The nature of heat

THE NATURE of HEAT
As we know, heat is a form of energy. In the form of
infrared radiation, heat from the sun travels through space
at the speed of 186,000 miles per second. Upon arriving on
earth, much of the radiant heat is absorbed by different
kinds of matter and is converted into heat that we can feel
(sensible heat). When you sit in the sun for a period of time
on a clear spring day, you may find that your clothing and
other objects around you become warm. Similarly, when
you walk barefoot across a beach on a summer day, you
may find the sand so hot that it burns your feet. In both
cases, radiant heat from the sun has been absorbed by
matter, and has been converted into heat that you can feel.
In this chapter, we will study the effect of heat upon matter.
HEAT AND THE MOTION OF MOLECULES
Have you ever tried to drill a hole through a piece of
metal? Both the drill and the metal become very hot.
Around 1800, an English scientist named Count Rumford
noted that, when a drill was used to bore a cannon, the bit
of the drill and the cannon both became very hot. To keep
the metals cool, he placed a cylinder of water around the
end of the cannon. As the boring continued, the water
became warmer and eventually boiled. Since the bit of the
drill and the cannon were cold at the start, Rumford
concluded that the heat produced probably came from the
friction created by the particles of the metal of the bit
rubbing against the particles of the metal of the cannon.

Further, he theorized that the motions of the particles in the
metals themselves (atoms or molecules) generated the heat.
Recall that various forms of energy can be converted into
other forms: When electrical energy passes through a thin
wire, as in a toaster, the wire becomes hot (electrical energy
to heat energy). When you rub your hands together, heat is
produced from the friction between the rubbed surfaces
(mechanical energy to heat energy). In general, when any
form of energy is absorbed by matter, the energy is
changed to heat. This may be explained by the kinetic-
molecular theory: The energy excites the molecules in the
matter, causing them to move faster and to collide more
frequently. As more collisions take place, more heat is;
produced.
The effect of heat energy on the motion of molecules can
demonstrated by using a sealed tube containing a little
mercury with some glass beads floating on the surface of
the mercury. At ordinary temperatures, the glass beads
merely float on the surface of the mercury. However, when
the tube of mercury is heated, the glass beads, bounce up
and down in a violent but random fashion. As still more
heat is supplied to the sealed tube, particles of mercury
begin to move more swiftly. The glass beads are repeatedly
struck by many mercury particles at the same time;
consequently, the glass beads begin to move randomly
themselves. Thus, Rumford's theory that heat is related to
the motions of molecules appears to be correct.
According to the kinetic-molecular theory, heat energy
acquired by a body is transformed into increased kinetic

energy of the molecules of the body. We observe this
increased kinetic energy whenever a solid, a liquid, or a gas
expands on heating. A further increase in kinetic energy
will eventually cause the particles of a solid or liquid to
become a gas.
Recall that when an ice cube (a solid) is heated, it melts
and becomes liquid water. When the water is heated, it
vaporizes and becomes gaseous water. According to the
kinetic-molecular theory, as increasing amounts of heat are
supplied to a piece of ice, the water molecules move more
rapidly until they gain sufficient energy to overcome the
attractive forces holding them together. This permits the ice
to liquefy and become water. Similarly, as still more energy
is received, the water molecules move at even greater
speeds. The attractive forces in the liquid are weakened and
the water is converted into gaseous water.
EXPANSION OF SOLIDS
Your laboratory experience with the ball and ring
apparatus indicated the effect of heat on volume. The
increase in size is not due to an increase in the size of the
particles that make up the solid ball, but rather to an
increase in the average distance between the particles.
When an object is heated, its particles vibrate faster, collide
more violently, and consequently move farther apart,
thereby increasing the volume of the object.
When the object is cooled, the opposite change occurs and
the volume of the object decreases. This decrease in
volume is calledcontraction.

The expansion of solids by heating may cause serious
practical problems. For example, the expansion of railroad
tracks, bridges, or the concrete in a roadbed can create
dangerous situations. Thus, allowance for the expansion of
solids daring hot weather must he made in the construction
of rails, bridges, and roads. For example, when rails arc
laid, gaps between the ends of the rails provide for
expansion. If this were not done, consider what would
happen to the railroad tracks on a very hot day. The metal
would expand, making the tracks bend andbuckle, which
might cause an oncoming train to be derailed. In bridge
construction, expansion joints allow for changes in the
length of the bridge. Concrete roadbeds are built with
spaces between the sections of concrete to allow for
expansion.
The contraction of solids, by cooling, may also present
problems. Thus telephone and electrical wires are strung
loosely to prevent their snapping as contraction takes place
during the colder times of the year.
UNEQUAL EXPANSION OF SOLIDS
Through extensive studies, scientists have found that
different metals expand at different rates when they are
heated. For example, when a piece of iron and a piece of
aluminum of equal size are heated together, we find that the
aluminum expands more than twice as much as the iron.
When two strips of different metals are fastened together,
they form a compound bar, or bimetallic strip, which is
employed in, useful devices such as thermostats and
metallic thermometers. In these devices the two different

