, –log�OH� Point Source Pollution pollution originat-ing from a specific location, for example, asmokestackPolar Covalent Bond a covalent bond inwhich the bonding pair of electrons are notshared equally between atoms, the atomwith the greater electronegativity has agreater affinity for the bonding electrons

Polar Molecule a molecule where the cen-ters of positive and negative charge differ,creating a permanent dipole momentPolarizability ability of an electron cloud ina neutral atom to be distortedPolarized Light light in which the electro-magnetic wave vibrates in only one planePolyatomic Ion an ion consisting of morethan one atomPolymer giant molecule formed by the link-ing of simple molecules called monomersPolypeptide a chain of more than 50 aminoacidsPolysaccharide a carbohydrate polymerconsisting of monosaccharides bondedtogetherPolyunsaturated Fatty Acid fatty acid con-taining multiple carbon-carbon doublebondsPositron a subatomic particle with the massof an electron and a charge of �1Potential Energy the ability to do work byvirtue of the position or arrangement of asystemPressure force per unit areaPrimary Amine an amine in which thenitrogen atom is bonded to a single carbonatom, when only one hydrogen in ammoniahas been replaced by an organic R groupPrimary Cells batteries that cannot berechargedPrimary Production conversion of matterinto biomass using sunlight by plants in anecosystemPrimary Structure the sequence of aminoacids bonded to each other in a peptidechainPrincipal Quantum Number a numberthat designates an electron’s energy leveland its average distance from the nucleus ofan atomProducts the resulting substances in achemical reactionProgestin pregnancy hormones

Glossary 347

Protein biomolecules consisting of poly-peptide chain with large molecular massProtostar early stage in the formation of astar when gases and dust start to contractdue to gravitational forcesP-V Work work associated with the expan-sion or compression of a gasPyrethroids synthetic forms of pyrethrins,insecticides based on extracts from chrysan-themumsQualitative Analysis general method usedto determine the composition of a com-poundQuality Control a system for maintainingproper standards in a product or collectionof dataQuantum Mechanics theory that explainsthe behavior of matter using wave functionsto characterize the energy of electrons inatomsQuark a fundamental particle hypothesizedto form the basic building blocks of all matter; there are six different quark types:up, down, strange, charm, top, and bot-tomQuaternary Structure describes the asso-ciation of the different chains in multichainproteinsRad a unit for an absorbed dose of radiationequal to 100 ergs of energy absorbed by agram of matterRadioactive unstable atomic nuclei sponta-neously emitting particles and energyRadioactive Tracer a radioactive substanceused to monitor the movement and behaviorof a chemical in biological processes andchemical reactionsRate Determining Step the slowest reac-tion in the reaction mechanismRate Law general mathematical expressionthat gives the rate of reaction, depends onrate constant and concentrations of reactantsReactants substances initially present in achemical reaction

Reaction Mechanism series of reactionsthat shows how reactants are converted intoproducts in a chemical reactionRedox Reaction reaction involving thetransfer of electrons, an oxidation-reductionreactionReducing Agent a substance that is oxi-dized and causes reduction to occur in aredox reactionReduction process that describes the gainof electronsReduction Reaction a chemical reaction inwhich an atom, molecule, or ion gains elec-tronsRelative Humidity ratio of the vapor pres-sure of water in air compared to the saturatedvapor pressure of pure water at the sametemperature, measure of the amount of watervapor in an air mass expressed as percent ofhow much water vapor that air mass can holdRelative Mass mass measured with respectto a standard, atomic masses based on C-12standardRem unit used to measure radiation doseequivalents based on the quantity of radia-tion absorbed and type of radiationResidual Chlorine chlorine that has notcombined with organic matterResins sticky, liquid organic substancesexuded from plants that harden upon expo-sure to airResonance Structures any of two or morestructures used to represent the overallstructure and behavior of a compound eventhough the compound does not exist in anyof the resonance formsReverse Osmosis process in which pressureis used to force a solvent of a concentratedsolution through a semipermeable mem-brane toward a diluted solution, often usedto purify waterReversible Reaction a reaction that occursin both directions: reactants to products andproducts to reactants

348 Glossary

Rod component of nuclear reactor holdingfuel (fuel rod) or substances used to controlnuclear reaction (control rod)Saccharides carbohydratesSacrificial Anode in cathodic protection,the metal connected to the structure to beprotected that is more readily oxidized thanthe structureSalt Bridge concentrated solution of elec-trolyte used to complete the circuit in anelectrochemical cell that helps to equalizecharge distribution in each half cellSaltpeter potassium nitrate, KNO3

Saponification conversion of a fat to soapby reacting with an alkaliSaturated solution that contains the maxi-mum amount of solute under a given set ofconditionsSaturated Fatty Acid fatty acids in whichthe carbon chain contains only single carbon-carbon bondsSaturated Hydrocarbon a hydrocarboncontaining only single carbon-carbon bondsScintillation process producing discreteflashes or sparks of lightScrubbing process used to remove air pol-lutants from industrial emissionsSecondary Amine an amine in which thenitrogen atom is bonded to two carbon atoms,when two hydrogen atoms in ammonia havebeen replaced by two organic R groupsSecondary Cell batteries capable of beingrechargedSecondary Production conversion of plantmatter into biomass by organisms feeding atthe second trophic level and aboveSecondary Structure designates howatoms in a protein’s backbone structure arearranged in spaceSemipermeable Membrane media thatallows the passage of solvent molecules butblocks solute moleculesShortwave Radiation radiation emitted bythe Sun that passes through and interactswith the atmosphere

