CHAPTER ONE

1
CHAPTER ONE

• The Foundations of Chemistry

2
Chapter Outline

• Matter and Energy
• States of Matter
• Chemical and Physical Properties
• Chemical and Physical Changes
• Mixtures, Substances, Compounds, and
  Elements
• Measurements in Chemistry
• Units of Measurement

3
Chapter Outline

1. Use of Numbers
2. The Unit Factor Method (Dimensional Analysis)
3. Percentage
4. Density and Specific Gravity
5. Heat and Temperature
6. Heat Transfer and the Measurement of Heat

4
Matter and Energy - Vocabulary

• Chemistry
• Science that describes matter its properties,
  the changes it undergoes, and the energy changes
  that accompany those processes
• Matter
• Anything that has mass and occupies space.
• Energy
• The capacity to do work or transfer heat.
• Scientific (natural) law
• A general statement based the observed behavior
  of matter to which no exceptions are known.

5
Natural Laws

• Law of Conservation of Mass
• Law of Conservation of Energy
• Law of Conservation of Mass-Energy
• Einsteins Relativity
• Emc2

6
Scientific Method

• Observation
• Hypothesis
• Observation or experiment
• Theory
• Observation or experiment
• Law

7
States of Matter

• Solids

8
States of Matter

• Solids
• Liquids

9
States of Matter

• Solids
• Liquids
• Gases

10
States of Matter

• Change States
• heating
• cooling

11
States of Matter

• Illustration of changes in state
• requires energy

12
Chemical and Physical Properties

• Chemical Properties - chemical changes
• rusting or oxidation
• chemical reactions
• Physical Properties - physical changes
• changes of state
• density, color, solubility
• Extensive Properties - depend on quantity
• Intensive Properties - do not depend on quantity

13
Mixtures, Substances, Compounds, and Elements

• Substance
• matter in which all samples have identical
  composition and properties
• Elements
• substances that cannot be decomposed into simpler
  substances via chemical reactions
• Elemental symbols
• found on periodic chart

14
Mixtures, Substances, Compounds, and Elements
15
Mixtures, Substances, Compounds, and Elements

• Compounds
• substances composed of two or more elements in a
  definite ratio by mass
• can be decomposed into the constituent elements
• Water is a compound that can be decomposed into
  simpler substances hydrogen and oxygen

16
Mixtures, Substances, Compounds, and Elements
17
Mixtures, Substances, Compounds, and Elements

• Mixtures
• composed of two or more substances
• homogeneous mixtures
• heterogeneous mixtures

18
Measurements in Chemistry

• Quantity Unit Symbol
• length meter m
• mass kilogram kg
• time second s
• current ampere A
• temperature Kelvin K
• amt. substance mole mol

19
Measurements in ChemistryMetric Prefixes

• Name Symbol Multiplier
• mega M 106
• kilo k 103
• deka da 10
• deci d 10-1
• centi c 10-2

20
Measurements in ChemistryMetric Prefixes

• Name Symbol Multiplier
• milli m 10-3
• micro ? 10-6
• nano n 10-9
• pico p 10-12
• femto f 10-15

21
Units of Measurement

• Definitions
• Mass
• measure of the quantity of matter in a body
• Weight
• measure of the gravitational attraction for a
  body

22
Units of Measurement

• Common Conversion Factors
• Length
• 1 m 39.37 inches
• 2.54 cm 1 inch
• Volume
• 1 liter 1.06 qt
• 1 qt 0.946 liter
• See Table 1-7 for more conversion factors

23
Use of Numbers

• Exact numbers
• 1 dozen 12 things for example
• Accuracy
• how closely measured values agree with the
  correct value
• Precision
• how closely individual measurements agree with
  each other

24
Use of Numbers

• Significant figures
• digits believed to be correct by the person
  making the measurement
• Measure a mile with a 6 inch ruler vs. surveying
  equipment
• Exact numbers have an infinite number of
  significant figures
• 12.000000000000000 1 dozen
• because it is an exact number

25
Use of Numbers

• Significant Figures - Rules
• Leading zeroes are never significant
• 0.000357 has three significant figures
• Trailing zeroes may be significant
• must specify significance by how the number is
  written
• 1300 nails - counted or weighed?
• Use scientific notation to remove doubt
• 2.40 x 103 has ? significant figures

26
Use of Numbers

• Scientific notation for logarithms
• take the log of 2.40 x 103
• log(2.40 x 103) 3.380
• How many significant figures?
• Imbedded zeroes are always significant
• 3.0604 has five significant figures

27
Use of Numbers

• Piece of Black Paper with rulers beside the
  edges

28
Use of Numbers

• Piece of Paper Side B enlarged
• How long is the paper to the best of your ability
  to measure it?

29
Use of Numbers

• Piece of Paper Side A enlarged
• How wide is the paper to the best of your ability
  to measure it?

30
Use of Numbers

• Determine the area of the piece of black paper
  using your measured values.
• Compare your answer with your classmates.
• Where do your answers differ in the numbers?
• Significant figures rules for multiplication and
  division must help us determine where answers
  would differ.