metals, usually brass and steel, are welded together. When
the bar is heated, it bends because the brass expands more
than the iron and becomes longer than the iron.
Thus, the brass strip will be on the outside of the bend. As
it cools, the bar returns to its original shape.
The thermostat is a device containing a compound liar that
regulates the heating systems of our homes. When the
temperature in the house falls below the setting on the
thermostat, the compound bar, which contracts as it cools,
closes the circuit, turning on the heat. As the room is
warmed, the compound bar in the thermostat expands,
bends, and thereby breaks the circuit, shutting off the heat.
The bimetallic thermometer is often used as an oven
thermometer to indicate the temperature within an oven, or
within a piece of meat that is cooking.
EXPANSION OF LIQUIDS
Liquids, like solids, expand when heated. In the laboratory
experience we demonstrated that when water is heated, it
expands. When the same water is cooled to its original
temperature, the water contracts to its original volume.
Many other liquids, such as alcohol and mercury, behave in
the same way. At lower temperatures, however, the
behavior of water is an exception to this rule. As water is
cooled from 100° C to 4° C, it contracts-like other liquids
do. However, when water is cooled below 4° C, the water
expands-unlike other liquids. Water continues to expand
until it reaches 0° C, its freezing point. It has been found, as

shown in that the spaces between the water molecules in ice
are larger than the spaces between the water molecules in
liquid water. Ice is therefore said to have an open structure.
Thus, as ice is formed, the need for increased space
between the molecules causes the volume of the ice to be
greater than that of the water from which it was formed.
(This expansion in volume begins as liquid water is cooled
below 4° C.) Since the volume of ice is greater than the
volume of water from which the ice is formed, the density
of ice is less than the density of water. (Recall that density
equals weight divided by volume.) This is why ice floats on
water.
Like solids, different liquids expand at different rates. As
we will see later, the expansion of liquids is used in alcohol
and mercury thermometers.
The expansion of liquids must be considered in certain
heating systems. In a hot-water heating system, allowances
must be made for the expansion of heated water. As the
furnace heats the water in the heating system, the water
expands. If the expansion continues, pressure would build
up in the pipes and could damage the entire system. To
avoid this difficulty, an expansion tank is provided. Excess
heated water enters the expansion tank and thereby reduces
the pressure in the system.
EXPANSION OF GASES
Cases, like solids and liquids, expand when heated. Our
laboratory experience indicated that, as air is warmed, it
expands. Scientists have made similar observations with

other gases which indicate that gases confined in an elastic
container expand when they are heated and contract when
they are cooled.
The expansion of gases by heat must be considered by
automobile tire. manufacturers, since tires may burst if
allowed to remain in the sun indefinitely. A less serious
hazard caused by expanding gases is that bottles of soda
may crack or even explode if they are ex posed to heat for a
considerable length of time.
Products such as whipped cream, shaving cream,
deodorants, and insect repellents are now supplied in
aerosol cans. These cans contain the product itself and a
gas that forces the product out of the can when the valve, is
open. When the product is used up, the can still contains
unused gas. If this can is thrown into incinerator, the gas
becomes heated, expands, and may cause the can to
explode. Such aerosol cans should be discarded in a manner
that does not involve
heating.
Different solids and liquids expand at different rates
when heated. Gases, however, generally expand at the
same rate when heated to the same temperature, at a given
pressure.

TEMPERATURE
Heat and temperature are two terms that are often
confused. We know that the temperature of a small sample

of molten iron is considerably higher that the temperature
of the water in the ocean. However, the total heat in a
sample of molten iron is much less than the total heat of the
water in the ocean.
Scientist now accept Rumford’s theory that heat is
related to the motions of particles in matter. Thus, heat
depends on the total kinetic energy of the particles in a
body. Recall the equation relating kinetic energy with mass
and velocity:
K.E. = ½ m v2. Thus, the total kinetic energy of the
particles in a body depends on the number of particles
(mass) and the velocity of these particles.
Because the water in the ocean is colder than the
sample of molten iron, the velocity of the particles in the
water is less than the velocity of the particles in the molten
iron. However, the much larger quantity (mass0 of water
compensates for the smaller velocity of the particles and
thus the particles of water in the ocean possess greater
kinetic energy. This means that there is more heat in the
water in the ocean than in a small sample of molten iron.
But why is the temperature of molten iron
higher? Temperature, unlike heat, depends on the average
kinetic energy of the particles, that is, the kinetic energy per
particle. To find this average, we divide the total kinetic
energy by the number of particles. Thus, the large mass of
ocean water has a smaller average kinetic energy per
particle and consequently has a lower temperature than a
small sample of molten iron.