Simple Carbohydrate monosaccharidessuch as glucose, fructose, and galactose thatcannot be broken down by waterSingle Replacement Reaction type ofchemical reaction in which one elementreplaces another in a compoundSmelting to melt an ore in order to refine ametal from the oreSolute in a solution the component presentin the smaller amountSolution a homogeneous mixture of two ormore substancesSolvation process by which a solute dis-solves in a solventSolvay Process method used to producesodium carbonate using sodium chloride,ammonia, and carbon dioxideSolvent in a solution the component presentin the larger amountSpecific Heat Capacity the amount of heatrequired to raise the temperature of anobject by 1 degree CelsiusSpectator Ions ions that do not take part ina chemical reactionStandard Hydrogen Electrode standardelectrode consisting of hydrogen gas at 1atmosphere pressure, a platinum electrode,and 1 molar hydrochloric acid, the potentialof this electrode is defined as 0 and is usedto determine the potential of other half-reactionsStandard Reduction Potential the voltagemeasured for a half-cell under standard con-ditions when in reference to the standardhydrogen electrodeSteady-State a system that does not changewith timeStereoisomer isomers in which the atomsdiffer in their three-dimensional spatialarrangementSteroids lipids that contain four fused car-bon rings, three containing six carbons andone containing five carbons, include choles-terol and the sex hormonesSterols see Steroids

Glossary 349

Stoichiometry mass relationship in a chem-ical reactionStraight Chain hydrocarbon structure inwhich carbon is bonded in a linear fashionas opposed to branchingStrong Electrolyte a substance that is agood conductor of electricity when dis-solved in waterStrong Nuclear Force force responsible forbinding the nucleons of an atom’s nucleustogetherStructural Isomer compounds that havethe same formula but have different atomicarrangementSubcritical Mass when the quantity ofradioactive fuel is insufficient to produce aself-sustaining chain reactionSublimation process where a substancepasses directly from the solid to gaseousphase without going through the liquidphaseSubstrate the compound an enzyme actsuponSupercritical Mass when the quantity of aradioactive fuel is sufficient to produce aself-sustaining chain reactionSupernova massive explosion of giant starat the end of its lifeSupersaturated a solution containing moresolute than is present in its saturated stateSurface Tension property of liquid causingit to contract to the smallest possible areadue to the unbalance of forces at the surfaceof the liquidSurroundings the remainder of the universeoutside a defined systemSuspension a mixture of nonsettling solidparticles in a liquid, that is, gels and aerosolsSynthesis Reaction see combination reac-tionSystem that part of the universe of concernto usTemper heat treatment of metalsTemporary Hardness see Carbonate hard-ness

Tertiary Amine an amine in which thenitrogen atom is bonded to three carbonatoms, when three hydrogen atoms inammonia have been replaced by threeorganic R groupsTertiary Structure three-dimensionalstructure in protein resulting from interac-tion of amino acids in coiled chainThermochemistry area of chemistry thatdeals with energy changes that take place inchemical processesThermoplastics plastics that can be repeat-edly heated and reformedThermosetting Plastics type of plastics thatharden into final form after being heatedonceTrace Gas gaseous components of theatmosphere that occur in low concentra-tions, for example, methane, ozoneTransmutation nuclear transformation inwhich the atomic nucleus of one element isconverted into another elementTriazine group of herbicides that work bydisrupting photosynthesisTriglyceride principal component of fatsand oils formed from three molecules offatty acids and one glycerol moleculeTrihalomethane class of suspected car-cinogens that form when chlorine and otherhalogens react with humic acids, chloro-form, and bromoformTrihydric an alcohol with three hydroxylgroupsTyndall Effect scattering of light by fineparticles in a suspensionUnit Cell most basic repeating unit of a lat-tice structureUnsaturated solution that has not reachedthe point of saturation and can hold moresoluteUnsaturated Hydrocarbon a hydrocarbonthat contains at least one double or triplebondValence a positive number indicating theability of an atom or radical to combine with

350 Glossary

other atoms or radicals, the number of elec-trons an atom can donate or accept in achemical reactionValence Bond Theory theory of bondingbased on overlapping valence orbitalsValence Electron Configuration quantumnumbers of electrons that reside in the out-ermost shell of an atomVan Der Waal Force intermolecular forcethat can include dipole-dipole, ion-dipole,or London forceVan’t Hoff Factor ratio of the number ofactual particles dissolved in solution to thetheoretical number of particles predictedwhen a substance dissolvesVapor the gaseous state of a substancethat normally exists in the solid or liquidphaseVapor Pressure pressure exerted by a liq-uid or solid’s vapor when the liquid or solidis in equilibrium with its vaporVaporization evaporation, process thatoccurs when molecules pass from liquid togaseous stateVariable one of any number of factors thatchange and may influence an observation orexperimental outcomeViscosity the resistance to flowVitalism theory that a vital force associatedwith life was associated with all organicsubstances

VSEPR Model valence shell electron pairrepulsion model, model used to predict thegeometry of molecule based on distribu-tion of shared and unshared electron pairsdistributed around central atom of a mole-culeVulcanization heating rubber in the pres-ence of sulfur to remove its tackiness andimprove its qualityWeak Electrolyte a substance that is a poorconductor of electricity when dissolved inwaterWhite Dwarf old star that no longer under-goes thermonuclear reaction but continuesto radiate lightWien’s Law law that states the wavelengthof maximum energy radiation from anobject is inversely proportional to the sur-face temperature of the objectWork product of force times displace-mentX-Ray Crystallography method used todetermine structure of crystal by examiningdiffraction pattern produced by x-raysdirected at crystalZooplankton marine and freshwater ani-mals that float in the surface waters, manyare microscopicZwitterion a molecule with a negativecharge on one atom and a positive charge onanother atom

