Transcript Chapter 3

Chapter 3
Matter and
Energy
Assigned Problems
Recommended
 Exercises: 1-27 (odd)
Required
 Problems: 29-85 (odd)
 Cumulative Problems: 87-105 (odd)
Optional
 Highlight Problems: 107-113 (odd)
What Is Matter?
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Matter is any material that has mass and
occupies space
Matter is made up of small particles
 Atoms
 Molecules
Includes all things (living and nonliving) such
as plants, soil, and rocks and any material we
use such as water, wood, clothing, etc.
Classifications (of a sample of matter) is based
on whether its shape and volume are definite
or indefinite
Classifying Matter According to Its State
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Solid
 Has a rigid, definite shape and definite volume
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Crystalline solids have a regular, internal long-range
order of atoms, ions, or molecules
Amorphous solids have no long-range order of atoms,
ions, or molecules in their lattice structure
Liquid
 Has an indefinite shape and a definite volume.
 It will take the shape of the container it fills
Gas
 Has an indefinite shape and an indefinite volume.
 It will take the shape and completely fill the volume
of the container it fills
 Gases are compressible
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Water is one of the few substances commonly found
in all three physical states
(a) Solid (Ice)
Fig3_2
(b) Liquid (Water)
(c) Gas (Steam)
Classifying Matter by Its Composition
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Matter can also be classified in terms of its chemical
composition
Matter
Pure Substance
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Mixture
Pure Substances: Composed of only one atom or molecule
Mixtures: Composed of two or more different atoms or
molecules combined in various proportions
Pure Substances
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Matter that has a definite and constant
composition is a pure substance
Composed of the same substance; no
variation
6 million pure substances have been isolated:
112 are elements, the rest are compounds
The two classifications of pure substances:
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Elements: e.g., a pure sample of copper or a pure
sample of gold (one type of atom)
Compounds: e.g., example, a pure sample of
water or a pure sample of sucrose (one type of
molecule)
Pure Substances
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Elements
 Substances which can not be broken down into
simpler substances by chemical reactions
 Fundamental substances
Compounds
 Two or more elements combined chemically in a
definite and constant ratio
 Can be broken down into simpler substances
 Most of matter is in the compound form
Pure Substances
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Compounds
 Results from a chemical combination of two or more
elements
 Can be broken down into elements by chemical processes
 Properties of the compound not related to the properties of
the elements that compose it
 Water is composed of hydrogen and oxygen gases
(combined in a 2:1 ratio)
Mixtures
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Something of variable composition
Result from the physical combination of two
or more substances (elements or
compounds)
Made up of two or more types of
substances physically mixed
Each substance retains its identity because
the substances are not chemically mixed
Mixtures of the same components can vary
in composition
Mixtures
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Mixtures can be classified by the (visual)
uniformity of the mixture’s components
Homogeneous mixture:
 Same uniform composition throughout
 Not possible to see the two substances
present
Heterogeneous mixture:
 Composition is not uniform throughout the
sample.
 It contains visibly different parts or phases
Mixtures
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Homogenous mixtures
 A sugar solution
 14 karat gold, a mixture of copper and gold
 Air, a mixture of gases (oxygen, nitrogen)
Heterogeneous mixture
 Oil and vinegar
 Raisin cookies
 Sand
Pure substance
 e.g. copper (all elements are pure substances)
Compounds vs. Mixtures
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Compounds are not mixtures
 Cannot be separated by a physical process
 Can be subdivided by a chemical process into two
or more simpler substances
 Simpler substances have different properties from
the compound
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Mixtures
 Unlike compounds, mixtures can be separated by
a physical process
 Each substance in a mixture retain its own
individual properties
Classification of Matter
Matte r
Pur e Subs tance s
Ele m e nts
Com pounds
Chemical Methods
M ixtur e
Hom oge ne ous
Mixtur e
He te r oge ne ous
Mixtur e
Physical Methods
Pur e Subs tance s
Physical and Chemical Properties
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Various kinds of matter are differentiated by
their properties
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Properties are the characteristics of a substance
used to identify and describe it
Two general categories:
 Physical Properties
 Chemical Properties
Properties can be:
 Directly observable (physical)
 The interaction of the matter with other
substances (chemical)
Physical and Chemical Properties:
Physical Properties
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A physical property is a characteristic of
a substance that can be observed
without changing a substance into
another substance
Characteristics of matter that can be directly
observed or measured without changing its
identity or composition
 Color, odor, physical state, density, melting
point, boiling point
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Physical and Chemical Properties:
Chemical Properties
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A chemical property describes the way a substance
undergoes a change or resists change to form a
new substance
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Properties that matter exhibits as it undergoes
changes in chemical composition:
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Objects made from copper will turn green when
exposed to moist air for long periods
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Gold objects will resist change when exposed to
moist air for long periods
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Sodium metal will react strongly with water and
produce hydrogen gas.
Classifying Properties
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The boiling point of ethyl alcohol is 78 °C
 Physical property – describes an inherent
characteristic of alcohol, its boiling point
Diamond is very hard
 Physical property – describes inherent
characteristic of diamond – hardness
Sugar ferments to form ethyl alcohol
 Chemical property – describes behavior
of sugar, ability to form a new substance
(ethyl alcohol)
Physical and Chemical Changes
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Changes in matter are regular occurrences:
 Food is cooked
 Paper is burned
 Iron rusts
Matter undergoes changes as a result of the
application of energy
Changes in matter are also categorized as two types:
 Physical
 Chemical
Physical and Chemical Changes
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A physical change is a process that alters the
appearance of a substance but does not
change its chemical identity or composition
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Folding aluminum foil sheets
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Crushing ice cubes