MEASURING TEMPERATURE
Instruments designed to measure temperature are called
thermometers. Most thermometers are based on the
principle that matter, on heating, expands and, on cooling,
contracts. In general, matter expands and contacts
regularly. This means that the amount of expansion or
contraction in length are generally equal for the same
increase or decrease in temperature. This regular
expansion and contraction has made it possible to construct
three different types of thermometers: gas (air), liquid
(mercury and alcohol), and solid (bimetallic)
thermometers.
The gas (air) thermometer
In this thermometer, the glass bulb contains air. When the
bulb is warmed, the air in the bulb expands and forces some
of the colored water out of the tube. This changes the level
of the liquid in the tube. By placing a suitable scale
alongside the tube, temperature changes can be measured.
Air thermometers of this type, while interesting, are not
very accurate because the volume of a gas is also
influenced by the air pressure around it. (Note that the flask
contains a tube open at both ends. Why?).
LIQUID THERMOMETERS
Thermometers containing liquids such as mercury and
alcohol are useful and accurate because these liquids
usually expand and contract uniformly (regularly).

Mercury thermometers are made by filling a thin glass tube
with mercury at a temperature greater than the maximum to
be measured. The tube is then cut and sealed at the top.
When the mercury cools, it contracts, leaving a partial
vacuum above the mercury. (Liquids expand and contract
to a much greater extent than do solids. Thus, in the given
temperature range, the glass tube is scarcely affected by the
temperature change.) The partial vacuum eliminates the
effect of air resistance to the expansion of the mercury. The
scale of the. thermometer is generally fixed by locating the
boiling and freezing points of water on the scale. The
distance between the boiling and freezing points is then
divided into units depending on the temperature scale used.
This will be discussed in the next section.
Since mercury freezes at -39° C, it cannot be used to
measure very low temperatures. However, mercury boils at
357° C, which means that a mercury thermometer can be
used to measure temperatures above the boiling point of
water. On the other hand, alcohol freezes at -114° C.
Accordingly, alcohol thermometers are used to measure
low temperatures. However, since alcohol boils at 78° C,
alcohol thermometers cannot be used to measure high
temperatures.
SOLID (BIMETALLIC)
Recall that a bimetallic strip behaves as it does because
different metals expand at different rates. Because most
metals melt only at very high temperatures, a thermometer
that uses a bimetallic strip can measure temperatures as
high as 1000° C. The dial thermometer, used in most

homes as an oven thermometer, is an example of a
bimetallic thermometer. A curved bimetallic strip, with the
faster-expanding metal on the outside of the bend, is
attached to a pointer. Upon heating, the bimetallic strip
moves, causing the pointer to indicate the temperature on a
circular scale.
Other metallic thermometers, called resistance
thermometers, use the principle that the resistance of a wire
changes with temperature. Such thermometers also measure
high temperatures.
THE FAHRENHEIT
AND CELSIUS TEMPERATURE SCALES
Temperature markings on thermometers are indicated in
Fahrenheit degrees or Celsius degrees. The Fahrenheit and
Celsius scales are named after their originators, Gabriel
Fahrenheit and Anders Celsius. (The Celsius scale is also
called the centigrade scale.) Both Fahrenheit and Celsius
scales are calibrated by using the boiling and freezing
points of water. The Fahrenheit scale is used in the English
system of measurement and the Celsius scale in the metric
system. In the Fahrenheit scale, the freezing point of water
is 32.° F, and the boiling point of water is 212° F. The
remainder of the scale between these two points is marked
off into 180 equal divisions (212 - 32 = 180) . In the
Celsius scale, the freezing point of water is 0° C and the
boiling point of water is 100° C. The remainder of the scale
between these points is divided into 100 equal divisions
(100 - 0 = 100). Note that there are 180 divisions between
the freezing and boiling points of water in the Fahrenheit