Brief Timeline of Chemistry

30000 B.C. cave paintings

8000 B.C. agriculture develops

6000 B.C. copper smelting

4000 B.C. copper alloys

2000 B.C. bronze

600 B.C. Anaximander posits air as primary element; Thasles posits water as primaryelement

450 B.C. Empedocies posits air, earth, fire and water as four major elements

400 B.C. Democritus leads atomists school, atoms basic form of matter

350 B.C. Aristotle’s Meterologica

A.D. 300 Emperor Diocletian outlaws chemistry in Roman Empire

700–1100 height of Arab period; work of Geber, Al Razi, Avicenna

1000 gunpowder invented in China

1250 Albertus Magnus’s alchemical studies

1250 Roger Bacon studies on gunpowder

1525 Paracelsus uses chemistry in medicinal treatment

1540 Pirotechnia, Biringuccio’s work on metallurgy

1556 Agricola’s Of Metallaurgy, treatise on mining

1620 Van Helmont’s concept of gas

1640 Torricelli invents barometer

1660 Boyle’s law

352 Brief Timeline of Chemistry

1661 Boyle publishes The Skeptical Chymist

1670 Becher and Stahl lay groundwork for phlogiston theory

1750 Hales invents pneumatic trough

1766 Cavendish isolates hydrogen

1770 Priestley and Scheele isolate oxygen

1789 Lavoisier’s Elements of Chemistry

1794 Lavoisier beheaded

1770–1800 Volta studies electricity

1802 DuPont chemical company founded

1808 Dalton proproses atoms as basic building blocks of matter

1808 Humphrey Davy uses electrolysis to discover metallic elements

1810 Avogadro’s hypothesis

1828 Wohler synthesizes urea

1830 Faraday’s studies on electricity

1845 Kolbe synthesizes acetic acid

1846 ether is used as anesthetic

1852 idea of valence is introduced by Frankland

1856 Perkin perfects synthetic dyes

1859 first commercial oil well is drilled by Drake in Titusville, Pennsylvania

1863 Solvay process is discovered for making sodium bicarbonate

1865 Kekule determines structure of benzene

1866 Nobel invents dynamite

1870 Mendeleev’s periodic table

1876 American Chemical Society is founded

1884 Arrhenius’s theory that acids produce hydrogen ions

1893 synthesis of aspirin by Hoffman

1895 Roentgen discoveres x-rays

1896 Becquerel discovers process of radioactivity

1897 Thomson discovers electron

1908 Rutherford’s model of the atom

1910 rayon production begins in U.S.A.

Brief Timeline of Chemistry 353

1913 Bohr’s model of hydrogen atom

1919 Rutherford’s alpha particles support existence of protons

1920 covalent bond by Lewis and Langmuir

1923 Brønsted definition of acid and base

1927–30 quantum mechanical model of atom developed

1932 neutron discovered by Chadwick

1931 DuPont produces synthetic rubber neoprene

1936 Carothers’ DuPont team produces nylon

1939 Muller synthesizes DDT

1940 synthetic rubber is used in tires

1953 Crick and Watson present structure for DNA

1963 Rachael Carson publishes Silent Spring

1969 EPA is established

1976 evidence that chlorine compounds destroy ozone

1985 Smalley and Kroto discover buckminsterfullerenes

1989 Human Genome Project commences

1990 PCR method established

1996 cloning of Dolly the sheep

2000 conclusion of human genome sequencing

Nobel Laureates in Chemistry

Nobel laureates are listed consecutively by year starting with 1901. The country indicateswhere the chemist did major work. If the chemist’s country of birth is different than wheremajor work is done, this is indicated after the “b.” In many years, several chemists arerecognized. When several chemists are recognized in the same year, it may be either forwork in similar areas of chemistry, working collaboratively or independently, or for work inentirely different areas. When several laureates have been recognized for work in similarareas, the work is cited for the first laureate in the list.

Year Recipient(s) Country Work

1901 Jacobus Henricusvan’t Hoff

Germanyb. Netherlands

laws of chemical dynamics and osmotic pres-sure in solutions

1902 Hermann Emil Fi-scher

Germany sugar and purine syntheses

1903 Svante August Ar-rhenius

Sweden electrolytic dissociation theory

1904 William Ramsay Great Britain discovery of noble gases1905 Adolf von Baeyer Germany organic dyes and hydroaromatic compounds1906 Henri Moissan France isolation of fluorine and electric furnace1907 Eduard Buchner Germany fermentation in absence of cells and biochem-

istry1908 Ernest Rutherford Great Britain

b. New Zealandradioactive decay

1909 Wilhelm Ostwald Germany b. Russia chemical equilibrium, kinetics, and catalysis1910 Otto Wallach Germany pioneering work with alicyclic compounds1911 Marie Curie France b. Poland discovery of radium and polonium1912 Victor Grignard France discovery of Grignard’s reagent

Paul Sabatier France hydrogenation with metal catalysts1913 Alfred Werner Switzerland

b. Germanybonding of inorganic compounds

1914 Theodore Richards United States determination of atomic weights1915 Richard Willstatter Germany studies of plant pigments, especially chloro-

phyll

356 Nobel Laureates in Chemistry


1918 Fritz Haber Germany synthesis of ammonia1920 Walter H. Nernst Germany thermochemistry1921 Frederick Soddy Great Britain radioactive substances and isotopes1922 Francis W. Aston Great Britain mass spectrography and discovery of isotopes1923 Fritz Perl Austria organic microanalysis1925 Richard A.

ZsigmondyGermanyb. Austria

colloid chemistry

1926 Theodor Svedberg Sweden disperse systems1927 Heinrich Otto Wie-

landGermany bile acids

1928 Adolf Otto Windaus Germany sterols relationship with vitamins1929 Arthur Harden Great Britain fermentation of sugar and sugar enzymes

Hans von Euler-Chelpin

Sweden b. Germany

1930 Hans Fischer Germany synthesis of hemin1931 Carl Bosch Germany high pressure chemical processing

Friedrich Bergins Germany1932 Irving Langmuir United States surface chemistry1934 Harold Urey United States discovery of heavy hydrogen1935 Frederic Joliot France synthesis of new radioactive elements

Irene Joliot-Curie France1936 Peter Debye Germany b.Netherlands dipole moments and x-ray diffraction1937 Walter Norman Hay-

worthGreat Britain carbohydrates and vitamin C

Paul Karrer Switzerland vitamins A and B121938 Richard Kuhn Germany b. Austria carotenoids and vitamins1939 Adolf F. J.