No new substance is formed
 Most common is a change of a substance’s
physical state
 The freezing of liquid water
 Evaporation of liquid water to steam
Physical and Chemical Changes
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A chemical change is a process that changes
the chemical composition of a substance
 Also called a chemical reaction
 (At least) one new substance is produced
 Wood burning, iron rusting, alka-seltzer
tablet reacting with water
 During a chemical change, the original
substance is converted into one or more
new substances with different chemical and
physical properties
Classifying Changes
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Melting of snow
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Burning of gasoline
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Physical change – a change of state but
not a change in composition
Chemical change – combines with
oxygen to form new compounds
Rusting of iron
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Chemical change – combines with
oxygen to form a new reddish-colored
substance (ferric oxide)
Classifying Changes
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Iron metal is melted
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Iron combines with oxygen to form rust
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Physical change – describes a state
change, but the material is still iron
Chemical change – describes how iron
and oxygen combine to make a new
substance, rust (ferric oxide)
Sugar ferments to form ethyl alcohol
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Chemical change – describes how sugar
forms a new substance (ethyl alcohol)
Conservation of Mass
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During a physical change: No new substance is
formed
During a chemical change: At least one new
substance is formed
Whether it is a physical or chemical change, the
amount of matter remains the same
The law of conservation of mass states that the total
mass of materials present after a chemical reaction is
the same as the total mass before the reaction
Matter is never created or destroyed
Energy
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Two major components of the universe:
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Matter
Energy
Energy is the capacity to do work or produce heat
 Electrical, radiant, mechanical, thermal, chemical,
nuclear
 Nearly all changes that matter undergoes involves
the release or absorption of energy
Chemistry is the study of matter
 The properties of different types of matter
 The way matter behaves when influenced by other
matter and/or energy
Energy
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Energy is the part of the universe that has the
ability to do work
Energy can be converted from one form to another
but it is neither created nor destroyed (the law of
conservation of energy)
Energy has two classifications
 Potential: Stored energy
 Kinetic: Motion energy
All physical changes and chemical changes
involve energy changes
Energy
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Potential energy:
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Determined by an objects position (or composition)
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Chemical energy (also potential energy) is stored in the bonds
contained within a molecule. It is released in a chemical reaction
Kinetic energy
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Energy that matter acquires due to motion
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Converted from the potential energy
All forms of energy can be quantified in the same units
Units of Energy
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The joule (J) is the SI unit of heat energy
The calorie (cal) is an older unit used for
measuring heat energy (not an SI unit)
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The amount of energy needed to raise the
temperature of one gram of water by 1°C
4.184 J = 1 cal
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1 kcal = 1000 cal
The Cal is the unit of heat energy in
nutrition
1 Cal = 1000 cal = 1 kcal
Energy: Chemical and Physical Change
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All physical changes and chemical changes involve
energy changes which convert energy from one form to
another
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In terms of a chemical reaction the universe is divided
into two parts:
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The system (chemical reaction)
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The surroundings (everything else)
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The potential energy differences between the reactants and
products determine whether heat flows into or out of a chemical
system
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Whether a reaction is exothermic or endothermic depends on how the
potential energy of the products compares to the PE of the reactants
Energy: Chemical and Physical Change
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Chemical systems with high
potential energy tend to change
in order to lower their potential
energy by the release of heat
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Chemical reactions that release
heat are called exothermic
Chemical systems with low
potential energy tend to change
in order to increase their
potential energy by the
absorption of heat
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Chemical reactions that absorb
heat are called endothermic
Temperature
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Temperature is a number related to the
average kinetic energy of the molecules of a
substance
In a substance, the temperature:
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measures the hotness or coldness of an object
measures the average molecular motions in a
system
relates (directly) to the kinetic energy of the
molecules
Temperature
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Fahrenheit Scale, °F
 Used in USA
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Water’s freezing point = 32°F, boiling point = 212°F
Celsius Scale, °C
 Used in science (USA) and everyday use in
most of the world
 Temperature unit larger than the Fahrenheit
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Water’s freezing point = 0°C, boiling point = 100°C
Temperature
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Kelvin Scale, K
SI Unit
 Used in science
 Temperature unit same size as Celsius
 Water’s freezing point = 273 K (0 ºC),
boiling point = 373 K (100 ºC)
 Absolute zero is the lowest temperature
theoretically possible
 No negative temperatures
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Converting °C to °F
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Units are different sizes
Fahrenheit scale: 180 degree intervals
between freezing and boiling
 Celsius scale: 100 degree intervals
between freezing and boiling
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180 F 9 F 1.8 F
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100 C 5 C 1 C
212ºF
100ºC
180
Fahrenheit
degrees
100
Celsius
degrees
32ºF
1.8 F
1 C
Fig2_9
Boiling point
0ºC
Freezing point
1.8 F
1 C
Converting °C to °F
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To convert from °C to °F
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Different values for the freezing points
32 °F
0 °C
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add 32 to the °F value
Different size of the degree intervals in each
scale
TF
1.81.8FF
T 
32
TTC

