Chapter 3
Matter and Energy
1
CHAPTER OUTLINE







Energy & Heat
Temperature Scales
Specific Heat
Classification of Matter
Physical & Chemical Properties
Physical & Chemical Changes
Conservation of Mass
2
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)
3
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
4
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.
7
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
8
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
11
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
16
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
17
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.
18
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
19
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
20
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 = ???
21
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.
22
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
23
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
24
ENERGY IN
CHEMICAL CHANGES
 In
Higher
all chemical
energy systems
changes,are
matter
less stable
eitherthan
absorbs
or releases
lower
energy
energy.
systems.
25
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
26
EXOTHERMIC vs.
ENDOTHERMIC
higher
energy
potential
is absorbed
energy
lower
energy
potential
is given
energy
off
Which is exothermic and
which is endothermic?
27
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.
28
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.
29
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.
30
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.
31
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
33
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)
34
MIXTURES
MIXTURE
Variable composition &
properties
HOMOGENEOUS
HETEROGENEOUS
Uniform composition
& properties
Non-uniform composition
& properties
Tea, Coke
Ink
Also called
solutions
Salad dressing
Cement
35
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
36
PURE SUBSTANCES
separation of
compound through
chemical methods
(electrolysis)
37
CONCEPT
CHECK
Classify each substance below as element, compound or
mixture.
Mixture:composition
made
of two
or
more
Element:
only
one
of atom
Compound:
is
fixed
Element: only one typetype
of atom
types of substances
38
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
39
HETEROGENEOUS
MIXTURES
 Heterogeneous mixtures are non-uniform in
their composition.
Heterogeneous
 Examples include vegetable soup, cement and
salad dressing.
40
HOMOGENEOUS
MIXTURES

 Homogeneous
Homogeneous mixtures
mixtures are
are uniform
called solutions.
in their
composition.
Homogeneous
 Examples include gasoline, soda pop and salt
solution.
41
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
42
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)
43
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.
44
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.
45
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
46
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.
47
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.
48
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
49
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
50
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.
51
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.
52
THE END
53