01 Lecture KM

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1

Chem istry : Atoms First

Julia Burdge & Jason Overby

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

Chapter 1

Chemistry :

The Science of Change

Kent L. McCorkle

Cosumnes River College

Sacramento, CA

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Chemis try: The Science of Change1

1.1 The Study of ChemistryChemistry You May Already Know

The Scientific Method

1.2 Classif ication of Matter

States of Matter Triple Point

Mixtures

1.3 The Properties of Matter

Physical Properties

Chemical Properties

Extensive and Intensive Properties

1.4 Scienti f ic Measurement

SI Base UnitsAnimationMass

TemperatureCritical Thinking

Derived Units: Volume and Density

1.5 Uncertainty in MeasurementSignificant FiguresHandout

Calculations with Measured Numbers

Accuracy and Precision

1.6 Using Uni ts and Solving Problems

Conversion Factors

Dimensional Analysis

Tracking Units

Gasoli ne Project

Sample Spreadsheet
http://www.fccj.us/Lab/ControlledExperimentLab2.htmhttp://www.youtube.com/watch?v=BLRqpJN9zeAhttp://www.lsua.info/MetricSystem/MetricPrefix.htmlhttp://www.lsua.info/chem1001/Chap2-3Movies/metric.htmlhttp://www.lsua.info/mathworkshop1/frametemp1.htmlhttp://www.northcampus.net/CHM1025/CriticalThinking/CriticalThinkingTempConversionExercise.htmhttp://www.lsua.info/chem1001/Chap2-3Movies/sigdigit.htmlhttp://www.northcampus.net/CHM1025/Lab/SignificantFiguresintheLaboratorySP2012.htmhttp://www.lsua.info/chem1001/dimanalysis/unitanalysisIntro.htmhttp://www.fscj.me/gasoline/GasolineProjectLab.htmhttp://www.fscj.me/gasoline/Gas_MileageSample.htmhttp://www.fscj.me/gasoline/Gas_MileageSample.htmhttp://www.fscj.me/gasoline/GasolineProjectLab.htmhttp://www.lsua.info/chem1001/dimanalysis/unitanalysisIntro.htmhttp://www.northcampus.net/CHM1025/Lab/SignificantFiguresintheLaboratorySP2012.htmhttp://www.lsua.info/chem1001/Chap2-3Movies/sigdigit.htmlhttp://www.northcampus.net/CHM1025/CriticalThinking/CriticalThinkingTempConversionExercise.htmhttp://www.lsua.info/mathworkshop1/frametemp1.htmlhttp://www.lsua.info/chem1001/Chap2-3Movies/metric.htmlhttp://www.lsua.info/MetricSystem/MetricPrefix.htmlhttp://www.youtube.com/watch?v=BLRqpJN9zeAhttp://www.fccj.us/Lab/ControlledExperimentLab2.htm

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The Study of Chem istry

Chemistryis the study of matter and the changes that matter undergoes.

Matteris anything that has mass and occupies space.

1.1

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The Study of Chemistry

Scientists follow a set of guidelines known as the scienti f ic method:

gather data via observations and experiments

identify patterns or trends in the collected data

summarize their findings with a law

formulate a hypothesis

with time a hypothesis may evolve into a theory

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Class i f icat ion of Matter

Chemists classify matter as either asubstanceor a mixtureof substances.

A substanceis a form of matter that has definite composition and distinct properties. Examples: salt (sodium chloride), iron, water, mercury, carbon dioxide, and

oxygen

Substances differ from one another in composition and may be identified by

appearance, smell, taste, and other properties.

1.2

A mixtureis a physical combination of two or more substances.

A homogeneousmixtureis uniform throughout. Also called a solution.

Examples: seawater, apple juice

Aheterogeneous mixtur eis notuniform throughout.

Examples: trail mix, chicken noodle soup

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Class i f icat ion o f Matter

All substances can, in principle,exist as a solid, liquid or gas.

We can convert a substance from

one state to another without

changing the identity of the

substance.

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Class i f icat ion o f Matter

Solid particles are held

closely together in an

ordered fashion.

Liquid particles are close

together but are not held

rigidly in position.

Gas particles have

significant separation

from each other and

move freely.

Solids do not conform

to the shape of their

container.

Liquids do conform to

the shape of their

container.

Gases assume both the

shape and volume of

their container.

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Class if ication of Matter

A mixture can be separated by physical means into its components

without changing the identities of the components.

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The Properties of Matter

There are two general types of properties of matter:

1)Quanti tative properties are measured and expressed with a

number.

