Introductory Chemistry, 3rd Edition
Nivaldo Tro
Chapter 3
Matter and Energy
Basic Principles of Chemistry Online
Southeast Missouri State University
Cape Girardeau, MO
2009, Prentice Hall
In Your Room
• Everything you can see,
touch, smell or taste in
your room is made of
matter.
• Chemists study the
differences in matter and
how that relates to the
structure of matter.
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What Is Matter?
• Matter is defined as
anything that occupies
space and has mass.
• Even though it appears to
be smooth and continuous,
matter is actually composed
of a lot of tiny little pieces
we call atoms and
molecules.
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Atoms and Molecules
• Atoms are the tiny particles
that make up all matter.
• In most substances, the
atoms are joined together in
units called molecules.
The atoms are joined in
specific geometric
arrangements.
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Structure Determines Properties
• The properties of matter are determined by the atoms and
molecules that compose it.
1.
2.
3.
4.
Carbon Monoxide
Composed of one carbon atom
and one oxygen atom.
Colorless, odorless gas.
Burns with a blue flame.
Binds to hemoglobin.
1.
2.
3.
4.
Carbon Dioxide
Composed of one carbon atom
and two oxygen atoms.
Colorless, odorless gas.
Incombustible.
Does not bind to hemoglobin.
5
Classifying Matter
by Physical State
• Matter can be classified as solid, liquid, or
gas based on what properties it exhibits.
State
Shape
Volume
Compress
Flow
Solid
Fixed
Fixed
No
No
Liquid
Indefinite
Fixed
No
Yes
Gas
Indefinite
Indefinite
Yes
Yes
•Fixed = Property doesn’t change when placed in a container.
•Indefinite = Takes the property of the container.
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Structure Determines Properties
• The atoms or molecules have different
structures in solids, liquids, and gases −
leading to different properties.
7
Solids
• The particles in a solid are packed
close together and are fixed in
position.
 Although they may vibrate.
• The close packing of the particles
results in solids being
incompressible.
• The inability of the particles to
move around results in solids
retaining their shape and volume
when placed in a new container
and prevents the particles from
flowing.
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Solids, Continued
• Some solids have their particles
arranged in an orderly geometric
pattern—we call these crystalline
solids.
Salt and diamonds.
• Other solids have particles that do
not show a regular geometric
pattern over a long range—we call
these amorphous solids.
Plastic and glass.
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Liquids
• The particles in a liquid are closely packed,
but they have some ability to move around.
• The close packing results in liquids being
incompressible.
• The ability of the particles to move allows
liquids to take the shape of their container
and to flow. However, they don’t have
enough freedom to escape and expand to fill
the container.
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Gases
• In the gas state, the particles have complete
freedom from each other.
• The particles are constantly flying around,
bumping into each other and the container.
• In the gas state, there is a lot of empty space
between the particles.
On average.
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Gases, Continued
• Because there is a lot of empty
space, the particles can be
squeezed closer together.
Therefore, gases are
compressible.
• Because the particles are not
held in close contact and are
moving freely, gases expand to
fill and take the shape of their
container, and will flow.
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Classification of Matter
by Appearance
• Homogeneous = Matter that is uniform throughout.
 Appears to be one thing.
 Every piece of a sample has identical properties, though another sample
with the same components may have different properties.
 Solutions (homogeneous mixtures) and pure substances.
• Heterogeneous = Matter that is non-uniform throughout .
 Contains regions with different properties than other regions.
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Practice—Classify the Following as
Homogeneous or Heterogeneous
• Table sugar.
• A mixture of table sugar and black pepper.
• A mixture of sugar dissolved in water.
• Oil and vinegar salad dressing.
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Practice—Classify the Following as
Homogeneous or Heterogeneous,
Continued
• Table sugar = homogeneous
• A mixture of table sugar and black pepper =
heterogeneous
• A mixture of sugar dissolved in water =
homogeneous
• Oil and vinegar salad dressing = heterogeneous
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Classifying Matter
by Composition
• Matter that is composed of only one kind of
atom or molecule is called a pure substance.