scale and 100 divisions between these points in the Celsius
scale. Thus, each Celsius division ( degree ) is 9/5 as large
as Fahrenheit division. This relationship, together with the
fact that there are 32 Fahrenheit divisions between 0 °F and
32° F, makes it possible to convert one scale into the other
by using the following formulas:
° C = 5/9(° F – 32) ° F = 9/5 °C + 32
THE KELVIN SCALE
Confined gases, like most solids and liquids, expand and
contract uniformly. For this statement to be true, however,
a gas must be heated or cooled in such a way that the
pressure remains constant. ( Recall that the air thermometer
is inaccurate because it is affected by surrounding air
pressure.) If we start at 0° C, we find that, for every Celsius
degree rise in temperature, the volume of a gas increases
273 of its original volume ( provided the pressure does not
change). Similarly, if we again start from 0° C, we find that
for every Celsius degree drop in temperature, the volume
decreases 273 of its original volume. At -273° C, the
volume of a gas would shrink to zero and all molecular
motion would cease. This, in turn, means that the gas
would contain no heat.
(Actually, gases generally liquefy before this temperature is
reached.) Scientists refer to -273°C as absolute zero, a
temperature that has never been attained, although some
scientists have come very close to this point.

Absolute zero, -273°C, is also called 0 Kelvin ( 0 K ). The
Kelvin scale, named after its originator, Lord Kelvin, is
based on absolute temperatures. Since the Kelvin scale
begins with absolute zero (-273° C), we use the following
formula to convert the Celsius scale to the Kelvin scale:
degrees Kelvin = degrees Celsius + 273 degrees
This formula may be written as K = ° C + 273
Let us find the freezing point of water (0° C) in the Kelvin
scale: K=0 + 273=273 K
Thus, 0° C is equivalent to 273 K.
Now, let us find the boiling point of water: K =100
+ 273 = 373 K
Thus, 100° C is equivalent to 373 K.
TRANSFER of HEAT
When a metal spoon is placed in a bowl of hot soup, the
entire spoon soon becomes hot because the heat travels
from the soup to the bowl-shaped part of the spoon, and
then to the handle. When ice is placed in warm water, the
ice soon melts. Both of these examples show that heat
travels from one body to another. Generally, when objects
are at different temperatures, heat is transferred from the
warmer object to the cooler object until both objects are at
the same temperature. Heat transfer can occur through one
of three methods: conduction, convection, or radiation.
CONDUCTION

When one end of a metal rod is held in a flame, the entire
rod will become hot enough to burn the hand. The heat
from the flame reaches the hand by traveling through the
rod. Substances that allow heat to travel through them are
called conductors. In general, as we learned before, metals
are good conductors. However, some metals conduct heat
more readily than others. This can be demonstrated by
inserting rods of aluminum, copper, iron, nickel, and brass
into a brass sphere or disk and then attaching a small ball of
wax to the end of each rod. When the center of the brass
disk is heated, the wax at the tip of each metal melts in the
order in which the different metals conduct heat. The wax
at the tip of the copper melts first and the wax at the tip of
the iron melts last.
Conduction in most materials can be explained by the
kinetic-molecular theory. When one end of a rod is heated,
the molecules in that end of the rod vibrate faster and strike
other nearby molecules, causing them to vibrate faster. In
this manner, the increased molecular motion is transferred
from one end of the rod to the other, permitting the heat to
travel through the rod.
Substances that do not readily allow heat to pass through
them are called insulators. Gases and liquids are generally
poor conductors of heat because their molecules are farther
apart than are the molecules in solids. Therefore,
neighboring molecules in a gas or in a liquid are less
affected by the increased motions of heated molecules, and
consequently heat is not conducted rapidly.

Substances like wood or plastic are poor conductors of
heat, so they are used to make handles for metallic objects
that are to be heated. The clothing we wear is also a poor
conductor of heat, enabling us to retain body warmth.
Porous material is generally non-conducting because it
contains layers of trapped air which do not permit heat
transfer.
CONVECTION
Although gases and liquids are poor conductors of heat,
heat is transferred through them by the process of
convection. Convection is the transfer of heat due to the
motion of the liquid or gas itself. For example, when a
beaker of water is heated the water layer closest to the heat
source is warmed slowly by conduction. As the water
becomes warmer, it expands, becomes less dense, and rises.
This brings heat to the upper layer. At the same time,
cooler water from the upper portion of the beaker moves
down, takes the place of the rising water, and becomes
heated itself. When warm enough, this water rises and
carries heat upward. As these processes continue, heat that
enters the bottom of the beaker is distributed throughout the
beaker until all the water is at the same temperature. The
moving water in such a case is said to have set up a
convection current.
Heat is also transferred through gases by convection. It is
by this means, in part, that a stove or a radiator heats a
room. Heat from the radiator warms the air above it,
causing the air to expand, become less dense, and rise. The
cooler air that moves in to take the place of the warmed air