2) Qualitative properties do not require measurement and are

usually based on observation.

1.3

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The Properties o f Matter

A physical propertyis one that can be observed and measuredwithout changing the identity of the substance.

Examples: color, melting point, boiling point

A physical changeis one in which the state of matter changes, but

the identity of the matter does not change.

Examples: changes of state (melting, freezing, condensation)

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A chemical propertyis one a substance exhibits as it interacts withanother substance.

Examples: flammability, corrosiveness

A chemical changeis one that results in a change of composition;

the original substances no longer exist.

Examples: digestion, combustion, oxidation

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An extensive propertydepends on the amount of matter.Examples: mass, volume

An intensive propertydoes notdepend on the amount of matter.Examples: temperature, density

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Scient i f ic Measurement1.4

Properties that can be measured

are called quantitative

properties.

A measured quantity must

always include a unit.TheEnglish systemhas units

such as the foot, gallon, pound,

etc.

The metric systemincludesunits such as the meter, liter,

kilogram, etc.

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SI Base Units

The revised metric system is called the I nternational System of

Units(abbreviated SI Uni ts) and was designed for universal use byscientists.

There are seven SI base units

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SI Base Units

The magnitude of a unit may be tailored to a particular application

using prefixes.

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Mass

Massis a measure of the amount of matter in an object or sample.

Because gravity varies from location to location, the weight of an

object varies depending on where it is measured. But mass doesnt

change.

The SI base unit of mass is the kilogram (kg), but in chemistry the

smaller gram (g) is often used.

1 kg = 1000 g = 1103g

Atomic mass unit (amu)is used to express the masses of atoms andother similar sized objects.

1 amu = 1.660537810-24g

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Temperature

There are two temperature scales used in chemistry:

The Celsius scale (C)

Freezing point (pure water): 0C

Boiling point (pure water): 100C

The Kelvin scale (K)

The absolute scale

Lowest possible temperature: 0 K (absolute zero)

K = C + 273.15

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Normal human body temperature can range over the course of a day from about

36C in the early morning to about 37C in the afternoon. Express these two

temperatures and the range that they span using the Kelvin scale.

Worked Example 1.1

Strategy Use K = C + 273.15 to convert temperatures from Celsius to Kelvin.

Solution 36C + 273 = 309 K

37C + 273 = 310 K

What range do they span?

310 K - 309 K = 1 K

Depending on the precision

required, the conversion from

C to K is often simply done by

adding 273, rather than 273.15.

Think About It Remember that converting a temperature from C to K is

different from converting a range or differencein temperature from C to K.

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The Fahrenheit scale is common in the United States.

Freezing point (pure water): 32C

Boiling point (pure water): 212C

There are 180 degrees between freezing and boiling in Fahrenheit

(212F-32F) but only 100 degrees in Celsius (100C-0C). The size of a degree on the Fahrenheit scale is only of a

degree on the Celsius scale.

Temperature

Temp in F = ( temp in C ) + 32F

5

9

5

9

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A body temperature above 39C constitutes a high fever. Convert this temperature

to the Fahrenheit scale.

Worked Example 1.2

Solution Temp in F = ( 39C ) + 32F

Temp inF = 102

F

Think About It Knowing that normal body temperature on the Fahrenheitscale is approximately 98.6F, 102F seems like a reasonable answer.

Temp in F = ( temp in C ) + 32F5

9

5

9

Strategy We are given a temperature and asked to convert it to degrees

Fahrenheit. We will use the equation below:

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Derived Units: Volume and Dens ity

There are many units (such as

volume) that require units not

included in the base SI units.

The derived SI unit for volume is

the meter cubed (m3).

A more practical unit for volume is

the liter(L).

1 dm3= 1 L

1 cm3= 1 mL

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d = density

m = mass

V= volume

SI-derived unit: kilogram per cubic meter (kg/m3)

Other common units: g/cm3 (solids)

g/mL (liquids)g/L (gases)

Derived Units: Volume and Dens ity

The densityof a substance is the ratio of mass to volume.

md =

V

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Ice cubes float in a glass of water because solid water is less dense than liquid

water. (a) Calculate the density of ice given that, at 0C, a cube that is 2.0 cm oneach side has a mass of 7.36 g, and (b) determine the volume occupied by 23 g of

ice at 0C.

