Chapter 1

Chemistry: Methods and Measurement

Deniston Topping Caret

7th Edition

Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

1.1 The Discovery Process

• Chemistry - The study of matter…

– Matter - Anything that has mass and occupies space• A table• A piece paper

– What about air? • Yes, it is matter

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Chemistry:• the study of matter• its chemical and physical properties• the chemical and physical changes it

undergoes• the energy changes that accompany

those processes

• Energy - the ability to do work to accomplish some change

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cess THE SCIENTIFIC METHOD

• The scientific method - a systematic approach to the discovery of new information

Characteristics of the scientific process

1. Observation

2. Formulation of a question

3. Pattern recognition

4. Developing theories

5. Experimentation

6. Summarizing information

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Models in Chemistry

• To aid in understanding of a chemical unit or system– a model is often used– good models are based on

everyday experience

• Ball and stick methane model– color code balls– sticks show attractive forces

holding atoms together

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1.2 Matter and Properties

• Properties - characteristics of matter– chemical vs. physical

• Three states of matter1. gas - particles widely separated, no definite

shape or volume solid

2. liquid - particles closer together, definite volume but no definite shape

3. solid - particles are very close together, define shape and definite volume

Three States of Water

(a) Solid (b) Liquid (c) Gas

Comparison of the Three Physical States

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ies • Physical property - is observed

without changing the composition or identity of a substance

• Physical change - produces a recognizable difference in the appearance of a substance without causing any change in its composition or identity- conversion from one physical state to

another

- melting an ice cube

Separation by Physical Properties

Magnetic iron is separated from other nonmagnetic substances, such as sand. This property is used as a large-scale process in the recycling industry.

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• Chemical property - result in a change in composition and can be observed only through a chemical reaction

• Chemical reaction (chemical change) - a process of rearranging, removing, replacing, or adding atoms to produce new substances

hydrogen + oxygen water

reactants products

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ies Classify the following as either a

chemical or physical property:

a. Color

b. Flammability

c. Hardness

d. Odor

e. Taste

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ies Classify the following as either a

chemical or physical change:

a. Boiling water becomes steam

b. Butter turns rancid

c. Burning of wood

d. Mountain snow pack melting in spring

e. Decay of leaves in winter

• Intensive properties - a property of matter that is independent of the quantity of the substance- Density

- Specific gravity

• Extensive properties - a property of matter that depends on the quantity of the substance- Mass

- Volume1.2

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• Pure substance - a substance that has only one component

• Mixture - a combination of two or more pure substances in which each substance retains its own identity, not undergoing a chemical reaction

Classification of Matter

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• Element - a pure substance that cannot be changed into a simpler form of matter by any chemical reaction

• Compound - a substance resulting from the combination of two or more elements in a definite, reproducible way, in a fixed ratio

Classification of Matter

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• Mixture - a combination of two or more pure substances in which each substance retains its own identity

• Homogeneous - uniform composition, particles well mixed, thoroughly intermingled

• Heterogeneous – nonuniform composition, random placement

Classification of Matter

Classes of Matter1.

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1.3 Measurement in ChemistryData, Results, and Units• Data - each piece is an individual result of a single

measurement or observation– mass of a sample

– temperature of a solution

• Results - the outcome of the experiment• Data and results may be identical, however usually

related data are combined to generate a result• Units - the basic quantity of mass, volume or

whatever quantity is being measured– A measurement is useless without its units

English and Metric Units

• English system - a collection of functionally unrelated units – Difficult to convert from one unit to

another – 1 foot = 12 inches = 0.33 yard = 1/5280 miles

• Metric System - composed of a set of units that are related to each other decimally, systematic – Units relate by powers of tens– 1 meter = 10 decimeters = 100 centimeters = 1000

millimeters

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Mass gram gLength meter mVolume liter L

• Basic units are the units of a quantity without any metric prefix

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• You must be able to convert between units

- within the metric system

- between the English system and metric system

• The method used for conversion is called the Factor-Label Method or Dimensional Analysis

!!!!!!!!!!! VERY IMPORTANT !!!!!!!!!!!