C 32
1 C
1 C
F
Converting °C to K
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Temperature units are the same size
Differ only in the value assigned to
their reference points
0 °C =
K273
= K°C
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add 273 to the °C value
+ 273
25°C is room temperature, what is the
equivalent temperature on the Kelvin
scale?
25ºC
ºC+++273
273===298
298K
25
ºC
273
298
KK
25
Example
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A cake is baked at 350 °F. What is
this in Centigrade/Celsius? In Kelvin?
TF 
1.8 F
TC  32
1 C
1 C
TF  32   TC 
1.8 F
1 C
350318
F°F
 32   TC 
1.8 F
TC  176.6667 C
176.7  273 =
449.7
450 K
K
Temperature Changes:
Heat Capacity
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Heat is the total amount of energy in a system
 It is function of the amount of motion (kinetic
energy) contained in molecules
 It is also a function of the potential energy of the
molecules
 It involves the exchange of thermal energy caused
by a temperature difference
Heat vs. Temperature
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Within a quantity of matter:
Heat has units of Joules and temperature has units in
degrees
Temperature relates only to kinetic energy within a
molecule
Heat is the total amount of energy in a molecule: It
contains a kinetic and potential energy component
Heat energy can be added or removed without a
change in temperature
As heat energy increases the temperature increases
Temperature Changes: Heat Capacity
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Heat energy is the form of energy most often
released or required for chemical and physical
changes
Every substance must absorb a different amount
of heat to reach a certain temperature
Different substances respond differently when
heat is applied
The amount of heat required to raise the
temperature of a given quantity of a substance by
1 ºC is called its heat capacity
Temperature Changes: Specific Heat
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If 4.184 J of heat is applied to:
 1 g of water, its temperature is raised by 1 °C
 1 g of gold, its temperature is raised by 32 °C
 Some substances requires large amounts of heat
to change their temperatures, and others require a
small amount
 The precise amount of heat that is required to
cause a given amount of substance (in grams) to
have a rise in temperature is called a substance’s
“specific heat”
Specific Heat
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The amount of heat energy (q) needed to
raise 1 gram of a substance by 1 °C
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Specific to the substance
The higher the specific heat value, the less its
temperature will change when it absorbs heat
SH values given in table 3.4, page 71
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Only for heating/cooling not for changes in state
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heat (qheat
)
J J (or cal)cal
C  SH 

