### Chapter 3 Lecture

Chapter 3
Matter and Energy
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CHAPTER OUTLINE
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Energy & Heat
Temperature Scales
Specific Heat
Classification of Matter
Physical & Chemical Properties
Physical & Chemical Changes
Conservation of Mass
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ENERGY & HEAT
 Energy is defined as the capacity of matter to
do work.
 Work is defined as the result of a force acting
on a distance.
 There are two types of energy:
Potential
(stored)
Kinetic
(moving)
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ENERGY & HEAT
 Energy possesses many forms
(chemical, electrical, thermal,
etc.), and can be converted
from one form into another.
 In chemistry, energy is
commonly expressed as heat.
PE is
converted to
KE
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ENERGY & HEAT
 Energy is conserved.
 The law of conservation of energy
states that energy is neither created
nor destroyed.
 The total amount of energy is
constant.
 Energy can be changed from one form
to another.
 Energy can be transferred from one
object to another.
UNITS OF ENERGY
 The SI unit of energy is the joule (J), named
after the English scientist James Joule (1818–
1889).
HEAT vs.
TEMPERATURE
 Heat is measured in SI units of joule or the
common unit of calorie.
Heat &
temperature
are NOT the
same thing!
1 cal = 4.184 J
Although the same amount of heat is added to
both containers, the temperature increases more in
the container with the smaller amount of water.
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HEAT vs.
TEMPERATURE
The difference between Heat and Temperature
AHeat
formis of
theenergy
total
associated
energy ofwith
all
particles of matter
Temperature
A measure ofisthe
the
intensity
averageofkinetic
heat or
energy
how hot
of particles
or cold aof
substance
matter is
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TEMPERATURE
 Temperature is a measure of how hot or cold a
substance is.
 Thermometer is an instrument used for
measuring temperature, and is based on
thermometric properties of matter (i.e.
expansion of solids or liquids).
 Three scales are used for measuring
temperature.
TEMPERATURE
SCALES
Fahrenheit
32 - 212
Celsius
0 - 100
Kelvin
273 - 373
TEMPERATURE
SCALES
 To convert from one scale to another the
following relationships can be used:
K = C + 273
F = (1.8 x C) + 32
C = (F - 32) ÷ 1.8
 Alternately,
F = [(C + 40) x 1.8]-40
C = [(F + 40) ÷ 1.8]-40
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Example 1:
The melting point of silver is 960.8 C. What is this
temperature in Kelvin?
TK = TC + 273
TK = 960.8 + 273 = 1234
1233.8
KK
Example 2:
Pure iron melts at 1800 K. What is this temperature
in Celsius?
TC = TK - 273
TC = 1800 - 273 = 1527 C
Example 3:
On a winter day, the temperature is 5 F. What is
this temperature on the Celsius scale?
TC = [(5 +40) ÷ 1.8]- 40 = -15 C
Example 4:
To make ice cream, rock salt is added to crushed
ice to reach temperature of -11 C. What is this
temperature in Fahrenheit?
TF = [(-11 + 40) x 1.8]- 40 = 12 F
SPECIFIC
HEAT
 Different
The specific
materials
heat of have
a substance
different
is capacities
the
for storing
amount
of heat
heat.required to change the
temperature of 1 g of that substance by 1C.
 Units of specific heat are:
s = J / g ºC
s = cal / g ºC
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SPECIFIC
HEAT
Specific Heat of Some Substances
Substance
Most substances
Aluminum
have substantially
Copper
lower specific
heats compared
Iron
to water
(cal/gC)
(J/gC)
0.214
0.897
0.0920
0.385
0.0308
0.129
Ammonia
0.488
2.04
Ethanol
0.588
2.46
Water
1.00
4.184
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SPECIFIC
HEAT
 When heated, substances with low specific
heat get hot faster, while substances with high
specific heat get hot at a slower rate.
 When cooled, substances with low specific heat
cool faster, while substances with high specific
heat cool at a slower rate.