31
Use of Numbers

• Multiplication Division rule
• Easier of the two rules
• Product has the smallest number of significant
  figures of multipliers

32
Use of Numbers

• Multiplication Division rule
• Easier of the two rules
• Product has the smallest number of significant
  figures of multipliers

33
Use of Numbers

• Multiplication Division rule
• Easier of the two rules
• Product has the smallest number of significant
  figures of multipliers

34
Use of Numbers

• Determine the perimeter of the piece of black
  paper using your measured values.
• Compare your answer with your classmates.
• Where do your answers differ in the numbers?
• Significant figures rules for addition and
  subtraction must help us determine where answers
  would differ.

35
Use of Numbers

• Addition Subtraction rule
• More subtle than the multiplication rule
• Answer contains smallest decimal place of the
  addends.

36
Use of Numbers

• Addition Subtraction rule
• More subtle than the multiplication rule
• Answer contains smallest decimal place of the
  addends.

37
Use of Numbers

• Addition Subtraction rule
• More subtle than the multiplication rule
• Answer contains smallest decimal place of the
  addends.

38
The Unit Factor Method

• Simple but important method to get correct
  answers in word problems.
• Method to change from one set of units to
  another.
• Visual illustration of the idea.

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The Unit Factor Method

• Change from a to a by obeying the
  following rules.

40
The Unit Factor Method

• Change from a to a by obeying the
  following rules.
• Must use colored fractions.

41
The Unit Factor Method

• Change from a to a by obeying the
  following rules.
• Must use colored fractions.
• The box on top of the fraction must be the same
  color as the next fractions bottom box.

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The Unit Factor Method
R

• Fractions to choose from

R
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The Unit Factor Method
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• Fractions to choose from

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The Unit Factor Method
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• Fractions to choose from

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The Unit Factor Method
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• Fractions to choose from

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The Unit Factor Method
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• Fractions to choose from

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The Unit Factor Method
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• Fractions to choose from

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The Unit Factor Method
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• Fractions to choose from

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The Unit Factor Method

• colored fractions represent unit factors
• 1 ft 12 in becomes or
• Example 1-1 Express 9.32 yards in millimeters.

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The Unit Factor Method
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The Unit Factor Method
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The Unit Factor Method
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The Unit Factor Method
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The Unit Factor Method
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The Unit Factor Method

• Example 1-2 Express 627 milliliters in gallons.
• You do it!

56
The Unit Factor Method

• Example 1-2. Express 627 milliliters in gallons.

57
The Unit Factor Method

• Area is two dimensional thus units must be in
  squared terms.
• Example 1-3 Express 2.61 x 104 cm2 in ft2.

58
The Unit Factor Method

• Area is two dimensional thus units must be in
  squared terms.
• Example 1-3 Express 2.61 x 104 cm2 in ft2.
• common mistake

59
The Unit Factor Method

• Area is two dimensional thus units must be in
  squared terms.
• Example 1-3 Express 2.61 x 104 cm2 in ft2.

O
R
P
60
The Unit Factor Method

• Area is two dimensional thus units must be in
  squared terms.
• Example 1-3 Express 2.61 x 104 cm2 in ft2.

O
R
R
61
The Unit Factor Method

• Area is two dimensional thus units must be in
  squared terms.
• Example 1-3 Express 2.61 x 104 cm2 in ft2.

62
The Unit Factor Method

• Area is two dimensional thus units must be in
  squared terms.
• Example 1-3 Express 2.61 x 104 cm2 in ft2.

63
The Unit Factor Method

• Volume is three dimensional thus units must be in
  cubic terms.
• Example 1-4 Express 2.61 ft3 in cm3.
• You do it!

64
The Unit Factor Method

• Volume is three dimensional thus units must be in
  cubic terms.
• Example 1-4 Express 2.61 ft3 in cm3.

65
Percentage

• Percentage is the parts per hundred of a sample.
• Example 1-5 A 335 g sample of ore yields 29.5 g
  of iron. What is the percent of iron in the ore?
• You do it!

66
Percentage

• Percentage is the parts per hundred of a sample.
• Example 1-5 A 335 g sample of ore yields 29.5 g
  of iron. What is the percent of iron in the ore?

67
Density and Specific Gravity

• density mass/volume
• What is density?
• Why does ice float in liquid water?

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Density and Specific Gravity

• density mass/volume
• What is density?
• Why does ice float in liquid water?

69
Density and Specific Gravity

• Example 1-6 Calculate the density of a substance
  if 742 grams of it occupies 97.3 cm3.

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Density and Specific Gravity

• Example 1-6 Calculate the density of a substance
  if 742 grams of it occupies 97.3 cm3.

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Density and Specific Gravity

• Example 1-7 Suppose you need 125 g of a corrosive
  liquid for a reaction. What volume do you need?
• liquids density 1.32 g/mL
• You do it!