 or
gramgram
s ΔT s Δtg  C g  Cg  C
Specific Heat Expression with
Calories and Joules
1 cal is the energy needed to
heat 1 g of water 1 °C
 1 cal is 4.184 J
 Make a conversion factor
from the statements
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1 cal
4.184 J
C


w ater 1g  1 C 1g  1 C
Specific Heat Equation
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The rearrangement of the SH equation gives the
expression called the “heat equation”
C
heat (q)
mass ( g )  ΔT
heat ( q )  C  mass ( g )  ΔT
C
 J 
  m(g) ΔT(C)
q(J)  
 g C 
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q = heat
C = specific heat (different for each substance)
m = mass (g)
∆T = change in temperature (°C)
 answer in joules
Specific Heat Equation

Energy (heat) required to change
the temperature of a substance
depends on:
The amount of substance being
heated (g)
 The temperature change (initial T
and final T in °C)
 The identity of the substance
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Energy and T
Heat (q)C  mass (g) Δt
2×
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2×
The amount the temperature of an object
increases depends on the amount of heat added
(q)
 If you double the added heat energy (q), the
temperature will increase twice as much.
 When a substance absorbs energy, q is
positive, temperature increases
 When a substance loses energy, q is
negative, temperature decreases
Energy and Heat Capacity Calculations
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Use same problem solving steps as before (Chapter 2)
 State the given and needed units
 Write the unit plan to convert the given unit to the final unit
 State the equalities and the conversion factors
 Set up the problem to cancel the units
Pepsi One™ contains 1 Calorie per can.
How many joules is this?
1 Cal = 1000 cal
4.184 J = 1 cal
1 Cal 1000 cal 4.184 J
 4184 J
1 Cal 1 cal
Energy and Heat Capacity Calculations

The 4184 J from the Pepsi One™
will heat how many grams of water
from 0°C to boiling?
q
m
C  T
4184 J 1C1g
4.184 J 100 C
10 g  10 mL
Energy and Heat Capacity Calculations

How many grams of water would reach
boiling if the water started out at room
temperature (25°C)?
m
q
C  T
4184 J 1 C  1 g
  25
C
4.184 J 10075

13.33 g  13.33 mL
Energy and Heat Capacity Calculations

If 50.0 J of heat is applied to 10.0 g
of iron, by how much will the
temperature of the iron increase?
50.0 J g  C
 11.11 C
0.45 J 10.0 g
Solve for ΔT
q  C  m  ΔT
T
q

Cm

end