Solution (a) A cube has three equal sides so the volume is (2.0 cm)3, or 8.0 cm3

d=

(b) Rearranging d = m/Vto solve for volume gives V = m/d

V=

Think About It For a sample with a density lessthan 1 g/cm3, the number

of cubic centimeters should begreaterthan the number of grams. In this

case, 25 cm3> 23 g.

7.36 g

8.0 cm3

23 g

0.92 g/cm3

= 0.92 g/cm3

= 25 cm3

Worked Example 1.3

Strategy (a) Determine density by dividing mass by volume, and (b) use the

calculated density to determine the volume occupied by the given mass.

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Uncertainty in Measurement

There are two types of numbers used in chemistry:

1)Exactnumbers:

a) are those that have defined values

1 kg = 1000 g

1 dozen = 12 objects

b) are those determined by counting

28 students in a class

2)Inexactnumbers:a) measured by any method other than counting

length, mass, volume, time, speed, etc.

1.5

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Uncertainty in Measu rement

An inexact number must be reported so as to indicate its uncertainty.

Signi f icant figuresare the

meaningful digitsin a reported

number.

The last digit in a measured numberis referred to as the uncertain digit.

When using the top ruler to measure

the memory card, we could estimate

2.5 cm. We are certain about the 2,

but we are notcertain about the 5.

The uncertainty is generally

considered to be + 1 in the last digit.

2.5 + 0.1 cm

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Uncertainty in Measu rement

When using the bottom ruler to

measure the memory card, we might

record 2.45 cm.

Again, we estimate one more digitthan we are certain of.

2.45 + 0.01 cm

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Signif icant Figures

The number of significant figures can be determined using the

following guidelines:

1) Any nonzero digit is significant.

2) Zeros between nonzero digits are significant.

3) Zeros to the left of the first nonzero digit are not significant.

112.1 4 significant figures

305 3 significant figures

0.0023 2 significant figure


0.000001 1 significant figure

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Signif icant Figures

The number of significant figures can be determined using the

following guidelines:

4) Zeros to the right of the last nonzero digit are significant if a

decimal is present.

5) Zeros to the right of the last nonzero digit in a number that does

not contain a decimal point may or may not be significant.


100 1, 2, or 3ambiguous

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Determine the number of significant figures in the following measurements: (a)

443 cm, (b) 15.03 g, (c) 0.0356 kg, (d) 3.00010-7L, (e) 50 mL, (f) 0.9550 m.

Worked Example 1.4

Solution (a) 443 cm (b)

15.03 g

(c) 0.0356 kg (d)

3.000 x 10-7L

(e) 50 mL (f)

0.9550 m

Strategy Zeros are significant between nonzero digits or after a nonzero digit

with a decimal. Zeros may or may not be significant if they appear to the right of

a nonzero digit without a decimal.

3 S.F. 4 S.F.

3 S.F. 4 S.F.

1 or 2, ambiguous 4 S.F.

Think About It Be sure that you have identified zeros correctly as either

significant or not significant. They are significant in (b) and (d); they are not

significant in (c); it is not possible to tell in (e); and the number in (f)

contains one zero that is significant, and one that is not.

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In addition and subtraction, the answer cannot have more digits to

the right of the decimal point than any of the original numbers.

102.50

+ 0.231102.731

143.29

- 20.1

123.19

round to two digits after the decimal point, 102.73

round to one digit after the decimal point, 123.2

two digits after the decimal point

three digits after the decimal point

two digits after the decimal point

one digit after the decimal point


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In multiplication and division, the number of significant figures in

the final product or quotient is determined by the original numberthat has the smallest number of significant figures.

1.48.011 = 11.2154

11.57/305.88 = 0.0378252

2 S.F.

fewest significant figures is 2, so

round to 114 S.F.

fewest significant figures is 4, so

round to 0.037834 S.F. 5 S.F.


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Exact numbers can be considered to have an infinite number of

significant figures and do not limit the number of significant figures

in a result.

Example: Three pennies each have a mass of 2.5 g. What is the

total mass?32.5 = 7.5 g


Exact

(counting number)

Inexact

(measurement)

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In calculations with multiple steps, round at the end of the

calculation to reduce any rounding errors.

Do not round after each step.

Compare the following:


1) 3.668.45 = 30.9

2) 30.92.11 = 65.2

1) 3.668.45 = 30.93

2) 30.932.11 = 65.3

Rounding after each step Rounding at end

In general, keep at least one extra digit until the end of a multistep

calculation.