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CALCULATING
HEAT
 The amount of heat lost or gained by a
substance is related to three quantities:
Heat
Q
=
Mass of
substance
=
m
x
Specific heat
of substance
x
x
s
x
Change in its
temperature
T
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Example 1:
How much heat is needed to raise the temperature
of 200. g of water by 10.0 C. (Specific heat of
water is 4.184 J/gC)
m = 200. g
s = 4.184 J/gC
T = 10.0 C
Q = ???
Q= m x s x
T
Q = (200. g)(4.184 J/gºC)(10.0 ºC)
Q = 8370 J or 8.37 kJ
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Example 2:
Ethanol has a specific heat of 2.46 J/gC. When 655 J
are added to a sample of ethanol, its temperature rises
from 18.2 C to 32.8 C. What is the mass in grams of
the ethanol sample?
Q = 655 J
s = 2.46 J/gC
T = 14.6 C
m =
Q
=
s x DT
655 J
o
o
(2.46 J/g C )(14.6 C )
m = 18.2 g
m = ???
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ENERGY & NUTRITION
 In the laboratory, foods are burned in a
calorimeter to determine their energy. A sample
of food is burned in the calorimeter, and the
energy released is absorbed by water
surrounding the calorimeter.
 The energy of the food can be calculated from
the mass of the food and the temperature
increase of the water.
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Example 3:
A 2.3-g sample of butter is placed in a calorimeter
containing 1900 g of water at a temperature of 17 C.
After the complete combustion of the butter, the water
has a temperature of 28 C. What is the energy value of
butter in Cal/g?
1. Calculate heat absorbed by water
Heat absorbed
by water
=
Heat released
by butter
2. Calculate energy value of butter
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Example 3:
1. Calculate heat absorbed by water
m = 1900 g
s = 1.00 cal/gC
T = 11 C
Q = ???
Q= m x s x
T
Q = (1900 g)(1.00 cal/gºC)(11 ºC)
Q = 21000 cal = 21 Cal
2. Calculate energy value of butter
21 Cal
= 9.1 Cal/g
2.3 g
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ENERGY IN
CHEMICAL CHANGES
 In
Higher
all chemical
energy systems
changes,are
matter
less stable
eitherthan
absorbs
or releases
lower
energy
energy.
systems.
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ENERGY IN
CHEMICAL CHANGES
 When energy is released during a chemical
change, it is said to be exothermic.
Exothermic reactions heat up
 When energy is gained during a chemical
change, it is said to be endothermic.
Endothermic reactions cool down
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EXOTHERMIC vs.
ENDOTHERMIC
higher
energy
potential
is absorbed
energy
lower
energy
potential
is given
energy
off
Which is exothermic and
which is endothermic?
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4.4
CLASSIFICATION
OF MATTER

 Matter
Matter is
can
anything
be classified
that has
by its
mass,
physical
and occupies
state as
space.
solid, liquid or gas.
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SOLIDS
 Solids are densely packed particles
with definite shape and volume.
 Solid particles have strong
forces of attraction towards
each other.
 Solids are not very
compressible.
 Ice, diamond, quartz, and
iron are examples of solid
matter.
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LIQUIDS
 Liquids are loosely packed particles with
definite volume but indefinite shape.
 Liquid particles have moderate
forces of attraction towards each
other and are mobile.
 Liquids are slightly compressible.
 Water, gasoline, alcohol, and
mercury are all examples of
liquid matter.
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GASES
 Gases are very loosely packed particles
with indefinite shape or volume.
 Gas particles have little or no
forces of attraction towards each
other.
 Gases are very compressible.
 Oxygen, helium, and carbon
dioxide are examples of gases.
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GASES ARE
COMPRESSIBLE
Since the atoms or
molecules that compose
gases are not in contact
with one another, gases
can be compressed.
