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Chapter Five:
The heat death of the universe seems to be a logical consequence of the second law of thermodynamics, but scientists are not
certain if the second law should apply to the whole universe. What
do you think? Will the universe with all its complexities of organization end with the simplicity of spread-out and slowly vibrating
molecules? As has been said, nature is full of symmetry—so why
should the universe begin with a bang and end with a whisper?
SUMMARY
The kinetic theory of matter assumes that all matter is made up of tiny,
ultimate particles of matter called molecules. A molecule is defined as
the smallest particle of a compound, or a gaseous element, that can exist
and still retain the characteristic properties of that substance. Molecules interact, attracting each other through a force of cohesion. Liquids, solids, and gases are the phases of matter that are explained by the
molecular arrangements and forces of attraction between their molecules. A solid has a definite shape and volume because it has molecules
that vibrate in a fixed equilibrium position with strong cohesive forces.
A liquid has molecules that have cohesive forces strong enough to give
it a definite volume but not strong enough to give it a definite shape.
The molecules of a liquid can flow, rolling over each other. A gas is composed of molecules that are far apart, with weak cohesive forces. Gas
molecules move freely in a constant, random motion.
Molecules can have vibrational, rotational, or translational kinetic
energy. The temperature of an object is related to the average kinetic
energy of the molecules making up the object. A measure of temperature tells how hot or cold an object is on two arbitrary scales, the
Fahrenheit scale and the Celsius scale. The absolute scale, or Kelvin scale,
has the coldest temperature possible (273°C) as zero (0 K).
The observable potential and kinetic energy of an object is the
external energy of that object, while the potential and kinetic energy of
the molecules making up the object is the internal energy of the object.
Heat refers to the total internal energy and is a transfer of energy that
takes place (1) because of a temperature difference between two objects,
or (2) because of an energy-form conversion. An energy-form conversion is actually an energy conversion involving work at the molecular
level, so all energy transfers involve heating and working.
A quantity of heat can be measured in joules (a unit of work or
energy) or calories (a unit of heat). A kilocalorie is 1,000 calories,
another unit of heat. A Btu, or British thermal unit, is the English system unit of heat. The mechanical equivalent of heat is 4,184 J 1 kcal.
The specific heat of a substance is the amount of energy (or heat)
needed to increase the temperature of 1 gram of a substance 1 degree
Celsius. The specific heat of various substances is not the same because
the molecular structure of each substance is different.
Energy transfer that takes place because of a temperature difference does so through conduction, convection, or radiation. Conduction
is the transfer of increased kinetic energy from molecule to molecule.
Substances vary in their ability to conduct heat, and those that are poor
conductors are called insulators. Gases, such as air, are good insulators.
The best insulator is a vacuum. Convection is the transfer of heat by the
displacement of large groups of molecules with higher kinetic energy.
Convection takes place in fluids, and the fluid movement that takes
place because of density differences is called a convection current. Radiation is radiant energy that moves through space. All objects with an
absolute temperature above zero give off radiant energy, but all objects
absorb it as well. Energy is transferred from a hot object to a cold one
through radiation.
Heat and Temperature
115
The transition from one phase of matter to another is called a phase
change. A phase change always absorbs or releases a quantity of latent
heat not associated with a temperature change. Latent heat is energy that
goes into or comes out of internal potential energy. The latent heat of
fusion is absorbed or released at a solid-liquid phase change. The latent
heat of fusion for water is 80.0 cal/g (144.0 Btu/lb). The latent heat of
vaporization is absorbed or released at a liquid-gas phase change. The
latent heat of vaporization for water is 540.0 cal/g (970.0 Btu/lb).
Molecules of liquids sometimes have a high enough velocity to
escape the surface through the process called evaporation. Evaporation
is a cooling process, since the escaping molecules remove the latent heat
of vaporization in addition to their high molecular energy. Vapor molecules return to the liquid state through the process called condensation.
Condensation is the opposite of evaporation and is a warming process.
When the condensation rate equals the evaporation rate, the air is said
to be saturated. The rate of evaporation can be increased by
(1) increased temperature, (2) increased surface area, (3) removal of
evaporated molecules, and (4) reduced atmospheric pressure.
Warm air can hold more water vapor than cold air, and the ratio
of how much water vapor is in the air to how much could be in the air
at that temperature (saturation) is called relative humidity.