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Perform the following arithmetic operations and report the result to the proper

number of significant figures: (a) 317.5 mL + 0.675 mL, (b) 47.80 L2.075 L,

(c) 13.5 g45.18 L, (d) 6.25 cm x 1.175 cm, (e) 5.46x102g4.991x103g

Worked Example 1.5

Solution (a) 317.5 mL+ 0.675 mL

318.175 mL

(b) 47.80 L

- 2.075 L45.725 L

Strategy Apply the rules for significant figures in calculations, and round each

answer to the appropriate number of digits.

round to 318.2 mL

round to 45.73 L

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(c) 13.5 g45.18 L, (d) 6.25 cm x 1.175 cm, (e) 5.46x102g4.991x103g

Worked Example 1.5 (cont .)

Solution

(c) 13.5 g

45.18 L

(d) 6.25 cm1.175 cm



round to 0.299 g/L= 0.298804781 g/L

3 S.F.

4 S.F.

round to 7.34 cm2= 7.34375 cm2

3 S.F. 4 S.F.

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(c) 13.5 g45.18 L, (d) 6.25 cm x 1.175 cm, (e) 5.46x102g4.991x103g

Worked Example 1.5 (cont .)

Solution (e) 5.46 x 102g

+ 49.91 x 102g

55.37 x 102g



= 5.537 x 103g

Think About It Changing the answer to correct scientific notation doesntchange the number of significant figures, but in this case it changes the number of

places past the decimal place.

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An empty container with a volume of 9.850 x 102 cm3is weighed and found to

have a mass of 124.6 g. The container is filled with a gas and reweighed. Themass of the container and the gas is 126.5 g. Determine the density of the gas to

the appropriate number of significant figures.

Worked Example 1.6

Solution 126.5 g

124.6 g

mass of gas = 1.9 g

density =

Strategy This problem requires two steps: subtraction to determine the mass of

the gas, and division to determine its density. Apply the corresponding rule

regarding significant figures to each step.

one place past the decimal point (two sig figs)

1.9 g

9.850 x 102cm3 round to 0.0019 g/cm3= 0.00193 g/cm3

Think About It In this case, although each of the three numbers we started

with hasfoursignificant figures, the solution only has twosignificant figures.

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Accuracy and Precis ion

Accuracytells us how close a

measurement is to the truevalue.

Precisiontells us how close a series of

replicate measurements are to one another.

Good accuracy and good precision

Poor accuracy but good precision

Poor accuracy and poor precision

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Accuracy and Precis ion

Three students were asked to find the mass of an aspirin tablet. The

true mass of the tablet is 0.370 g.

Student A: Results are precise but not accurate

Student B: Results are neither precise nor accurate

Student C: Results are both precise and accurate

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Using Uni ts and Solv ing Prob lems

A conversion factor is a fraction in which the same quantity is

expressed one way in the numerator and another way in the

denominator.

For example, 1 in = 2.54 cm, may be written:

1.6

1 in

2.54 cm

2.54 cm

1 in

or

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The Food and Drug Administration (FDA) recommends that dietary sodium

intake be no more than 2400 mg per day.

Worked Example 1.7

Solution

2400 mg

Strategy The necessary conversion factors are derived from the equalities

1 g = 1000 mg and 1 lb = 453.6 g.

1 lb

453.6 g

453.6 g

1 lbor

1 g

1000 mgor

1000 mg

1 g

1 g

1000 mg1 lb

453.6 g = 0.005291 lb

Think About It Make sure that the magnitude of the result is reasonable and

that the units have canceled properly. If we had mistakenly multiplied by 1000

and 453.6 instead of dividing by them, the result

(2400 mg1000 mg/g453.6 g/lb = 1.089109mg2/lb) would be

unreasonably large and the units would not have canceled properly.

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An average adult has 5.2 L of blood. What is the volume of blood in cubic

meters?

Worked Example 1.8

Solution

5.2 L

Strategy 1 L = 1000 cm3and 1 cm = 1x10-2m. When a unit is raised to a

power, the corresponding conversion factor must also be raised to that power in

order for the units to cancel appropriately.

1000 cm3

1 L

1 x 10-2m

1 cm = 5.2 x 10-3m3

3

Think About It Based on the preceding conversion factors, 1 L = 110-3m3.

Therefore, 5 L of blood would be equal to 510-3m3, which is close to the

calculated answer.

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Chapter Summary: Key Points1

The Scientific MethodStates of Matter

Substances

Mixtures

Physical Properties

Chemical PropertiesExtensive and Intensive Properties

SI Base Units

Mass

TemperatureVolume and Density

Significant Figures

Documents

01 Lecture KM