• Matter that is composed of different kinds of
atoms or molecules is called a mixture.
• Because pure substances always have only one
kind of piece, all samples show the same
properties.
• However, because mixtures have variable
composition, different samples will show
different properties.
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Copper—A Pure Substance
•
•
•
•
•
Color is brownish red.
Shiny, malleable, and ductile.
Excellent conductor of heat and electricity.
Melting point = 1084.62 °C
Density = 8.96 g/cm3 at 20 °C
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Brass—A Mixture
Type
Color
% Cu
% Zn Density
g/cm3
MP
°C
Tensile
Strength
psi
Uses
Gilding
reddish
95
5
8.86
1066
50K
pre-83 pennies,
munitions, and
plaques
Commercial
bronze
90
10
8.80
1043
61K
door knobs and
grillwork
Jewelry
bronze
87.5
12.5
8.78
1035
66K
costume jewelry
Red
golden
85
15
8.75
1027
70K
electrical sockets,
fasteners, and
eyelets
Low
deep
yellow
80
20
8.67
999
74K
musical instruments
and clock dials
Cartridge
yellow
70
30
8.47
954
76K
car radiator cores
Common
yellow
67
33
8.42
940
70K
lamp fixtures and
bead chain
Muntz metal
yellow
60
40
8.39
904
70K
nuts and bolts
Classification of Matter
Matter
Pure Substance
Constant Composition
Homogeneous
Mixture
Variable Composition
• Pure Substance = All samples are made of the same
pieces in the same percentages.
 Salt
• Mixtures = Different samples may have the same pieces in
different percentages.
 Salt water
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Pure Substances vs. Mixtures
1.
2.
3.
4.
5.
Pure Substances
All samples have the same
physical and chemical
properties.
Constant composition = All
samples have the same
pieces in the same
percentages.
Homogeneous.
Separate into components
based on chemical
properties.
Temperature stays constant
while melting or boiling.
1.
2.
3.
4.
5.
Mixtures
Different samples may show
different properties.
Variable composition =
Samples made with the same
pure substances may have
different percentages.
Homogeneous or
heterogeneous.
Separate into components
based on physical
properties.
Temperature usually
changes while melting or
boiling because composition
changes.
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Practice—Classify the Following as
Pure Substances or Mixtures
• A homogeneous liquid whose temperature stays
constant while boiling.
• Granite—a rock with several visible minerals in it.
• A red solid that turns blue when heated and
releases water that is always 30% of the solid’s
mass.
• A gas that when cooled and compressed, a liquid
condenses out but some gas remains.
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Practice—Classify the Following as
Pure Substances or Mixtures,
Continued
• A homogeneous liquid whose temperature stays
constant while boiling = pure substance.
• Granite—a rock with several visible minerals in it
= mixture.
• A red solid that turns blue when heated and
releases water that is always 30% of the solid’s
mass = pure substance.
• A gas that when cooled and compressed, a liquid
condenses out but some gas remains = mixture.
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Classification of Pure Substances
• Substances that cannot be broken down into simpler
substances by chemical reactions are called elements.
 Basic building blocks of matter.
 Composed of single type of atom.
 Although those atoms may or may not be combined into molecules.
• Substances that can be decomposed are called compounds.
 Chemical combinations of elements.
 Although properties of the compound are unrelated to the properties of the
elements in it!
 Composed of molecules that contain two or more different kinds
of atoms.
 All molecules of a compound are identical, so all samples of a
compound behave the same way.
• Most natural pure substances are compounds.
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• Atoms
Atoms and Molecules
 Are submicroscopic particles that are the
unit pieces of elements.
 Are the fundamental building blocks of all
matter.
• Molecules
 Are submicroscopic particles that are the
unit pieces of compounds.
 Two or more atoms attached together.
 Attachments are called bonds.
 Attachments come in different strengths.
 Molecules come in different shapes and
patterns.
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Classification of Pure Substances
Elements
Compounds
1. Made of one
type of atom.
(Some elements
are found as
multi-atom
molecules in
nature.)
