Chapter One: CHEMICAL FOUNDATIONS 講義. Copyright © Houghton Mifflin Company. All rights reserved.Chapter 1 | Slide 2 What we’ll learn in Chapter 1 Chemistry

Chapter One:

CHEMICALFOUNDATIONS

講義

Copyright © Houghton Mifflin Company. All rights reserved.Chapter 1 | Slide 2

What we’ll learn in Chapter 1• Chemistry overview• Science, theory and law, experiment• Measurement, SI units and unit conversion,

prefixes of numbers• Significant figures, rounding rules, precision

and accuracy, errors (systematic or random)• Classification of matter, states of matter,

purification and separation• Physical change, chemical change


Assignment

• 11,18,27,32,59,70,71,85


Chemistry: An Overview

• A main challenge of chemistry is to understand the connection between the macroscopic world that we experience and the microscopic world of atoms and molecules.

• You must learn to think on the atomic level.

1.1


Atoms vs. Molecules

1.1

user

Physicists and chemists may have different concerns about atoms and molecules. Physicists habor dispise on chemists: cheimists are shallow, pigs. Chemists have disgust on physicists: physicists are sick, nitty-picking, rats.


Oxygen and Hydrogen Molecules

1.1


A Chemical Reaction

1.1

user

this reaction has not been fully understood yet, perhaps in contrast to your guess.


A Chemical Reaction

1.1

user

this simple reaction has not been fully understood yet either.


Science

• Science is a framework for gaining and organizing knowledge.

• Science is a plan of action—a procedure for processing and understanding certain types of information.

1.2


The Various Parts of the Scientific Method

(?)

(!)

user

Newton didn't even observe and made a correct and great hypothesis when he was in a dream drinking free beer ('cause he didn't have money to drink in the taverns on Cambridge campus with his rich schoolmates). Boltzmann made a correct prediction which actually cost his life (because he was not mentally strong enough compared wt his enemies).


Law vs. Theory

• A law summarizes what happens.

• A theory (model) is an attempt to explain why it happens.

1.2


Post-it note: a nice story of doing chemistry

Making observationsFormulating hypothesesPerforming experiments

How to make sticky-but-not-too-sticky adhesives?

Too sticky

Not sticky enough

Right!


Nature of Measurement

• Measurement – quantitative observation consisting of two parts:

NumberScale (unit)

• Examples:20 grams6.63 × 106.63 × 10-34-34 joule·seconds

1.3


The Fundamental SI Units

Physical Quantity Name of Unit Abbreviation

Mass kilogram kg

Length meter m

Time second s

Temperature kelvin K

Electric current ampere A

Amount of substance mole mol

Luminous intensity candela cd

1.3


Table 1.1 The Fundamental SI Units


Table 1.2 The Prefixes Used in the SI System


Table 1.3 Some Examples of Commonly Used Units


Table 1.4 English-Metric Equivalents


Artist's Conception of the Lost Mars Climate Orbiter


Uncertainty in Measurement

• A digit that must be estimated is called uncertain.

• A measurement always has some degree of uncertainty.

• Record the certain digits and the first uncertain digit (the estimated number).

1.4


Measurement of Volume Using a Buret

1.4


Precision and Accuracy

• Accuracy – agreement of a particular value with the true value

• Precision – degree of agreement among several measurements of the same quantity

1.4


Precision and Accuracy


Rules for Counting Significant Figures

• Nonzero integers always count as significant figures:

3456 has 4 sig figs

1.5


Rules for Counting Significant Figures (continued)

• Leading zeros do not count as significant figures:

0.048 has 2 sig figs

1.5



• Captive zeros always count as significant figures:


1.5



• Trailing zeros are significant only if the number contains a decimal point:


150 has 2 sig figs

1.5



• Exact numbers have an infinite number of significant figures: 1 inch = 2.54 cm, exactly 9 pencils (obtained by counting)

1.5


Sig Figs in Mathematical Operations

• For multiplication or division, the number of significant figures in the result is the same as the number in the calculation that has the fewest significant figures:

1.342 × 5.5 = 7.381 7.4

1.5


Sig Figs in Mathematical Operations (continued)

• For addition or subtraction, the result has the same number of decimal places as the measurement with the fewest number of decimal places:

23.445

+ 7.83

=31.275 31.28

1.5


Concept Check

You have water in each graduated cylinder shown. You then add both samples to a beaker.

How would you write the number describing the total volume?

What limits the precision of this number?