is also soon warmed. As this air rises, a convection current
is established. The convection current continues to
distribute heat throughout the room until the entire room is
warmed.
The formation of a convection current in air is
demonstrated with a convection box apparatus. First the
candle is lighted, then smoking touch paper is placed
over the chimney, opposite the candle. The smoke, coloring
the air, can be seen to move down this chimney, across the
box, and out through the other chimney. This occurs
because the air over the candle is heated, becomes less
dense, and rises, leaving a partial vacuum. Cooler, more
dense air from the first chimney moves in to fill the partial
vacuum. This cycle continues as long as heat is given off
by the burning candle.
RADIATION
We know that light energy and heat energy travel from the
sun to the earth through space, which is an almost perfect
vacuum. These forms of energy, traveling without the aid
of molecular collisions, are transferred from the sun to the
earth by radiation, that is, by means of rays, or waves. You
can understand this method of heat transfer by standing a
short distance from an open fire or by placing your hand a
little to one side of, but not touching, a hot radiator. Since
neither source of heat is being touched, you cannot receive
heat by conduction. Since warm air rises vertically from the
heat source, the heat cannot reach you by convection. The
heat that is transferred to you from the fire or radiator
reaches you by radiation.

The heat radiated by one body ( the sun, for example) is
most rapidly absorbed by other bodies that are black in
color and rough in texture. In warm climates, white
clothing which reflects the radiant heat of the sun is cooler
than dark clothing which quickly absorbs the radiant heat.
Similarly, bodies that are rough and dark tend to radiate
heat better than shiny smooth bodies. This is why steam
radiators are often dark and have a roughened surface. It is
for the same reason that coal burning stoves are black.
Bodies that are shiny and smooth do not absorb heat
readily. Instead, these bodies reflect heat. Thus, aluminum
used for roofing keeps homes cool in the summer and warm
in the winter. This principle is utilized in the thermos
(vacuum) bottle, which is so constructed as to permit
liquids to retain their temperatures for a long time. A
thermos bottle is double walled, with a partial vacuum
between the walls to prevent heat transfer by conduction or
convection. A cork stopper also prevents heat transfer by
conduction. The inner glass walls are silvered to reflect
radiant heat back into the liquid, thereby minimizing heat
loss by radiation. Thus, a hot liquid remains hot because
heat is lost very slowly. A cold liquid remains cold in
thermos bottles because outside heat enters very slowly by
conduction, convection, or radiation.
MEASURING HEAT
We learned that temperature is a measure of
the average kinetic energy of the molecules of a substance.
This is the same as saying that temperature represents the
average intensity of the motion of the molecules, or the

degree of hotness of a substance. Average kinetic energy
means the total kinetic energy divided by the total number
of particles. Recall that the ocean contains much
more heat than does a small amount of molten iron.
However, since the ocean contains many more particles
than the molten iron, the temperature of the ocean (that is,
the total kinetic energy divided by the total number of
particles) is lower than that of the molten iron.
Temperature, therefore, does not tell us the quantity of heat
present. The quantity of heat represents the total kinetic
energy contained by allof the particles of the substance.
In the metric system, we measure the quantity of heat by a
unit called the calorie. The calorie is the amount of heat
needed to raise the temperature of 1 gram of water 1
Celsius degree. In the English system, heat is measured by
a unit called the British Thermal Unit (BTU). A BTU is the
amount of heat needed to raise the temperature of 1 pound
of water 1 Fahrenheit degree. The amount of heat energy
present in a substance cannot be measured directly with
simple measuring devices. Instead, it is measured by
observing its effect on a given quantity of water in a device
called a calorimeter. One type of calorimeter, consists of
two polished metal cups surrounded by air, a poor
conductor of heat. An insulating cover, holding a
thermometer, makes up the top of the calorimeter. The
polished cups reflect heat, thus maintaining the temperature
of the liquid in the container.
To determine the amount of heat energy absorbed (or lost)
by a given quantity of water, we multiply the weight of the