2. Combine
together to make
compounds.
1. Made of one
type of
molecule, or
array of ions.
2. Molecules
contain 2 or
more different
kinds of atoms.
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Practice—Classify the Following as
Elements or Compounds
• Chlorine, Cl2
• Table sugar, C12H22O11
• A red solid that turns blue when heated and
releases water that is always 30% of the solid’s
mass.
• A brown-red liquid that, when energy is applied to
it in any form, causes only physical changes in the
material, not chemical.
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Practice—Classify the Following as
Elements or Compounds, Continued
• Chlorine, Cl2 = element.
• Table sugar, C12H22O11 = compound.
• A red solid that turns blue when heated and
releases water that is always 30% of the solid’s
mass = compound.
• A brown-red liquid that, when energy is applied to
it in any form, causes only physical changes in the
material, not chemical = element.
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Classification of Mixtures
• Mixtures are generally classified based on
their uniformity.
• Mixtures that are uniform throughout are
called homogeneous.
Also known as solutions.
Mixing is on the molecular level.
• Mixtures that have regions with different
characteristics are called heterogeneous.
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Classification of Mixtures, Continued
Heterogeneous
Homogeneous
1. Made of
multiple
substances,
whose
presence can
be seen.
2. Portions of a
sample have
different
composition
and properties.
1. Made of
multiple
substances, but
appears to be
one substance.
2. All portions of
a sample have
the same
composition
and properties.
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Classifying Matter
30
Properties Distinguish Matter
• Each sample of matter is distinguished by
its characteristics.
• The characteristics of a substance are called
its properties.
• Some properties of matter can be observed
directly.
• Other properties of matter are observed
when it changes its composition.
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Properties of Matter
• Physical Properties are the characteristics of matter
that can be changed without changing its
composition.
 Characteristics that are directly observable.
• Chemical Properties are the characteristics that
determine how the composition of matter changes as
a result of contact with other matter or the influence
of energy.
 Characteristics that describe the behavior of matter.
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Some Physical Properties
Mass
Volume
Density
Solid
Liquid
Gas
Melting point
Boiling point
Volatility
Taste
Odor
Color
Texture
Shape
Solubility
Electrical
conductance
Malleability
Thermal
conductance
Ductility
Magnetism
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Specific heat
capacity
33
Some Physical Properties of Iron
• Iron is a silvery solid at room temperature with a
metallic taste and smooth texture.
• Iron melts at 1538 °C and boils at 4428 °C.
• Iron’s density is 7.87 g/cm3.
• Iron can be magnetized.
• Iron conducts electricity, but not as well as most other
common metals.
• Iron’s ductility and thermal conductivity are about
average for a metal.
• It requires 0.45 J of heat energy to raise the temperature
of one gram of iron by 1°C.
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Some Chemical Properties
Acidity
Causticity
Reactivity
Inertness
(In)Flammability
Oxidizing ability
Basicity (aka alkalinity)
Corrosiveness
Stability
Explosiveness
Combustibility
Reducing ability
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Some Chemical Properties of Iron
• Iron is easily oxidized in
moist air to form rust.
• When iron is added to
hydrochloric acid, it produces
a solution of ferric chloride
and hydrogen gas.
• Iron is more reactive than
silver, but less reactive than
magnesium.
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Practice—Decide Whether Each of the Observations
About Table Salt Is a Physical or Chemical Property
•
•
•
•
•
•
•
Salt is a white, granular solid.
Salt melts at 801 °C.
Salt is stable at room temperature, it does not decompose.
36 g of salt will dissolve in 100 g of water.
Salt solutions and molten salt conduct electricity.
When a clear, colorless solution of silver nitrate is added
to a salt solution, a white solid forms.
When electricity is passed through molten salt, a gray
metal forms at one terminal and a yellow-green gas at the
other.
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Practice − Decide Whether Each of the Observations
About Table Salt Is a Physical or Chemical Property
•
•
•
•
•
•
•
Salt is a white, granular solid = physical.
Salt melts at 801 °C = physical.