ButenandtGermany sex hormones

Leopold Ruzicka Switzerlandb. Hungary

polymethylenes and terpenes

1943 George DeHevesy Sweden b. Hungary isotope tracers1944 Otto Hahn Germany fission of heavy nuclei1945 Atturi I. Virtanen Finland agricultural and food chemistry and preserva-

tion of fodder1946 James B. Sumner United States crystallation of enzymes

John H. Northrop United States preparation of proteins and enzymes in pureform

Wendell M. Stanley United States1947 Robert Robinson Great Britain alkaloids1948 Arne W. K. Tiselius Sweden electrophoresis and serum proteins1949 William F. Giauque United States low temperature thermodynamics1950 Otto Diels Germany diene synthesis

Kurt Alder Germany1951 Edwin McMillan United States chemistry of transuranium elements

Glenn Seaborg United States



1952 Archer J. P. Martin Great Britain invention of partition chromatographyRichard L. M. Synge Great Britain

1953 Hermann Staudinger Germany macromolecular chemistry1954 Linus Pauling United States chemical bonding and molecular structure of

proteins1955 Vincent Du

VigneaudUnited States sulfur compounds of biological importance,

synthesis of polypeptide hormone1956 Cyril Hinshelwood Great Britain mechanisms of chemical reaction

Nikolay Semenov USSR1957 Alexander Todd Great Britain nucleotides and their co-enzymes1958 Frederick Sanger Great Britain protein structure, insulin1959 Jaroslav Heyrovsky Czechoslovakia polarographic methods of analysis1960 Willard Libby United States carbon-14 dating1961 Melvin Calvin United States CO2 assimilation in plants1962 Max Preutz Great Britain

b. Austriastructure of globular proteins

John Kendrew Great Britain1963 Giulio Natta Italy high polymers

Karl Ziegler Germany1964 Dorothy Crowfoot

HodgkinGreat Britain x-ray techniques of the structures of biochem-

ical substances1965 Robert Woodward United States organic synthesis1966 Robert Mulliken United States chemical bonds and electronic structure of

molecules1967 Manfred Eigen Germany study of very fast chemical reactions

Ronald Norrish Great BritainGeorge Porter Great Britain

1968 Lars Onsager United Statesb. Norway

thermodynamic of irreversible processes

1969 Derek Barton Great Britain conformationOdd Hensel Norway

1970 Luis Leloir Argentina sugar nucleotides and carbohydrate biosynthe-sis

1971 Gerhard Herzberg Canada b. Germany structure and geometry of free radicals1972 Christian Anfinsen United States ribonuclease, amino acid sequencing and bio-

logical activityStanford Moore United States chemical structure and catalytic activity of ri-

bonucleaseWilliam Stein United States

1973 Ernst Fischer Germany organometallic sandwich compoundsGeoffrey Wilkinson Great Britain

1974 Paul Flory United States macromolecules1975 John Cornforth Great Britain stereochemisitry of enzyme-catalyzed reac-

tionsVladmir Prelog Switzerland b. Bosnia stereochemistry of organic molecules

1976 William Lipscomb United States structure of borane and bonding

358 Nobel Laureates in Chemistry


1977 Ilya Prigogine Belgium b. Russia theory of dissipative structures1978 Peter Mitchell Great Britain chemiosmotic theory1979 Herbert Brown United States

b. Great Britainorganic synthesis of boron and phosphoruscompound

George Wittig Germany1980 Paul Berg United States recombinant DNA

Walter Gilbert United States nucleic acid base sequencesFrederick Sanger Great Britain

1981 Kenichi Fukui Japan chemical reactions and orbital theoryRoald Hoffman United States

b. Poland1982 Aaron Klug Great Britain crystallographic electron microscopy applied

to nucleic acids and proteins1983 Henry Taube United States

b. Canadaelectron transfer in metal complexes

1984 Robert Merrifield United States chemical synthesis1985 Herbert Hauptman United States crystal structures

Jerome Karle United States1986 Dudley Herschbach United States chemical elementary processes

Yuan Lee United Statesb. Taiwan

John Polanyi Canada1987 Donald Cram United States development of molecules with highly selec-

tive structure specific interactionsJean-Marie Lehn FranceCharles Pedersen United States

b. Korea, of Norwegiandescent

1988 Johann Deisenhofer Germany and UnitedStates b. Germany

photosynthesis

Robert Huber GermanyMichel Hartmut Germany

1989 Sidney Altman United Statesb. Canada

catalytic properties of RNA

Thomas Cech United States1990 Elias James Corey United States organic synthesis1991 Richard Ernst Switzerland nuclear resonance spectroscopy1992 Rudolph Marcus United States

b. Canadaelectron transfer in chemical systems

1993 Kary Mullis United States invention of PCR methodMichael Smith Canada

b. Great Britainmutagenesis and protein studies

1994 George Olah United Statesb. Hungary

carbocation chemistry



1995 Paul Crutzen Germanyb. Netherlands

atmospheric chemistry

Mario Molina United Statesb. Mexico

stratospheric ozone depletion

Rowland F.Sherwood

United States

1996 Robert Kurl United States discovery of fullerenesHarold Kroto Great BritainRichard Smalley United States

1997 Paul Boyer United States enzyme mechanism of ATPJohn Walker Great BritainJens Skou Denmark discovery of ion transport enzyme Na�, K�-

ATPase1998 Walter Kohn United States density function theory

John Pople Great Britain computational methods in quantum chemistry1999 Ahmed Zewail United States

b. Egypttransition states using femto spectroscopy

2000 Alan J. Heeger United States discovery of conductive polymersAlan MacDiarmid United StatesHideki Shirakawa Japan

2001 William Knowles United States chirally catalyzed hydrogenation reactionsRyoji Noyori JapanK. Barry Sharpless United States chirally catalyzed oxidation reactions

2002 John B. Fenn United States identical of biological macromoleculesKoichi Tanaka JapanKurt Wuthrich Switzerland NMR analysis of biological macromolecules