SUMMARY OF
PROPERTIES OF MATTER
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CLASSIFICATION
OF MATTER
MATTER
Anything that has mass
PURE SUBSTANCE
MIXTURE
Fixed composition &
properties
Variable composition &
properties
Mixtures can be converted into pure substances by
simple physical processes (e.g. filtration, evaporation)
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MIXTURES
MIXTURE
Variable composition &
properties
HOMOGENEOUS
HETEROGENEOUS
Uniform composition
& properties
Non-uniform composition
& properties
Tea, Coke
Ink
Also called
solutions
Cement
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PURE SUBSTANCES
PURE SUBSTANCE
Fixed composition &
properties
ELEMENTS
COMPOUNDS
Composed of one type
of atom
Propertiesparticle
are unique
Smallest
is a
chemically
combined
compared to their
molecule
2 or more elements
components
hydrogen,
Compounds can be converted into elements
water, salt
by
copper,
gold
chemical
processes or reactions (e.g. electrolysis)
aspirin
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PURE SUBSTANCES
separation of
compound through
chemical methods
(electrolysis)
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CONCEPT
CHECK
Classify each substance below as element, compound or
mixture.
Mixture:composition
of two
or
more
Element:
only
one
of atom
Compound:
is
fixed
Element: only one typetype
of atom
types of substances
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MIXTURES
 Mixtures are 2 or more substances physically
combined together.
 Mixtures possess properties similar to those of
their components.
 Mixtures can be separated easily by a physical
process.
 Two types of mixtures are possible:
heterogeneous
homogeneous
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HETEROGENEOUS
MIXTURES
 Heterogeneous mixtures are non-uniform in
their composition.
Heterogeneous
 Examples include vegetable soup, cement and
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HOMOGENEOUS
MIXTURES

 Homogeneous
Homogeneous mixtures
mixtures are
are uniform
called solutions.
in their
composition.
Homogeneous
 Examples include gasoline, soda pop and salt
solution.
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MIXTURES vs.
COMPOUNDS
List 3 differences between compounds & mixtures.
Composition
Compounds have fixed composition while
mixtures have varied composition
Properties
Compounds have unique properties while
mixtures have blended properties
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MIXTURES vs.
COMPOUNDS
List 3 differences between compounds & mixtures.
Make-up
Compounds are chemically combined
(cannot be easily separated) while mixtures
are physically combined (easily separated)
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PHYSICAL & CHEMICAL
PROPERTIES
 The characteristics of a substance are called its
properties.
 Physical properties are those that describe the
matter without changing its composition.
Examples are density, color, melting and
boiling points, and electrical conductivity.
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PHYSICAL & CHEMICAL
PROPERTIES
 The characteristics of a substance are called its
properties.
 Chemical properties are those that describe
how matter behave in combination with other
matter, and involve change in its composition.
Examples are flammability, corrosion, and
reactivity with acids.
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Examples:
Identify each of the following properties as
physical or chemical:
Physical
1. Oxygen is a gas
Chemical
2. Helium is un-reactive
Physical
3. Water has high specific heat
4. Gasoline is flammable
Chemical
5. Sodium is soft & shiny
Physical
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PHYSICAL
CHANGES
 Changes in physical properties of matter that
do not involve change in its composition are
called physical changes.
Examples are melting,
evaporation and other
phase changes.
 Physical changes are
easily reversible.
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CHEMICAL
CHANGES
 A
Chemical
change that
changes
alters
are
the
not
chemical
easily reversible,
composition
of matter,
and
are commonly
and forms
called
new chemical
substance is called a
chemical change.
reactions.
Examples are burning, rusting, and reaction
with acids.
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Examples:
Identify each of the following changes as
physical or chemical:
Chemical
1. Cooking food
Physical
2. Mixing sugar in tea
Physical
3. Carving wood
Chemical
4. Burning gas
5. Food molding
Chemical
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CONSERVATION
OF MASS
 Similar to the law of conservation of energy,
the law of the conservation of mass states that
matter is neither created nor destroyed.
 The total mass of substances does not change
during a chemical reaction.
Mass of
Reactants
=
Mass of
Products
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CONSERVATION
OF MASS
Suppose that we burn 58 g of butane in a lighter. It
will react with 208 g of oxygen to form ??? g of
carbon dioxide and 90 g of water.
266 g reactant
mass reactants
→
266 g product
=
mass products
The number of substances and their properties may
change, but the total amount of matter remains constant.
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Burning: a Chemical Change
 The butane molecules
react with oxygen
molecules in air to
form new molecules,
carbon dioxide and
water.
 This is a chemical
change.
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THE END
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