Salt is stable at room temperature, it does not decompose =
chemical.
36 g of salt will dissolve in 100 g of water = physical.
Salt solutions and molten salt conduct electricity = physical.
When a clear, colorless solution of silver nitrate is added to a salt
solution, a white solid forms = chemical.
When electricity is passed through molten salt, a gray metal forms
at one terminal and a yellow-green gas at the other = chemical.
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Changes in Matter
• Changes that alter the state or appearance of the
matter without altering the composition are
called physical changes.
• Changes that alter the composition of the matter
are called chemical changes.
During the chemical change, the atoms that are
present rearrange into new molecules, but all of the
original atoms are still present.
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Changes in Matter, Continued
• Physical Changes—Changes in
the properties of matter that do
not effect its composition.
Heating water.
Raises its temperature, but it is still
water.
Evaporating butane from a lighter.
Dissolving sugar in water.
Even though the sugar seems to
disappear, it can easily be separated
back into sugar and water by
evaporation.
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Changes in Matter, Continued
• Chemical Changes involve a change
in the properties of matter that change
its composition.
 A chemical reaction.
 Rusting is iron combining with oxygen to
make iron(III) oxide.
 Burning results in butane from a lighter to
be changed into carbon dioxide and water.
 Silver combines with sulfur in the air to
make tarnish.
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Is it a Physical or Chemical Change?
• A physical change results in a different form of
the same substance.
 The kinds of molecules don’t change.
• A chemical change results in one or more
completely new substances.
 Also called chemical reactions.
 The new substances have different molecules than the
original substances.
 You will observe different physical properties because
the new substances have their own physical properties.
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Phase Changes Are
Physical Changes
•
•
•
•
•
•
•
Boiling = liquid to gas.
Melting = solid to liquid.
Subliming = solid to gas.
Freezing = liquid to solid.
Condensing = gas to liquid.
Deposition = gas to solid.
State changes require heating or cooling the substance.
 Evaporation is not a simple phase change, it is a solution
process.
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Practice—Classify Each Change as Physical
or Chemical
•
•
•
•
•
•
Evaporation of rubbing alcohol.
Sugar turning black when heated.
An egg splitting open and spilling out.
Sugar fermenting.
Bubbles escaping from soda.
Bubbles that form when hydrogen peroxide is
mixed with blood.
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Practice—Classify Each Change as Physical
or Chemical, Continued
• Evaporation of rubbing alcohol = physical.
• Sugar turning black when heated = chemical.
• An egg splitting open and spilling out =
physical.
• Sugar fermenting = chemical.
• Bubbles escaping from soda = physical.
• Bubbles that form when hydrogen peroxide is
mixed with blood = chemical.
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Separation of Mixtures
• Separate mixtures based on different
physical properties of the components.
Physical change.
Different Physical Property
Technique
Boiling point
Distillation
State of matter (solid/liquid/gas)
Filtration
Adherence to a surface
Chromatography
Volatility
Evaporation
Density
Centrifugation and
decanting
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Distillation
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Filtration
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Law of Conservation of Mass
• Antoine Lavoisier
• “Matter is neither created nor destroyed in a
chemical reaction.”
• The total amount of matter present before a
chemical reaction is always the same as the
total amount after.
• The total mass of all the reactants is equal to
the total mass of all the products.
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Conservation of Mass
• Total amount of matter remains constant in a
chemical reaction.
• 58 grams of butane burns in 208 grams of oxygen to
form 176 grams of carbon dioxide and 90 grams of
water.
butane + oxygen  carbon dioxide + water
58 grams + 208 grams  176 grams + 90 grams
266 grams
=
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266 grams
50
Practice—A Student Places Table Sugar and
Sulfuric Acid into a Beaker and Gets a Total
Mass of 144.0 g. Shortly, a Reaction Starts that
Produces a “Snake” of Carbon Extending from
the Beaker and Steam Is Seen Escaping. If the
Carbon Snake and Beaker at the End Have a
Total Mass of 129.6 g, How Much Steam Was
Produced?