Table of the Elements

Element SymbolAtomicNumber

AtomicMass Element Symbol

AtomicNumber

AtomicMass

Actinium Ac 89 227 Mercury Hg 80 200.6Aluminum Al 13 26.98 Molybdenum Mo 42 95.94Americium Am 95 243 Neodymium Nd 60 144.2Antimony Sb 51 121.8 Neon Ne 10 20.18Argon Ar 18 39.95 Neptunium Np 93 237Arsenic As 33 74.92 Nickel Ni 28 58.69Astatine At 85 210 Niobium Nb 41 92.91Barium Ba 56 137.3 Nitrogen N 7 14.01Berkelium Bk 97 247 Nobelium No 102 259Beryllium Be 4 9.012 Osmium Os 76 190.2Bismuth Bi 83 209.0 Oxygen O 8 16.00Bohrium Bh 107 264 Palladium Pd 46 106.4Boron B 5 10.81 Phosphorus P 15 30.97Bromine Br 35 79.90 Platinum Pt 78 195.1Cadmium Cd 48 112.4 Plutonium Pu 94 242Calcium Ca 20 40.08 Polonium Po 84 210Californium Cf 98 251 Potassium K 19 39.10Carbon C 6 12.01 Praseodymium Pr 59 140.9Cerium Ce 58 140.0 Promethium Pm 61 147Cesium Cs 55 132.9 Protactinium Pa 91 231Chlorine Cl 17 35.45 Radium Ra 88 226Chromium Cr 24 52.00 Radon Rn 86 222Cobalt Co 27 58.93 Rhenium Re 75 186.2Copper Cu 29 63.55 Rhodium Rh 45 102.9Curium Cm 96 247 Rubidium Rb 37 85.47Dubnium Db 105 262 Ruthenium Ru 44 101.1Dysprosium Dy 66 162.5 Rutherfordium Rf 104 261

362 Table of the Elements

Element SymbolAtomicNumber

AtomicMass Element Symbol

AtomicNumber

AtomicMass

Einsteinium Es 99 254 Samarium Sm 62 150.4Erbium Er 68 167.3 Scandium Sc 21 44.96Europium Eu 63 152.0 Seaborgium Sg 106 263Fermium Fm 100 253 Selenium Se 34 78.96Flourine F 9 19.00 Silicon Si 14 28.09Francium Fr 87 223 Silver Ag 47 107.9Gadolinium Gd 64 157.3 Sodium Na 11 22.99Gallium Ga 31 69.72 Strontium Sr 38 87.62Germanium Ge 32 72.59 Sulfur S 16 32.07Gold Au 79 197.0 Tantalum Ta 73 180.9Hafnium Hf 72 178.5 Technetium Tc 43 99Hassium Hs 108 265 Tellurium Te 52 127.6Helium He 2 4.003 Terbium Tb 65 158.9Holmium Ho 67 164.9 Thallium Tl 81 204.4Hydrogen H 1 1.008 Thorium Th 90 232.0Indium In 49 114.8 Thulium Tm 69 168.9Iodine I 53 126.9 Tin Sn 50 118.7Iridium Ir 77 192.2 Titanium Ti 22 47.88Iron Fe 26 55.85 Tungsten W 74 183.9Krypton Kr 36 83.80 Ununnilium Uun 110 269Lanthanum La 57 138.9 Unununium Uun 111 269Lawrencium Lr 103 257 Unumbium Uub 112 277Lead Pb 82 207.2 Uranium U 92 238.0Lithium Li 3 6.941 Vanadium V 23 50.94Lutetium Lu 71 175.0 Xenon Xe 54 131.3Magnesium Mg 12 24.31 Ytterbium Yb 70 1773.0Manganese Mn 25 54.94 Yttrium Y 39 88.91Meitnerium Mt 109 268 Zinc Zn 30 65.39Mendelevium Md 101 256 Zirconium Zr 40 91.22

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1

Author Index

Abegg, Richard, 74Agricola, Georgius, 16Al-Razi, 13Anaximander of Miletus, 9Andrews, Samuel, 305Aquinas, Thomas, 14Aristotle, 10, 13, 14, 18Arrhenius, Svante August, 141, 156, 157, 158, 159Avicenna, 13Avogadro, Amedeo, 35, 106

Bacon, Roger, 14Baekeland, Leo Hendrik, 298Bayen, Pierre, 26Bayer, Friedrich, 168Bayer, Karl Josef, 192Becher, John, 19, 20Becquerel, Antoine Henri, 38Bergman, Tobern, 27, 72Bernigaut, Louis Marie Hilaire, 297, 298Berthelot, Marcellin, 196Berthollet, Claude Louis, 27, 32, 33, 72Berzelius, Jöns Jacob, 34, 51, 73, 196, 197, 198Biringuccio, Vannoccio, 16Black, Joseph, 20, 21, 22, 25, 26Bohr, Niels, 40Boltzmann, Ludwig, 103Bosch, Karl, 153Bose, Satyendra Nath, 85Boullay, Pierre Francois Guillaume, 196Boyle, Robert, 18, 19, 21, 25, 102Bragg, William Henry, 91Bragg, William Lawrence, 91

Brønsted, Johannes, 158Bunsen, Robert, 296

Cannizzaro, Stanislao, 62, 63Carlisle, Anthony, 176Carothers, Wallace Hume, 298, 299, 300Carson, Rachael, 283Cavendish, Henry, 20, 25, 26, 27, 36Celsius, Anders, 104Chadwick, James, 37, 38Chapman, Sydney, 262Charles, Jacques Alexandre, 105, 107Chesebrough, Robert, 218Colgate, William, 306Couper, Archibald Scott, 198Crick, Francis, 234, 236Crookes, William, 35, 36Crutzen, Paul, 266Curie, Marie, 38Curie, Pierre, 38

Dalton, John, 10, 33, 34, 35, 38, 45, 46, 51, 56, 107,131, 176

Daniell, John Fredrick, 180Davy, Humphrey, 72, 73, 156, 176de Broglie, Louis, 40de Buffon, Comte, 72de Fourcroy, Antoine Francois, 27de Mestral, George, 299de Morveau, Guyton, 27Democritus, 9, 10, 47, 71Diocletian, 12Döbereiner, Johann, 61, 62

366 Author Index

Dow, Herbert Henry, 302, 303Drake, Edwin L., 216Du Fay, Charles, 172Dumas, Jean-Baptiste André, 196, 197, 198DuPont, Henry, 292, 293DuPont, LaMont, 293DuPont, Pierre Samuel, 292DuPont de Nemours, Eleuthère Irènèe, 292