water in grams by the change in temperature of the water in
Celsius degrees. Thus: amount of heat = weight of water X
change in temperature
In a calorimeter, when 20 grams of water at 20°C are
heated to a temperature of 30°C, how much heat is
absorbed?
The temperature change = the final temperature - the initial
temperature = 30°C - 20°C = 10 Celsius degrees.
Substituting, amount of heat = 20 grams X 10 C° = 200
calories
We conclude that 200 calories have been ab sorbed. (We
assume that no heat has escaped from the calorimeter.)
HEAT EXCHANGE IN WATER
In a calorimeter, when a quantity of water at a given
temperature is mixed with a quantity of water at a different
temperature, the amount of heat lost by the "hot" water is
equal to the amount of heat gained by the "cold" water.
Suppose we mix 100 grams of water at 90° C with 100
grams of water at 40° C and find that the final temperature
of the mixture is 65° C. Let us calculate the number of
calories lost. by the hot water and gained by the cold water.
The temperature of the hot water dropped from 90° C to
65° C, a decrease of 25 Celsius degrees. Since we began
with 100 grams of hot water that underwent a temperature
change of 25° C, we determine the amount of heat lost:

amount of heat = weight of water X change in
temperature
(calories) (grams) (Celsius degrees)
amount of heat = 100 grams X 25° C amount of heat =
2500 calories
The minus sign in the answer indicates that 2500 calories
of heat have been lost.
The temperature of the cold water increased from 40° C
to 65° C, an increase of 25 Celsius degrees. Since we began
with 100 grams ofcold water that underwent a
temperature change of 25° C, we determine the amount of
heat gained:
amount of heat = weight of water X change in
temperature
(calories) (grams) (Celsius
degrees)
amount of heat = 100 grams X 25° C amount of heat =
2500 calories
Note that the amount of heat lost by the hot water (2500
calories) is the same as the amount of heat gained by the
cold water (2500 calories). We assume that the heat
exchange was "perfect" and that no heat escaped from the
calorimeter.
The quantity of heat needed to raise the temperature of 1
gram of a substance 1 Celsius degree is called the specific
heat of the substance. For water, the specific heat is 1. This

means that 1 calorie of heat will raise the tem perature of 1
gram of water 1° C. Water is the only substance for which
this is true. Other substances vary in the quantity of heat
needed to raise 1 gram of the substance 1 Celsius degree.
Consequently the formula
amount of heat =weight of water X change in
temperature applies only to water.
CALORIES AND FOOD
Your body requires energy in order to per form its daily
tasks. Most of this energy comes from energy-rich foods
such as carbohydrates and fats. This energy is released
when the body utilizes these foods. Using special calo
rimeters, scientists have measured the energy content, or
the number of calories present, in fixed quantities of certain
foods. For example, a slice of white bread contains about
60 000 calories; a typical chocolate bar may contain about
300 000 calories.
Nutritionists use a special kind of notation when
discussing the calorie content of foods. They define a food
Calorie ( written with a capital letter) as 1000 calories. On a
calorie table, therefore, we would read that a slice of white
bread contains about 60 Calories and that a chocolate bar
contains about 300 Calories.
This information comes from "Physical Science
Workbook", 1961

The History of Heat

Aristotle and the Greeks had their idea of fire as one
of the 4 Primal Elements. Even the ancients realized that
heat and light were not alike as aspects of fire,
though. After the fire had gone out and the light gone, the
heat of the kettle and its contents remained.

First modern chemist to study heat was Joseph
Black (1728 - 1799). Black tried to explain heat in terms of
a fluid. He explained how a kettle of water placed over a
fire increased in temperature but a kettle filled with water
and ice placed over a fire did not change in temperature till
all the ice was melted. He said that until the ice was
saturated with the heat-fluid and thus became melted could
its temperature rise. Lavoisier accepted this theory and
gave the name for this heat-fluid “caloric” from the Latin
word for heat.

Another idea competed with the caloric
theory. Scientist knew that kinetic energy of motion plus
the stored energy called potential energy was given the
name mechanical energy and that friction was a part of the
conservation of these energies. They knew friction could
warm up an object so maybe the invisible motion of
invisible particles was what we call heat. Summed up;
friction was converting mechanical energy into heat. The

problem was this idea of really small particles of matter
(i.e., atoms and molecules).

Count Rumford (really Benjamin Thompson - a spy
for the British authorities during the Revolutionary War)
was to supervise the boring of cannon for Bavarian
army. Using a horse to work a treadmill he realized that
the solid block of brass grew hot as the borer cut its way
in. Rumford calculated that if the caloric theory were
correct the heat released during the boring would have
melted the entire block of metal first. He pointed out that
heat was produced without fire, without light, without
chemical combustion. It came just out of motion.