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Practice—A Student Places Table Sugar and
Sulfuric Acid into a Beaker and Gets a Total
Mass of 144.0 g. Shortly, a Reaction Starts that
Produces a “Snake” of Carbon Extending from
the Beaker and Steam Is Seen Escaping. If the
Carbon Snake and Beaker at the End Have a
Total Mass of 129.6 g, How Much Steam Was
Produced?
• Total of reactants and beaker = 144.0 g.
• Conservation of mass says total of products and
beaker must be 144.0 g.
• Mass of steam = 144.0 g − 129.6 g = 14.4 g.
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Energy
• There are things that do not have mass and
volume.
• These things fall into a category we call energy.
• Energy is anything that has the capacity to do
work.
• Although chemistry is the study of matter, matter
is effected by energy.
It can cause physical and/or chemical changes in
matter.
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Law of Conservation of Energy
• “Energy can neither be created nor destroyed.”
• The total amount of energy in the universe is
constant. There is no process that can increase
or decrease that amount.
• However, we can transfer energy from one
place in the universe to another, and we can
change its form.
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Matter Possesses Energy
• When a piece of matter
possesses energy, it can
give some or all of it to
another object.
It can do work on the other
object.
• All chemical and physical
changes result in the matter
changing energy.
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Kinds of Energy
Kinetic and Potential
• Potential energy is energy that is
stored.
 Water flows because gravity pulls it
downstream.
 However, the dam won’t allow it to
move, so it has to store that energy.
• Kinetic energy is energy of motion,
or energy that is being transferred
from one object to another.
 When the water flows over the dam,
some of its potential energy is converted
to kinetic energy of motion.
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Some Forms of Energy
• Electrical
 Kinetic energy associated with the flow of electrical
charge.
• Heat or Thermal Energy
 Kinetic energy associated with molecular motion.
• Light or Radiant Energy
 Kinetic energy associated with energy transitions in an
atom.
• Nuclear
 Potential energy in the nucleus of atoms.
• Chemical
 Potential energy in the attachment of atoms or because of
Chemistry",
57
their position. Tro's "Introductory
Chapter 3
Converting Forms of Energy
• When water flows over the dam, some of its
potential energy is converted to kinetic energy.
 Some of the energy is stored in the water because it is
at a higher elevation than the surroundings.
• The movement of the water is kinetic energy.
• Along the way, some of that energy can be used to
push a turbine to generate electricity.
 Electricity is one form of kinetic energy.
• The electricity can then be used in your home.
For example, you can use it to heat cake batter you
mixed, causing it to change chemically and storing
some of the energy in the new molecules that are
made.
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Using Energy
• We use energy to accomplish all kinds of
processes, but according to the Law of
Conservation of Energy we don’t really use
it up!
• When we use energy we are changing it
from one form to another.
For example, converting the chemical energy
in gasoline into mechanical energy to make
your car move.
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“Losing” Energy
• If a process was 100% efficient, we could
theoretically get all the energy transformed
into a useful form.
• Unfortunately we cannot get a 100%
efficient process.
• The energy “lost” in the process is energy
transformed into a form we cannot use.
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There’s No Such Thing as a Free
Ride
• When you drive your car, some of the
chemical potential energy stored in the
gasoline is released.
• Most of the energy released in the
combustion of gasoline is transformed into
sound or heat energy that adds energy to the
air rather than move your car down the
road.
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Units of Energy
• Calorie (cal) is the amount of energy needed to
raise one gram of water by 1 °C.
kcal = energy needed to raise 1000 g of water 1 °C.
food calories = kcals.
Energy Conversion Factors
1 calorie (cal)
1 Calorie (Cal)
=
=
4.184 joules (J)
1000 calories (cal)
1 kilowatt-hour (kWh)
=
3.60 x 106 joules (J)
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Energy Use
Unit
joule (J)
Energy
Energy Required
Required to
to Raise
Light 100-W
Temperature of 1 g Bulb for 1
of Water by 1°C
Hour
Energy
Used by
Average
U.S. Citizen
in 1 Day
4.18
3.6 x 105
9.0 x 108
calorie (cal)
1.00
8.60 x 104
2.2 x 108
Calorie (Cal)
1.00 x 10-3
86.0
2.2 x 105
kWh
1.1 x 10-6
0.100
2.50 x 102
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Example 3.5—Convert 225 Cal to Joules
1.