Dimensional Analysis

• Use when converting a given result from one system of units to another:– Use the equivalence statement that relates

the two units– Consider the direction of the required change

to select the correct unit factor (cancel unwanted units)

– Multiply the quantity to be converted by the unit factor

1.6


Concept Check

What data would you need to estimate the money you would spend on gasoline to drive your car from New York to Chicago? Provide estimates of values and a sample calculation.

1.6


Temperature

• Three systems for measuring temperature:– Fahrenheit– Celsius– Kelvin

1.7


The Three Major Temperature Scales

1.7


Converting Between Scales

K = °C + 273.15 °C = K – 273.15

°C = (°F – 32)(5/9) °F = °C(9/5) + 32

1.7


Exercise

At what temperature does C = F?


Density

• Mass of substance per unit volume of the substance:

Density = mass/volume

1.8


Table 1.5 Densities of Various Common Substances* at 20° C


Classification of Matter

• Matter – anything occupying space and having mass.

• Matter exists in three states:

– Solid

– Liquid

– Gas

1.9


Classification of Matter

• Solid – rigid; has fixed volume and shape

• Liquid – has definite volume but no specific shape; assumes shape of container

• Gas – has no fixed volume or shape; takes on the shape and volume of its container

1.9


The Three States of Water

1.9


Structure of a Solid

1.9


Structure of a Liquid

1.9


Structure of a Gas

1.9


Mixtures

• Mixtures have variable composition: Homogeneous – having visibly

indistinguishable parts; solution Heterogeneous – having visibly

distinguishable parts

1.9


Homogeneous Mixtures

1.9


Homogeneous vs. Heterogeneous

1.9


Compound vs. Mixture

1.9


Simple Laboratory Distillation Apparatus

1.9


The Organization of Matter

1.9





1.9



1.9


Concept Check

Sketch a magnified view (showing atoms/molecules) of each of the following:

– A heterogeneous mixture of two different compounds

– A homogeneous mixture of an element and a compound

Chapter One:

CHEMICALFOUNDATIONS

案例 /討論


Figure 1.4 The Fundamental Steps of the Scientific Method


Figure 1.5 The Various Parts of the Scientific Method


Figure 1.6 Measurement of Volume


Figure 1.7 Common Types of Laboratory Equipment Used to Measure Liquid Volume


Figure 1.9 Measurement of Volume Using a Buret


Figure 1.10 The Results of Several Dart Throws Show the Difference Between Precise

and Accurate


Figure 1.11 The Three Major Temperature Scales


Figure 1.12 Normal Body Temperature


Figure 1.13 The Three States of Water


Figure 1.16 The Organization of Matter


Figure 1.1a The Surface of a Single Grain of Table Salt


Figure 1.1b An Oxygen Atom on a Gallium Arsenide Surface


Figure 1.1c Scanning Tunneling Microscope Image


Figure 1.2 A Charged Mercury Atom


Figure 1.3a Each Grain of Sand is Composed of Tiny Atoms


Figure 1.3b Beach at Big Sur


Robert Boyle


Soda is Sold in 2-Liter Bottles- an Example of SI Units in Everyday Life


Figure 1.8 An Electronic Analytic Balance


Rounding Numbers


Liquid Nitrogen


Figure 1.14 Simple Laboratory Distillation Apparatus


Figure 1.15a A Line of the Mixture to be Separated is Placed at One End of a Sheet


Figure 1.15b The Paper Acts as a Wick to Draw up the Liquid


Figure 1.15c Component with the Weakest Attraction for the Paper Travels Faster


Mercury and Iodine Combine to Form Mercuric Iodide


Table 1.1 The Fundamental SI Units


Table 1.2 The Prefixes Used in the SI System


Table 1.3 Some Examples of Commonly Used Units


Table 1.4 English-Metric Equivalents


Table 1.5 Densities of Various Common Substances* at 20° C

Chapter One

Chemical Foundations

問答


Question

• Which of the following is an example of a quantitative observation?– Solution A is a darker red color than solution B. – The grass is green.– Substance A has a greater mass than

substance B.– The temperature of the water is 45°C.


Answer

• d) The temperature of the water is 45°C.

• A quantitative observation includes a measurement (numerical) and a unit.


Question

• The glassware shown below is called a buret. The buret is filled to the zero mark (at the top) with a solution and the solution is transferred to a beaker. What volume of transferred solution should be reported? – 20 mL – 22 mL– 22.0 mL– 22.00 mL– 25 mL


Answer

• c) 22.0 mL

• In a measurement, we always include one uncertain digit. The graduations on this buret are in 1-mL units, so we can estimate the volume to the tenths place.