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simply memorizing connections between units.– For example: How many donuts are in

one dozen?

– We say: “Twelve donuts are in a dozen.”– Or: 12 donuts = 1 dozen donuts

• What does any number divided by itself equal?

• ONE! 1 dozen 1

donuts 12

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dozen 1

donuts 12

• This fraction is called a unit factor

• What does any number times one equal?

• That number

• Multiplication by a unit factor does not change the amount – only the unit

• We use these two mathematical facts to use the factor label method– a number divided by itself = 1– any number times one gives that number

back

• Example: How many donuts are in 3.5 dozen?

• You can probably do this in your head but try it using the Factor-Label Method.1.

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3.5 dozen

Then set up your unit factor...

dozen 1

donuts 12

See that the units cancel...

Then multiply and divide all numbers...

= 42 donuts

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yCommon English System Units

• Convert 12 gallons to units of quarts

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yIntersystem Conversion Units

• Convert 4.00 ounces to kilograms

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2. Convert 50.0 milliliters to pints

3. Convert 1.8 in2 to cm2

1.4 Significant Figures and Scientific Notation

• Information-bearing digits or figures in a number are significant figures

• The measuring devise used determines the number of significant figures a measurement has

• The amount of uncertainty associated with a measurement is indicated by the number of digits or figures used to represent the information

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Significant figures - all digits in a number representing data or results that are known with certainty plus one uncertain digit

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n Recognition of Significant Figures

• All nonzero digits are significant

• 7.314 has four significant digits

• The number of significant digits is independent of the position of the decimal point

• 73.14 also has four significant digits

• Zeros located between nonzero digits are significant

• 60.052 has five significant digits

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nUse of Zeros in Significant

Figures• Zeros at the end of a number (trailing zeros) are

significant if the number contains a decimal point.

• 4.70 has three significant digits

• Trailing zeros are insignificant if the number does not contain a decimal point.

• 100 has one significant digit; 100. has three

• Zeros to the left of the first nonzero integer are not significant.

• 0.0032 has two significant digits

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How many significant figures are in the following?

1. 3.400

2. 3004

3. 300.

4. 0.003040

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n Scientific Notation

• Used to express very large or very small numbers easily and with the correct number of significant figures

• Represents a number as a power of ten

• Example:

4,300 = 4.3 x 1,000 = 4.3 x 103

• To convert a number greater than 1 to scientific notation, the original decimal point is moved x places to the left, and the resulting number is multiplied by 10x

• The exponent x is a positive number equal to the number of places the decimal point moved

5340 = 5.34 x 104

• What if you want to show the above number has four significant figures?

= 5.340 x 104

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• To convert a number less than 1 to scientific notation, the original decimal point is moved x places to the right, and the resulting number is multiplied by 10-x

• The exponent x is a negative number equal to the number of places the decimal point moved

0.0534 = 5.34 x 10-2

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Types of Uncertainty

• Error - the difference between the true value and our estimation

– Random

– Systematic

• Accuracy - the degree of agreement between the true value and the measured value

• Precision - a measure of the agreement of replicate measurements1.

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correct answer 152.83 liters

Significant Figures in Calculation of Results

Rules for Addition and Subtraction

• The result in a calculation cannot have greater significance than any of the quantities that produced the result

• Consider:

37.68 liters 6.71862 liters

108.428 liters 152.82662 liters

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)calculator(on 109688692.210255.2

)94.15(102.4 84

3

Which number has the fewest significant figures? 4.2 x 103 has only 2

The answer is therefore, 3.0 x 10-8

Rules for Multiplication and Division

• The answer can be no more precise than the least precise number from which the answer is derived

• The least precise number is the one with the fewest significant figures

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n Exact and Inexact Numbers

• Inexact numbers have uncertainty by definition

• Exact numbers are a consequence of counting

• A set of counted items (beakers on a shelf) has no uncertainty

• Exact numbers by definition have an infinite number of significant figures

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=3.35 x 104

Rules for Rounding Off Numbers

• When the number to be dropped is less than 5 the preceding number is not changed

• When the number to be dropped is 5 or larger, the preceding number is increased by one unit

• Round the following number to 3 significant figures: 3.34966 x 104

How Many Significant Figures?