Write down the Given
quantity and its unit.
2.
Write down the quantity
Find:
you want to Find and unit.
Write down the appropriate Conversion
Conversion Factors.
Factors:
3.
4.
5.
6.
7.
Write a Solution Map.
Follow the solution map to
Solve the problem.
Given:
Solution
Map:
225 Cal
3 sig figs
?J
1 Cal = 1000 cal
1 cal = 4.184 J
Cal
cal
J
1000 cal
1 Cal
4.184 J
1 cal
Solution:
1000 cal 4.184 J
225 Cal 

 9.41 105 J
1 Cal
1 cal
Significant figures and
round.
Round:
Check.
Check:
225 Cal = 9.41 x 105 J
3 significant figures
Units and magnitude are
correct.
Chemical Potential Energy
• The amount of energy stored in a material is its
chemical potential energy.
• The stored energy arises mainly from the
attachments between atoms in the molecules and
the attractive forces between molecules.
• When materials undergo a physical change, the
attractions between molecules change as their
position changes, resulting in a change in the
amount of chemical potential energy.
• When materials undergo a chemical change, the
structures of the molecules change, resulting in a
change in the amount of chemical potential
energy.
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Energy Changes and
Chemical Reactions
• Chemical reactions happen most readily
when energy is released during the reaction.
• Molecules with lots of chemical potential
energy are less stable than ones with less
chemical potential energy.
• Energy will be released when the reactants
have more chemical potential energy than
the products.
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Exothermic Processes
• When a change results in the release of energy it is
called an exothermic process.
• An exothermic chemical reaction occurs when the
reactants have more chemical potential energy
than the products.
• The excess energy is released into the surrounding
materials, adding energy to them.
 Often the surrounding materials get hotter from the
energy released by the reaction.
Tro's "Introductory Chemistry",
Chapter 3
74
An Exothermic Reaction
Surroundings
reaction
Potential energy
Reactants
Tro's "Introductory Chemistry",
Chapter 3
Amount
of energy
released
Products
75
Endothermic Processes
• When a change requires the absorption of energy
it is called an endothermic process.
• An endothermic chemical reaction occurs when
the products have more chemical potential energy
than the reactants.
• The required energy is absorbed from the
surrounding materials, taking energy from them.
 Often the surrounding materials get colder due to the
energy being removed by the reaction.
Tro's "Introductory Chemistry",
Chapter 3
76
An Endothermic Reaction
Surroundings
reaction
Potential energy
Products
Tro's "Introductory Chemistry",
Chapter 3
Amount
of energy
absorbed
Reactants
77
Temperature Scales
• Fahrenheit scale, °F.
Used in the U.S.
• Celsius scale, °C.
Used in all other countries.
A Celsius degree is 1.8
times larger than a
Fahrenheit degree.
• Kelvin scale, K.
Absolute scale.
Tro's "Introductory Chemistry",
Chapter 3
78
Temperature Scales
100°C
373 K
212°F
671 R
Boiling
point water
25°C
298 K
75°F
534 R
Room temp
0°C
273 K
32°F
459 R
-38.9°C
234.1 K
-38°F
421 R
Melting
point ice
Boiling
point
mercury
-183°C
90 K
-297°F
162 R
-269°C
-273°C 4 K
0 K -452°F
-459 °F 7 R
Celsius
Kelvin
Fahrenheit
Boiling
point
oxygen
BP helium
0 R Absolute
Rankine zero
Temperature Scales
• The Fahrenheit temperature scale used as its
two reference points the freezing point of
concentrated saltwater (0 °F) and average
body temperature (96 °F).
More accurate measure now sets average body
temperature at 98.6 °F.
• Room temperature is about 72 °F.