John Dalton comes along with his ideas of atoms
and the kinetic energy idea of heat began to gain favor. An
English physicist, James Prescott Joule (1818-1889) was
attempting to find the mechanical equivalent of heat. In the
end he found that a given amount of energy of whatever
form always yielded that same amount of heat (at 4.18
joules per calorie). The relationship of the motion of atoms
to temperature and heat was placed on firm theoretical
basis about 1860 by the Scottish physicist James Clerk
Maxwell.

Specific Heat

The ability of water to stabilize temperature
depends on its relatively high specific heat. The specific
heat of a substance is defined at the amount of heat that
must be absorbed or lost for 1 g of that substance to change
its temperature by 1º C. The specific heat of water is 1.00
cal/g ºC. Compared with most other substances, water has
an unusually high specific heat. For example, ethyl
alcohol, the type in alcoholic beverages, has a specific heat
of 0.6 cal/g ºC.
Because of the high specific heat of water relative
to other materials, water will change its temperature less
when it absorbs or loses a given amount of heat. The reason
you can burn your finger by touching the metal handle of a
pot on the stove when the water in the pot is still lukewarm
is that the specific heat of water is ten times greater than
that of iron. In other words, it will take only 0.1 cal to raise
the temperature of 1 g of iron 1ºC. Specific heat can be
thought of as a measure of how well a substance resists
changing its temperature when it absorbs or releases
heat. Water resists changing its temperature; when it does
change its temperature, it absorbs or loses a relatively large
quantity of heat for each degree of change.
We can trace water’s high specific heat, like many
of its other properties, to hydrogen bonding. Heat must be
absorbed in order to break hydrogen bonds, and heat is
released when hydrogen bonds form. A calorie of heat
causes a relatively small change in the temperature because
must of the heat energy is used to disrupt hydrogen bonds
before the water molecules can begin moving faster. And

when the temperature of water drops slightly, many
additional hydrogen bonds form, releasing a considerable
amount of energy in the form of heat.
What is the relevance of water’s high specific heat
to life on Earth? By warming up only a few degrees, a
large body of water can absorb and store a huge amount of
heat from the sun in the daytime and during summer. At
night and during winter, the gradual cooling water can
warm the air. This is the reason coastal areas generally
have milder climates than inland regions. The high specific
heat of water also makes ocean temperatures quite stable,
creating a favorable environment for marine life. Thus,
because of its high specific heat, the water that covers most
of planet Earth keeps temperature fluctuations within limits
that permit life. Also, because organisms are made
primarily of water, they are more able to resist changes in
their own temperatures than if they were made of a liquid
with a lower specific heat.
Water is one of the few substances that are less dense as a
solid than as a liquid. While other materials contract when
they solidify, water expands. The cause of this exotic
behavior is, once again, hydrogen bonding. At
temperatures above 4º C, water behaves like other liquids,
expanding as it warms and contracting as it cools. Water
begins to freeze when its molecules are no longer moving
vigorously enough to break their hydrogen bonds. As the
temperature reaches 0º C, the water becomes locked into a
crystalline lattice, each water molecule bonded to the
maximum of four partners. The hydrogen bonds keep the

molecules far enough apart to make ice about 10% less
dense than liquid water at 4º C. When ice absorbs enough
heat for its temperature to increase to above 0º C, hydrogen
bonds between molecules are disrupted. As the crystal
collapses, the ice melts, and molecules are free to slip
closer together. Water reaches it greatest density at 4º C
and then begins to expand as the molecules move faster.
The ability of ice to float because of the expansion
of water as it solidifies is an important factor in the fitness
of the environment. If ice sank, then eventually all ponds,
lakes, and even the oceans would freeze solid, making life
as we know it impossible on Earth. During summer, only
the upper few inches of the ocean would thaw. Instead,
when a deep body of water cools, the floating ice insulates
the liquid water below, preventing it from freezing and
allowing life to exist under the frozen surface.

Specific Thermal Electrical
Metal Density
Heat Conductivity Conductivity

k
cp g/cm3 1E6/Ωm
watt/cm K
cal/g° C
Brass 0.09 1.09 8.5
Iron 0.11 0.803 7.87 11.2
Nickel 0.106 0.905 8.9 14.6
Copper 0.093 3.98 8.95 60.7
Aluminum 0.217 2.37 2.7 37.7
Lead 0.0305 0.352 11.2

Heat and Temperature Teaching Notes
1) Heat flows from hot to cold areas due to a temperature
difference only.
example: A small hot block of a material is placed next
to a larger, cooler block. Heat flows from the
small hot block to the larger block till
equilibrium is reached.
2) Note the difference between the heat content and
temperature. A lake may be cooler in temperature than a
liter flask of water but the lake has a much greater heat
content due to the vast number of particles and their
associated motion.
3) Human perception of heat and temperature is not
adequate for scientific work. So we must investigate tools
that provide the accuracy and repeatability we need.
4) For temperature we will be using thermometers. There
are other devices that allow you to measure temperature.
5) For heat we will be working with simple calorimeters.
We will look at these devices when we look at the transfer
of heat.