Eastman, George, 298Ehrlich, Paul, 296Einstein, Albert, 85, 246Empedocles of Agrigentum, 9Epicurus, 10

Fahrenheit, Daniel, 104Faraday, Michael, 73, 176, 205, 294Fermi, Enrico, 46Firestone, Harvey Samuel, 305Fischer, Ernest Gottfried, 32Ford, Henry, 305Frankland, Edward, 73, 198Franklin, Benjamin, 23, 172Fresenius, Karl, 296Fuller, F. Buckminster, 96

Galen, 14Galilei, Galileo, 99Galvani, Luigi, 172, 173, 174, 180, 193Gamble, James, 305, 306Gamble, William, 306Gay-Lussac, Joseph Louis, 34, 35, 105, 107, 196,

290Geiger, Johannes Wilhelm, 39Gell-Mann, Murray, 44Geoffroy, Etienne Francois, 72Gerhardt, Charles Fredric, 168, 197, 198Gilbert, William, 171Glover, John, 290Goodrich, Benjamin Franklin, 304, 305Goodyear, Charles, 300, 304Gore, Wilbert, 299Graham, Thomas, 136Grove, Sir William, 188Guldberg, Cato Maximilian, 151

Haber, Fritz, 153, 154Hales, Stephen, 21, 22Hall, Charles Martin, 191, 192, 193Heisenberg, Werner, 40Henry, William, 33, 131Heraclitus of Ephesus, 9Héroult, Paul L. T., 191, 192, 193

Hippocrates, 167Hoffman, Felix, 168Hooke, Robert, 6, 18, 19, 297Hyatt, John Wesley, 298

Jabir ibn Hayyan (Geber), 13, 14Joule, James Prescott, 115

Kekulé, Friedrich August, 73, 198, 206, 296Kirwan, Richard, 27Koch, Robert, 296Kolbe, Adolph Wilhelm Hermann, 168, 196Kossel, Walther, 74Kroto, Harold, 96

Langmuir, Irving, 74Laurent, Auguste, 73, 197, 198Lavoisier, Antoine, 18, 20, 21, 23, 25, 26, 27, 28, 29,

31, 32, 35, 70, 156, 193, 196, 292Le Châtelier, Henri Louis, 149, 153LeBlanc, Nicholas, 291Leclanché, George, 186Leucippus, 9, 71Lewis, Gilbert Newton, 74, 75, 76, 159Libavius, Andreas, 16Libby, Williard, 246Liebig, Justus, 156, 197, 198, 295, 296Lister, Joseph, 296London, Fritz, 89Lowry, Thomas, 158, 159Lucretius, 10

Magnus, Albertus, 14Marsden, Ernest, 39Maxwell, James Clerk, 103Mayow, John, 19Mendeleev, Dmitri, 63, 64, 65Meyer, Julius Lothar, 63Millikan, Robert Andrew, 37Mitscherlich, Eilhardt, 205Molina, Mario, 265, 266Morehead, James T., 304Morton, William T. J., 209Moseley, Henry, 64Müller, Paul, 282, 283Mullis, Kary B., 236

Newlands, John, 62Newton, Isaac, 6, 25, 71, 72Nicholson, William, 176Nieuwland, Julius, 300Nobel, Alfred, 293Nobel, Immanuel, 293

Author Index 367

Paracelsus, 14, 15, 18Parmenides, 9Pasteur, Louis, 296Pauli, Wolfgang, 42Pauling, Linus, 77Perkins, William Henry, 295Perrin, Jean Baptiste, 56Planté, Gaston, 186Priestley, Joseph, 20, 23, 24, 25, 26, 27, 31, 172Procter, Harley, 306Procter, William, 305, 306Proust, Joseph Louis, 32, 33

Queeny, John Francis, 303

Richter, Jermias Benjamin, 31, 32Rockefeller, John D., 305Roebuck, John, 290Roentgen, Wilhelm Conrad, 38, 46Rouelle, Giullame Francois, 25Rowland, F. Sherwood, 265, 266Rutherford, Daniel, 22Rutherford, Ernest, 37, 39

Scheele, Carl, 24, 25, 26Schrödinger, Edwin, 40Schweppe, Jacob, 23Seaborg, Glenn, 69Seiberling, Charles, 304Seiberling, Frank A., 304Sennert, Daniel, 18

Smalley, Richard, 96Smith, Robert Angus, 267Sobrero, Ascanio, 293Solvay, Ernest, 292Sorenson, Soren Peer Lauritz, 162Stahl, George, 20, 21Stoney, George, 37

Thompson, William (Lord Kelvin), 104, 105Thomson, Joseph John, 36, 37, 38, 73, 176Topham, Charles, 297Torricelli, Evangelista, 99, 100Tyndall, John, 137

van der Waals, Johannes Diederick, 90, 91van Helmont, Jan Baptista, 17, 18, 21van Musschenbroek, Pieter, 172van’t Hoff, Jacobus Hendricus, 142Volta, Alessandro, 72, 172, 174, 175, 176, 180, 193von Guericke, Otto, 171von Hofmann, Wilhelm, 295, 296von Laue, Max Theodor Felix, 91

Waage, Peter, 151Watson, James, 234, 236Watt, James, 22, 26Williams, Charles Greville, 300Willson, Thomas L., 304Wöhler, Friedrich, 195, 196, 197, 294

Zweig, George, 44

1

Subject Index

Absolute zero, 105Acetone, 210Acid ionization constant (Ka), 160Acid precipitation, 164, 266Acid rain, 266–70, 314; effect on aquatic

organisms, 269Acids: common, 157; definitions of, 156–59; gen-

eral properties of, 156; history of 155–56,290–91; stock solutions, 299

Actinides, 67Activated complex, 141Activation energy, 141Activity series, 181Addition polymerization, 203Adhesion, 95Alchemy, 11Alcohols, 208, 217; common, 386; denatured, 386Aldehyde, 209–10Alizarin, 294Alkali metal, 64Alkaline earth metal, 64Alkaloid, 215Alkanes, 199–201, 205Alkenes, 202–3, 205Alkyl group, 207Alkynes, 202–3, 205Allotrope, 495Alloy, 8, 317Alpha: decay, 39, 243–44; particle, 39, 243–44Alpha helix, 231–32Aluminum, 191–93American Chemical Society, 332–33Amine, 215–17