Round off each number to 3 significant figures:

1. 61.40

2. 6.171

3. 0.066494

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1.5 Experimental Quantities

• Mass - the quantity of matter in an object– not synonymous with weight– standard unit is the gram

• Weight = mass x acceleration due to gravity

• Mass must be measured on a balance (not a scale)

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• Units should be chosen to suit the quantity described – A dump truck is measured in tons– A person is measured in kg or pounds– A paperclip is measured in g or ounces– An atom?

• For atoms, we use the atomic mass unit (amu)– 1 amu = 1.661 x 10-24 g

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• Length - the distance between two points– standard unit is the meter– long distances are measured in km– distances between atoms are measured in nm,

1 nm = 10-9 m

• Volume - the space occupied by an object– standard unit is the liter– the liter is the volume occupied by 1000

grams of water at 4 oC– 1 mL = 1/1000 L = 1 cm3

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The milliliter (mL) and the cubic centimeter (cm3) are equivalent

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- metric unit is the second

• Temperature - the degree of “hotness” of an object

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1.8

32-FC

oo

32C)(1.8F oo

1. Convert 75oC to oF

2. Convert -10oF to oC

1. Ans. 167 oF 2. Ans. -23oC

Conversions Between Fahrenheit and Celsius

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K = oC + 273

Kelvin Temperature Scale

• The Kelvin scale is another temperature scale.

• It is of particular importance because it is directly related to molecular motion.

• As molecular speed increases, the Kelvin temperature proportionately increases.

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• Energy - the ability to do work

• kinetic energy - the energy of motion

• potential energy - the energy of position (stored energy)

• Energy is also categorized by form:• light• heat• electrical• mechanical• chemical

Energy

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Characteristics of Energy

• Energy cannot be created or destroyed

• Energy may be converted from one form to another

• Energy conversion always occurs with less than 100% efficiency

• All chemical reactions involve either a “gain” or “loss” of energy

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• Basic Units:• calorie or joule• 1 calorie (cal) = 4.184 joules (J)

• A kilocalorie (kcal) also known as the large Calorie. This is the same Calorie as food Calories.

• 1 kcal = 1 Calorie = 1000 calories

• 1 calorie = the amount of heat energy required to increase the temperature of 1 gram of water 1oC.

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Density and Specific Gravity

• Density – the ratio of mass to volume– an extensive property– use to characterize a substance as

each substance has a unique density

– Units for density include:• g/mL• g/cm3

• g/cc

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liquid mercury

brass nut

water

cork

Calculating the Density of a Solid

• 2.00 cm3 of aluminum are found to weigh 5.40g. Calculate the density of aluminum in units of g/cm3.– Use the formula– Substitute our values

5.40 g

2.00 cm3

= 2.70 g / cm3

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Air has a density of 0.0013 g/mL. What is the mass of 6.0-L sample of air?

Calculate the mass in grams of 10.0 mL if mercury (Hg) if the density of Hg is 13.6 g/mL.

Calculate the volume in milliliters, of a liquid that has a density of 1.20 g/mL and a mass of 5.00 grams.

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(g/mL) water ofdensity

(g/mL)object ofdensity gravity specific

Specific Gravity

• Values of density are often related to a standard

• Specific gravity - the ratio of the density of the object in question to the density of pure water at 4oC

• Specific gravity is a unitless term because the 2 units cancel

• Often the health industry uses specific gravity to test urine and blood samples