72
Density and Specific Gravity

• Example 1-7 Suppose you need 125 g of a corrosive
  liquid for a reaction. What volume do you need?
• liquids density 1.32 g/mL

73
Density and Specific Gravity

• Example 1-7 Suppose you need 125 g of a corrosive
  liquid for a reaction. What volume do you need?
• liquids density 1.32 g/mL

74
Density and Specific Gravity

• Waters density is essentially 1.00 at room T.
• Thus the specific gravity of a substance is very
  nearly equal to its density.
• Specific gravity has no units.

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Density and Specific Gravity

• Example 1-8 A 31.0 gram piece of chromium is
  dipped into a graduated cylinder that contains
  5.00 mL of water. The water level rises to 9.32
  mL. What is the specific gravity of chromium?
• You do it

76
Density and Specific Gravity

• Example1-8 A 31.0 gram piece of chromium is
  dipped into a graduated cylinder that contains
  5.00 mL of water. The water level rises to 9.32
  mL. What is the specific gravity of chromium?

77
Density and Specific Gravity

• Example1-8 A 31.0 gram piece of chromium is
  dipped into a graduated cylinder that contains
  5.00 mL of water. The water level rises to 9.32
  mL. What is the specific gravity of chromium?

78
Density and Specific Gravity

• Example 1-9 A concentrated hydrochloric acid
  solution is 36.31 HCl and 63.69 water by mass.
  The specific gravity of the solution is 1.185.
  What mass of pure HCl is contained in 175 mL of
  this solution?
• You do it!

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Density and Specific Gravity
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Density and Specific Gravity
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Density and Specific Gravity
82
Try This One

• Battery acid is 40.0 sulfuric acid and 60 water
  by mass. Its specific gravity is 1.31.

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Solution

• From the given specific value number 1.31
• We may write Density 1.31g/mL
• Next

84

• The solution is 40 sulfuric acid and 60 water
  by massfrom this information we get

Because 100g of solution contains 40.0 g of
sulfuric acid
sulfuric acid
solution
85
We can now solve the problem

• __g H2SO4

86
Heat and Temperature

• Heat and Temperature are not the same thing
• T is a measure of the intensity of heat in a body
• 3 common temperature scales - all use water as a
  reference

87
Heat and Temperature

• Heat and Temperature are not the same thing
• T is a measure of the intensity of heat in a body
• 3 common temperature scales - all use water as a
  reference

88
Heat and Temperature

• MP water BP water
• Fahrenheit 32 oF 212 oF
• Celsius 0.0 oC 100 cC
• Kelvin 273 K 373 K

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Relationships of the Three Temperature Scales
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Relationships of the Three Temperature Scales
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Relationships of the Three Temperature Scales
92
Relationships of the Three Temperature Scales

• Easy method to remember how to convert from
  Centigrade to Fahrenheit.
• Double the Centigrade temperature.
• Subtract 10 of the doubled number.
• Add 32.

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Heat and Temperature

• Example 1-10 Convert 211oF to degrees Celsius.

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Heat and Temperature

• Example 1-11 Express 548 K in Celsius degrees.

95
Heat Transfer and The Measurement of Heat

• SI unit J (Joule)
• calorie
• Amount of heat required to heat 1 g of water 1 oC
• 1 calorie 4.184 J
• Calorie
• Large calorie, kilocalorie, dietetic calories
• Amount of heat required to heat 1 kg of water 1
  oC
• English unit BTU
• Specific Heat
• amount of heat required to raise the T of 1g of a
  substance by 1oC
• units J/goC

96
Heat Transfer and the Measurement of Heat

• Heat capacity
• amount of heat required to raise the T of 1 mole
  of a substance by 1oC
• units J/mol oC
• Example 1-12 Calculate the amt. of heat to raise
  T of 200.0 g of water from 10.0oC to 55.0oC

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Heat Transfer and the Measurement of Heat

• Heat transfer equation
• necessary to calculate amounts of heat
• amount of heat amount of substance x
  specific heat x?DT

98
Heat Transfer and the Measurement of Heat

• Heat transfer equation
• necessary to calculate amounts of heat
• amount of heat amount substance x specific
  heat x?DT

99
Heat Transfer and the Measurement of Heat

• Heat transfer equation
• necessary to calculate amounts of heat
• amount of heat amount substance x specific
  heat x??T

100
Heat Transfer and the Measurement of Heat

• Example 1-13 Calculate the amount of heat to
  raise the temperature of 200.0 grams of mercury
  from 10.0oC to 55.0oC. Specific heat for Hg is
  0.138 J/g oC.
• You do it!

101
Heat Transfer and the Measurement of Heat

• Example 1-13 Calculate the amount of heat to
  raise T of 200.0 g of Hg from 10.0oC to 55.0oC.
  Specific heat for Hg is 0.138 J/g oC.
• Requires 30.3 times more heat for water
• 4.184 is 30.3 times greater than 0.138

102
Heating Curve for 3 Substances
Which substance has the largest specific heat?
Which substances T will decrease the most after
the heat has been removed?
103
Heating Curve for 3 Substances