Amino acid, 229; essential, 230Ammonia, 152–54Ammonium ion, 52Amorphous solids, 92Amylopectin, 224Amylose, 224Anabolism, 221Androgen, 228Anion, 46Annealing, 8Antacids, 167Antimatter, 254Aqua regia, 155Aromatic hydrocarbons, 205–7, 217Arsenic, 276; use as pesticide, 282Arteriosclerosis, 228Aryl group, 207Aspirin, 168–69Atmospheric pressure, 99–100, 318–20; composi-

tion of, 100Atom: Bohr model, 40; history of, 9, 33–34; plane-

tary model, 39Atomic mass, 45Atomic number, 42, 45Atomic orbitals, 40–43Atomic theory, 31–34Avogadro’s law, 106Avogadro’s number, 56

Bakelite, 298Base ionization constant (Kb), 160BASF, 306–7Bayer Company, 168, 306–7

370 Subject Index

Becquerel: radiation unit, 244Benzene, 205–7Beta: decay, 39, 243–44; particle, 39, 243–44B. F. Goodrich Company, 301–4Binding energy, 246–47Biological magnification, 283Bipyridyliums, 286Boiling point elevation, 132–33Bond energies, 79Bonds: covalent, 76–78; early ides, 71–74; ionic,

74–76; metallic, 80; polar covalent, 76–77Bose-Einstein condensate, 85Boyle’s Law, 102Breeder reactor, 249Buckminsterfullerenes, 96–97Buffer, 166–67

Caffeine: structure, 215Calcination, 19–20Calorimetry, 118–20Calx, 20Carbamates, 284Carbohydrates, 221–26Carbon dioxide: emissions by country, 273Carbon monoxide, 279–80Carbon-14 dating, 246Carbonyl, 209Carboxyl, 210Carboxylic acids, 210–12, 217Catabolism, 221Catalysts, 144–47Catalytic converter, 146Cathode-ray tube, 35–36Cathode rays, 35–36Cathodic protection, 190Cation, 46Cellulase, 223–24Cellulose, 223–24Celsius temperature scale, 104Chain reaction, 248Chapman reactions, 262Charles’ law, 105Chemical change, 4Chemical composition: of Earth, 70; of human body,

70Chemical compound: naming, 51–53Chemical equilibrium, 147–49Chemical formula, 54Chemist’s code of conduct, 333Chirality, 222–23Chlorine: as water disinfectant, 275Cohesion, 95Colligative properties, 131–34Collision theory, 140–42

Colloids, 136–37Compounds, 49; common names, 53; cyclic, 205;

naming, 51–53Concentration units: of solutions, 129Conjugate acid-base pairs, 158Corrosion, 189–90Cracking process, 218Crenation, 133Crookes tube, 35Crystalline solid, 90–92Curie: radiation unit, 255Cyclic compounds, 205

Dalton’s Law of Partial Pressure, 107d-Block elements, 67DDT, 282–84Dephlogisticated air, 23Diabetes, 224–25Dialysis, 133Diamond: structure, 4; synthetic, 97Dioxins, 285Dipole-dipole force, 89Dipole moment, 271Dispersion force, 89DNA, 2; replication, 235; structure, 233–35Dry cell battery, 185–86Dry ice, 324Dupont Chemical Company, 292–93, 302Dyes, 293–97Dynamite, 293

Electrochemical cell, 179–84Electrolysis: of water, 185, 315Electrolyte, 128Electromagnetic radiation, 262; spectrum of, 263Electromotive force, 182Electron: charge to mass ratio, 36–37; discovery of,

36–37; size, 45Electron configuration, 42–45Electronegativity, 76–78Electroplating, 190Element, 3, 28Elementary steps, 145Elements: discovery of, 61–62; periodic trends, 63;

table of, 361–62Emulsifier, 227Endothermic reaction, 121Energy content: of selected fuels, 122Enzymes, 146–47Error: random, 312; systematic, 312Esters, 211–12, 217Estrogens, 228Ether, 208–9, 217; as anesthetic, 209Eutrophication, 274

Subject Index 371

Exothermic reaction, 120Exxon, 301, 305

Fahrenheit temperature scale, 104Fatty acids, 211, 226–27; common, 212; in common

fats and oils, 227; in soap, 213Fermentation, 8Firestone Tire and Rubber Company, 305First law of thermodynamics, 121Fission, 246–47Fixed air, 22Formaldehyde, 209Fossil fuels, 216Fractional distillation, 217Free radical, 265Freezing point depression, 132Fructose, 222Fuel cells, 188Fullerene, 96–97Functional group, 198, 207, 217Fundamental particle, 44–45Fusion, 247

Gamma: decay, 39, 243; energy, 39, 243Gasoline, 218; combustion of, 279Gay Lussac’s law, 105–6Geiger counter, 256Genetic engineering, 238Glass, 92–93Glucose, 222–23Glycogen, 224Goodyear Tire and Rubber Company, 300, 304Gore-Tex, 299Gunpowder, 292–93

Haber process, 152–53Half-Life: chemical, 144; of common isotopes, 245;

radioactive, 244Half reactions, 177Halogen, 64Halogenation, 202Hard water, 274Heat capacity, 116Heat of: fusion, 117, 119; reaction, 120;

vaporization, 117, 119Henry’s law, 110, 131Histamine, 215Human genome project, 236Hydration, 127Hydrocarbons, 199; saturated, 199; unsaturated, 199Hydrogen, 25Hydrogen bond, 87–88Hydrogenation, 202Hydrolysis, 213