Tro's "Introductory Chemistry",
Chapter 3
80
Temperature Scales, Continued
• The Celsius temperature scale used as its
two reference points the freezing point of
distilled water (0 °C) and boiling point of
distilled water (100 °C).
More reproducible standards.
Most commonly used in science.
• Room temperature is about 22 °C.
Tro's "Introductory Chemistry",
Chapter 3
81
Fahrenheit vs. Celsius
• A Celsius degree is 1.8 times larger than a
Fahrenheit degree.
• The standard used for 0° on the Fahrenheit
scale is a lower temperature than the
standard used for 0° on the Celsius scale.

F - 32 
C 
1.8
Tro's "Introductory Chemistry",
Chapter 3
82
The Kelvin Temperature Scale
• Both the Celsius and Fahrenheit scales have
negative numbers.
 Yet, real physical things are always positive amounts!
• The Kelvin scale is an absolute scale, meaning it
measures the actual temperature of an object.
• 0 K is called absolute zero. It is too cold for
matter to exist because all molecular motion
would stop.
 0 K = -273 °C = -459 °F.
 Absolute zero is a theoretical value obtained by
following patterns mathematically.
Tro's "Introductory Chemistry",
Chapter 3
83
Kelvin vs. Celsius
• The size of a “degree” on the Kelvin scale is the
same as on the Celsius scale.
 Although technically, we don’t call the divisions on the
Kelvin scale degrees; we call them kelvins!
 That makes 1 K 1.8 times larger than 1 °F.
• The 0 standard on the Kelvin scale is a much lower
temperature than on the Celsius scale.
• When converting between kelvins and °C, remember
that the kelvin temperature is always the larger
number and always positive!
K  C  273
Tro's "Introductory Chemistry",
Chapter 3
84
Example 3.7—Convert –25 °C to Kelvins
1.
Write down the Given
quantity and its unit.
2.
Write down the quantity
you want to Find and unit.
Write down the appropriate
Equations.
3.
4.
5.
6.
7.
Write a Solution Map.
Follow the solution map to
Solve the problem.
Given:
-25 °C
units place
Find:
Equation:
Solution
Map:
K
K = ° C + 273
°C
K
K   C  273
Solution:
K  ( 25  C)  273  258 K
Significant figures and
round.
Round:
Check.
Check:
258 K
units place
Units and magnitude are
correct.
Example 3.8—Convert 55° F to Celsius
1.
Write down the Given
quantity and its unit.
2.
Write down the quantity
you want to Find and unit.
Write down the appropriate
Equations.
3.
4.
5.
6.
7.
Write a Solution Map.
Follow the solution map to
Solve the problem.
Given:
55 °F
Find:
°C
Equation:
C 
Solution
Map:
F - 32 
1.8
°C
°F
C 
units place
and 2 sig figs
F - 32
1.8
Solution:
C 
55 F - 32 
Significant figures and
round.
Round:
Check.
Check:
1.8
 12.778 C
12.778 °C = 13 °C
units place and 2 sig figs
Units and magnitude are
correct.
Example 3.9—Convert 310 K to Fahrenheit
1.
Write down the Given
quantity and its unit.
2.
Write down the quantity
you want to Find and unit.
Write down the appropriate
Equations.
3.
4.
Write a Solution Map.
Given:
310 K
Find:
°F
Equation:
Solution
Map:
5.
Follow the solution map to
Solve the problem.
6.
Significant figures and
round.
Round:
Check.
Check:
7.
C 
K = °C + 273
K
units place
and 3 sig figs
F - 32
°C
1.8
°F
°C = K - 273 F  1.8C   32
Solution:
C  310  273  37 C
F  1.837  C   32  98 .6 F
98.6 °F = 99 °F
units place and 2 sig figs
Units and magnitude are
correct.
Practice—Convert 0 °F into Kelvin
Tro's "Introductory Chemistry",
Chapter 3
95
Practice—Convert 0 °F into Kelvin,
Continued
°C = 0.556(°F-32)
°C = 0.556(0-32)
°C = -18 °C
K = °C + 273
K = (-18) + 273
K = 255 K
Tro's "Introductory Chemistry",
Chapter 3
96
Energy and the Temperature of Matter
• The amount the temperature of an object
increases depends on the amount of heat
energy added (q).