Question

• The boiling point of a liquid was measured in the lab, with the following results:

• Trial Boiling Point• 1 22.0°C ± 0.1• 2 22.1°C ± 0.1• 3 21.9°C ± 0.1

• The actual boiling point of the liquid is 28.7°C. The results of the determination of the boiling point are

– accurate and precise. – precise but inaccurate.– accurate but imprecise.– inaccurate and imprecise.


Answer

• b) precise but inaccurate.

• The measurements are precise because they are all in close agreement with one another. However, they are relatively far from the true value, so they are inaccurate.


Question

• _______ reflects the reproducibility of a given type of measurement.– Accuracy – Precision– Certainty– Systematic error– Random error


Answer

• b) Precision

• Measurements are precise if they are relatively close to one another, regardless of how close they are to the true answer.


Question

• _______ is the agreement of a particular value with the true value.– Accuracy – Precision– Certainty– Systematic error– Random error


Answer

• a) Accuracy

• If a measurement is in close agreement with the true value, it is an accurate measurement.


Question

• After performing a calculation in the lab, the display on your calculator reads “0.023060070”. If the number in the answer is to have five significant figures, what result should you report?– 0.0230 – 0.00231– 0.023060– 0.2367– 0.02306


Answer

• c) 0.023060

• The leading zeros are not significant, but the captive zero and the trailing zero are significant.


Question

• How many significant figures are in the number 0.03040?– 1 – 2– 3– 4– 5


Answer

• c) 4

• The leading zeros are not significant, but the captive zero and the trailing zero are significant.


Question

•The beakers below have different precisions.

•You pour the water from these three beakers into one container. What is the volume in this container reported to the correct number of significant figures?•

– 78.817 mL – 78.82 mL – 78.8 mL – 79 mL


Answer

• d) 79 mL

• In a measurement, we always include one uncertain digit. In this case, the first measurement could be 26.4 mL ± 0.1 mL, the second could be 26 mL ± 1 mL, and the third could be 26.42 mL ± 0.01 mL. When adding, the result has the same number of decimal places as the least precise measurement—in this case, to the ones place. So the answer is 26.4 + 26 + 26.42 = 78.82 mL, which must be rounded to 79 mL.


Question

• Express 3140 in scientific notation.– 3.14 × 103 – 3.14 × 10-3

– 3.140 × 103

– 3.140 × 10-3


Answer

• a) 3.14 × 103

• 103 = 1000, and 3.14 × 1000 = 3140. We lose the zero because it is not significant (it is a placeholder). If the zero was significant, we should write the number as “3140.”.


Question

• A solution is also a – heterogeneous mixture. – homogeneous mixture.– compound.– distilled mixture.– pure mixture.


Answer

• b) homogeneous mixture.

• Solutions can be liquids in liquids (such as rubbing alcohol, which consists of isopropanol in water), solids in liquids (such as sugar water), a mixture of gases (such as air), or even a mixture of solids (such as brass, which consists of a mixture of copper and zinc).


Question

•Which of the following statements is false? – Solutions are always homogeneous

mixtures. – Atoms that make up a solid are mostly

open space. – Elements can exist as atoms or

molecules. – Compounds can exist as elements or

molecules.


Answer

•d) Compounds can exist as atoms or molecules.

•Elements can be atoms (such as He) or molecules (such as O2), but compounds must exist as molecules.


Richard III of England • Richard III (2 October 1452 – 22 August 1485) was King of England from 1483 until his death. He was the last king from the House of York, and his defeat at the Battle of Bosworth marked the culmination of the Wars of the Roses and the end of the Plantagenet dynasty. After the death of his brother King Edward IV, Richard briefly governed as regent for Edward's son King Edward V with the title of Lord Protector, but he placed Edward and his brother Richard in the Tower (see Princes in the Tower) and seized the throne for himself, being crowned on 6 July 1483.

• Two large-scale rebellions rose against Richard. The first, in 1483, was led by staunch opponents of Edward IV and, most notably, Richard's own 'kingmaker', Henry Stafford, 2nd Duke of Buckingham. The revolt collapsed and Buckingham was executed at Salisbury, near the Bull's Head Inn. However, in 1485, another rebellion arose against Richard, headed by Henry Tudor, 2nd Earl of Richmond (later King Henry VII) and his uncle Jasper. The rebels landed troops and Richard fell in the Battle of Bosworth Field, then known as Redemore or Dadlington Field, as the last Plantagenet king and the last English king to die in battle.


Raider’s lost ark

Gold idol

Booby-trapped pedestal

A bag of sand


Documents

Chapter One: CHEMICAL FOUNDATIONS 講義. Copyright © Houghton Mifflin Company. All rights reserved.Chapter 1 | Slide 2 What we’ll learn in Chapter 1 Chemistry