Ideal gas, 102Ideal gas law, 106–7Indigo, 294Internal energy, 114Invert sugar, 223Iodine number, 226; of common fats and oils, 227Ion product constant (Kw), 162Ion-dipole force, 89Ionization, 46Ionizing radiation, 256Ions, 45–46; present in seawater, 125Iron: corrosion of, 189Isotope, 45; half-lives of, 245IUPAC, 69Ivory soap, 305–6

Joule: unit of energy, 115

Ka, 160Kb, 160Kelvin temperature scale, 104–6Ketone, 209–10, 217

Latex, 299Lathanides, 67Law of Definite Proportions, 33–34Law of mass action, 151Law of Multiple Proportions, 33–34Lead: toxic effects, 266–67Lead-chamber process, 290Lead storage battery, 186–87LeBlanc method, 291LeChâtelier’s Principle, 149–50Lewis acid, 159Lewis base, 159Lewis electron dot formula, 74–76Limewater, 321Lipids, 226Liquid crystals, 95–96Lock-and-key model, 147London force, 89Long-range order, 93

Main group elements, 64–67Mass defect, 246Mass number, 45Matter, 3, 10Mauve, 295Mercury: toxic effects, 276–77Metallic bond, 80Metalloids, 67Metals, 67–69Mixture, 49Mole, 56

372 Subject Index

Molecular orbital theory, 83Molecule, 35, 49–50Monsanto, 303–4MTBE, 209

Neutron, 37, 45Nicotine: as pesticide, 282; structure, 215Nitroglycerin, 293Nitrous gas, 34Nobel prize, 293Nobel prizes: in chemistry, 355–59Noble gas, 64, 68Nomenclature: ancient symbols, 50; common

names, 53; history of, 49–52; for naming com-pounds, 51–54

Non-ionizing radiation, 256Nonmetals, 67Nuclear fission, 246–47Nuclear fusion, 247; and stellar evolution, 251–52Nuclear power, 248–51Nuclear stability, 241–42Nucleon, 241Nylon, 298–99

Octane number, 201–2Octet rule, 74Oil, 94Optical isomer, 296Organic Foods Production Act, 219Organic substitution, 197Organochloride, 284Organophosphates, 284Osmosis, 133Osmotic pressure, 133Oxidation numbers, 177–79Oxidation reaction, 176–77Oxidizing agent, 177Oxygen: discovery of, 23–25Ozone, 262; depletion, 262–66

Partial pressure, 107; and diving, 109–10Pauli Exclusion Principle, 42PCBs, 277–78Peptide, 229Perfect gas, 102Periodic table, inside cover; history of, 63–64; mod-

ern, 64; recently discovered elements and, 69Pesticides, 281–87; history of, 281–82Petrochemicals, 218pH, 161–63; of common substances, 163; indicator,

163, 310Phase changes, 113–16Phenol, 208Phenyl, 207

Pheromones, 212Philosopher’s stone, 11Phlogiston theory, 19–20Physical change, 4Plasmolysis, 133Plastics, 297–99; recycling symbols, 204Pneumatic trough, 21pOH, 161–63Polarizability: and intermolecular forces, 90Polarized light, 96Polyatomic ions, 52Polyethylene terephthalate (PET), 214Polymerase chain reaction, 236–37Polymerization, 203Polysaccharide, 223Potash, 290–91Potential energy, 114; chemical, 120Pressure, 99; units, 101Procter and Gamble, 305–6; and Ivory soap, 306Progestin, 228Protein, 229; structure, 229–32; synthesis, 235Proton, 44–45; discovery of, 37Pyrethroids, 284–85

Quantum numbers, 40–42Qualitative analysis, 136Quarks, 44Quaternary structure, 231–32

Radiation: background, 259; biological effects,257–59; detection, 256–57; units, 255–56

Radical theory, 196Radiometric dating, 244–46Random error, 312Rate law, 143Reaction order, 143Reaction rates: factors affecting, 142Reactions, 54; balancing, 55; complete, 147; oxida-

tion/reduction, 176–77; reversible; 147; typesof, 55

Recombinant DNA, 238Recycling symbols, 204Reducing agent, 177Relative humidity, 108Relative mass, 45; Dalton’s, 33Resins, 209Resonance structures, 206Reverse osmosis, 134RNA, 235–36Rubber: synthetic, 299–301

Saccharides, 221, 223Saccharin, 2SAE number, 94

Subject Index 373

Saponification, 213–14Scientific method, 310–13Sickle-cell anemia, 231Smog, 280–81Soap, 213Soda ash, 291. See also LeBlanc method; Solvay

processSolids: amorphous, 92; crystalline, 90; ionic, 91;

metallic, 91; molecular, 91Solubility: of gases, 130–131; rules for ionic com-

pounds, 128; and temperature, 130Solutions: concentration units, 129; saturated, 129Solvation, 127Solvay process, 292Specific heat, 116; of common substances, 116Specific heat capacity, 116Standard hydrogen electrode, 182Standard reduction potential, 182; of common half-

reactions, 184Starch, 224States of matter, 85–87Stereoisomers, 223Steroids, 228Stoichiometry, 32Structural isomers, 201Sublimation, 113–14Sulfuric acid, 155, 157; lead chamber process, 290Surface tension, 94

Temperature scales, 104Tokamak, 251

Transmutation, 252Triazines, 285–86Trihalomethanes, 275Tyndall effect, 137

Ultraviolet radiation, 263; in ozone depletion,262–63

Union Carbide, 304Urea: synthesis of, 195

Valence, 73Valence bond theory, 83Valence electron configuration, 43–44Valence shell electron pair repulsion (VSEPR),

80–83Vapor pressure, 107Vaporization, 113–14; heat of, 117Vaseline, 218Velcro, 299Viscosity, 93Vitalism, 195Vulcanization, 300

Water: hydrogen bonding in, 87; quality, 273–78

Wien’s Law, 270Work: P-V, 115

X-Ray crystallography, 91–92

Zwitterion, 231

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The Basics of Chemistrys2.bitdl.ir/Ebook/Chemistry/Myers - The Basics of Chemistry (Greenwood... · Colligative Properties 131 Reactions of Solutions 134 Colloids and Suspensions