Thermodynamics is the study of heat and its relationship to
mechanical energy, and the laws of thermodynamics describe these relationships: The first law of thermodynamics states that the energy supplied to a thermodynamic system in the form of heat, minus the work
done by the system, is equal to the change in internal energy. The second law of thermodynamics states that heat flows from objects with a
higher temperature to objects with a cooler temperature. The second
law implies a degradation of energy as high-quality (more ordered)
energy sources undergo degradation to low-quality (less ordered)
sources. Entropy is a thermodynamic measure of disorder, and entropy
is seen as continually increasing in the universe and may result in the
maximum disorder called the heat death of the universe.
Summary of Equations
5.1
TF 9
T 32º
5 C
TC 5
T 32º 9 F
5.2
5.3
TK TC 273
5.4
Quantity of heat (mass)(specific heat)(temperature change)
Q mcT
5.5
Heat absorbed or released (mass)(latent heat of fusion)
5.6
Q mLf
Heat absorbed or released (mass)(latent heat of vaporization)
Q mLv
116
Chapter Five:
Heat and Temperature
5.7
relative humidity water vapor in air
100%
capacity at present temperature
R.H. gm3 present
100%
gm3 max
5.8
(heat) (work) internal energy difference between two states
Q W U 2 U1
5.9
work (mechanical equivalence of heat)(difference in heat
input and heat output)
W J(QH QL)
KEY TERMS
absolute scale (p. 95)
absolute zero (p. 95)
British thermal unit (p. 98)
calorie (p. 98)
Celsius scale (p. 95)
conduction (p. 100)
convection (p. 101)
entropy (p. 113)
external energy (p. 97)
Fahrenheit scale (p. 94)
first law of
thermodynamics (p. 112)
heat (p. 96)
heat engine (p. 109)
internal energy (p. 97)
Kelvin scale (p. 95)
kilocalorie (p. 98)
kinetic molecular
theory (p. 90)
latent heat of fusion (p. 106)
latent heat of
vaporization (p. 106)
molecule (p. 91)
phase change (p. 103)
quality of energy (p. 113)
radiation (p. 103)
relative humidity (p. 109)
second law of
thermodynamics (p. 112)
specific heat (p. 99)
temperature (p. 93)
APPLYING THE CONCEPTS
1. The temperature of a gas is proportional to the
a. average velocity of the gas molecules.
b. internal potential energy of the gas.
c. number of gas molecules in a sample.
d. average kinetic energy of the gas molecules.
2. The kinetic molecular theory explains the expansion of a
solid material with increases of temperature as basically the
result of
a. individual molecules expanding.
b. increased translational kinetic energy.
c. molecules moving a little farther apart.
d. heat taking up the spaces between molecules.
3. Two degree intervals on the Celsius temperature scale are
a. equivalent to 3.6 Fahrenheit degree intervals.
b. equivalent to 35.6 Fahrenheit degree intervals.
c. twice as hot as 1 Celsius degree.
d. none of the above.
4. A temperature reading of 2°C is
a. equivalent to 3.6°F.
b. equivalent to 35.6°F.
c. twice as hot as 1°C.
d. none of the above.
5. The temperature known as room temperature is
nearest to
a. 0°C.
b. 20°C.
c. 60°C.
d. 100°C.
6. Using the absolute temperature scale, the freezing point of
water is correctly written as
a. 0 K.
b. 0°K.
c. 273 K.
d. 273°K.
7. The metric unit of heat called a calorie is
a. the specific heat of water.
b. the energy needed to increase the temperature of 1 gram of
water 1 degree Celsius.
c. equivalent to a little over 4 joules of mechanical work.
d. all of the above.
8. Which of the following is a shorthand way of stating that “the
temperature change of a substance is directly proportional to the
quantity of heat added”?
a. Q m
b. m Tf Ti
c. Q ∆T
d. Q Tf Ti
9. The quantity known as specific heat is
a. any temperature reported on the more specific absolute
temperature scale.
b. the energy needed to increase the temperature of 1 gram of a
substance 1 degree Celsius.
c. any temperature of a 1 kg sample reported in degrees
Celsius.
d. the heat needed to increase the temperature of 1 pound of
water 1 degree Fahrenheit.
10. Table 5.1 lists the specific heat of soil as 0.20 kcal/kgC° and
the specific heat of water as 1.00 kcal/kgC°. This means that if
1 kg of soil and 1 kg of water each receive 1 kcal of energy,
ideally,
a. the water will be warmer than the soil by 0.8°C.
b. the soil will be 5°C warmer than the water.
c. the water will be 5°C warmer than the soil.
d. the water will warm by 1°C, and the soil will warm by
0.2°C.