If you double the added heat energy the
temperature will increase twice as much.
• The amount the temperature of an object
increases depending on its mass.
If you double the mass, it will take twice as
much heat energy to raise the temperature the
same amount.
Tro's "Introductory Chemistry",
Chapter 3
97
Heat Capacity
• Heat capacity is the amount of heat a substance
must absorb to raise its temperature by 1 °C.
cal/°C or J/°C.
Metals have low heat capacities; insulators have
high heat capacities.
• Specific heat = heat capacity of 1 gram of the
substance.
cal/g°C or J/g°C.
Water’s specific heat = 4.184 J/g°C for liquid.
Or 1.000 cal/g°C.
It is less for ice and steam.
98
Specific Heat Capacity
• Specific heat is the amount of energy required to raise
the temperature of one gram of a substance by 1 °C.
• The larger a material’s specific heat is, the more
energy it takes to raise its temperature a given amount.
• Like density, specific heat is a property of the type of
matter.
 It doesn’t matter how much material you have.
 It can be used to identify the type of matter.
• Water’s high specific heat is the reason it is such a
good cooling agent.
 It absorbs a lot of heat for a relatively small mass.
Tro's "Introductory Chemistry",
Chapter 3
99
Specific Heat Capacities
Substance
Specific Heat
J/g°C
Aluminum
0.903
Carbon (dia)
0.508
Carbon (gra)
0.708
Copper
0.385
Gold
0.128
Iron
0.449
Lead
0.128
Silver
0.235
Ethanol
2.42
Water (l)
4.184
Water (s)
2.03
Water (g)
2.02
Tro's "Introductory Chemistry",
Chapter 3
100
Heat Gain or Loss by an Object
• The amount of heat energy gained or lost by an
object depends on 3 factors: how much material
there is, what the material is, and how much the
temperature changed.
Amount of Heat = Mass x Heat Capacity x Temperature Change
q = m x C x DT
Tro's "Introductory Chemistry",
Chapter 3
101
Example 3.10—Calculate Amount of Heat Needed to
Raise Temperature of 2.5 g Ga from 25.0 to 29.9 °C
1.
Write down the Given
quantity and its unit.
2.
Write down the quantity
you want to Find and unit.
Write down the appropriate Equation:
Equations.
3.
4.
5.
Write a Solution Map.
Follow the solution map to
Solve the problem.
Given:
m = 2.5 g, T1 = 25.0 °C,
T2= 29.9 °C, C = 0.372 J/g°C
Find:
q, J
q  m  C  DT
m, C, DT
Solution
Map:
Solution:
q
q  m  C  DT


q  2.5 g  0.372 g JC  29.9 - 25.0C 
q  4.557 J
6.
7.
Significant figures and
round.
Round:
Check.
Check:
4.557 J = 4.6 J
2 significant figures
Units and magnitude are
correct.
Practice—Calculate the Amount of Heat Released
When 7.40 g of Water Cools from 49° to 29 °C
Tro's "Introductory Chemistry",
Chapter 3
110
Practice—Calculate the Amount of Heat Released
When 7.40 g of Water Cools from 49° to 29 °C,
Continued
•
•
Sort
Information
Strategize
Given:
T1 = 49 °C, T2 = 29 °C, m = 7.40 g
Find:
q, J
Solution Map:
Cs m, DT
q
q  m  C s  ΔT
Relationships:
•
•
q = m ∙ Cs ∙ DT
Cs = 4.18 J/gC (Table 3.4)
Solution:
Follow the
DT  T2  T1
concept
plan to
DT  29 C - 49C 
solve the
 - 20 C
problem.
Check:
Check.
q  m  C s  ΔT


 7.40 g   4.18 g C  - 20 C 
J
 618.64 J  6.2  10 2 J
The unit and sign are correct.