Thermometers and Temperature Scales see
comparisons charts

thermometers = instruments to measure temperature
see drawing: gas (air) thermometers
see drawing: liquid (Hg and alcohol)
see drawing: solid (bimetallic)
Mercury thermometer -- fill thin glass tube with Hg at a
temp. greater than the maximum to be measured; tube is cut
and sealed; the Hg cools and contracts leaving a partial
vacuum above the Hg (eliminates effects of air resistance
on expansion of Hg); then calibrate thermometer
Hg freezes at -39° C so it cannot be used for low
temperatures (use alcohol which freezes at -114° C)
can use Hg for high temperature boils at 357° C (alcohol
cannot be used for high temperature work due to its low
boiling point - 78° C)
use alcohol thermometers in schools unless extreme
accuracy is needed due to safety factor. The alcohol may be
inaccurate by 1 - 2 degrees but this may not be a problem if
you are looking at changes in temp.

Calibration
Both Celsius (Anders Celsius) and Fahrenheit
(Gabriel Fahrenheit) scales are established by using
the boiling and freezing point of water at 1 atmosphere of
pressure.
step 1 -- establish a mixture of ice and water in equilibrium
(0° C) mark point of liquid in thermometer at 0° C
establish a mixture of steam and water at equilibrium (both
at a pressure of 1 atm.) steam condensing and water
vaporizing)
label this point of liquid as 100 °C
step 2 -- divide the interval between 0° C and 100° C into
100 equal parts, each representing a change in temperature
of 1° C.
using this scale you can extend your marks below 0 °C and
above 100 °C as far as ,you wish.
The Fahrenheit scale labels the freezing point at 32°
(his label for the temperature he could achieve with an ice
and water mixture and labeled the boiling point temperature
of water at 212 ° which was a number chosen for
convenience apparently creating 180 divisions.

You might ask your self about the amount of heat energy
need to cause a 1 degree change in temperature on a
Celsius scale compared to that needed on a Fahrenheit
scale. (more heat needed to cause change of 1 degree on
Celsius scale.)
KELVIN scale
We know that gases decrease in volume 1/273 of
its original volume for each degree drop in
temperature. Thus at -273° C the volume of gas would
shrink to zero and the gas would have no molecular
motion. We know this is impossible (particles have zero-
point energy). To label these very low temperatures a
scale called the absolute or Kelvin scale is often used. It
designates -273° C at the zero point, and is called called
absolute zero.
Extrapolation to absolute zero: A good research project
for students or a quick review of graphing can be done by
using the method to calculate absolute zero. Use a capillary
tube with a trapped bubble of air between light machine oil.
Measure the length of the bubble after placing it in different
temperature mixtures. Plot the length versus temperature.
Since it is not easy to obtain very cold temperatures, the
linear series of points that you did obtain should allow you
to extrapolate, (extend a curve beyond the known data
points following the apparent pattern of the curve) until it
intersects the temperature axis. See actual plot!

Transfer of Heat by Conduction, Convection, Radiation
Conduction is a consequence of the kinetic behavior of
matter. Faster vibrating particles collide with less energetic
neighbors and transfer some of their kinetic energy to the
slower moving particle.
Through successive molecular collisions energy travels
through a material without the average position of the
particles being changed. There must be a temperature
differential (one end of some object at a higher temperature
than the other) for heat to be conducted.
Gases are poor conductors of heat (compared to liquids
and solids) because the molecules are relatively far apart
and collisions are infrequent.
Metals have the greatest ability to conduct heat (for the
same reason as their high electrical conductivity). This is
due to a significant number of electrons being able to move
about freely instead of being bound permanently to
particular atoms.
Thermal conductivity of a material is a measure of its
ability to conduct heat.
Example: wood and metal (see class discussion)
Convection involves the actual motion of a hot fluid from
one place to another, displacing a colder fluid in its path
and setting up a convection current. Convection is the chief
mechanism of heat transfer in fluids.

Natural convection occurs when the buoyancy of heated
fluids leads to motion. Heated fluids (gas or liquid) expand
and becomes less dense than surrounding cooler fluids. It
then rises.
Radiation is defined as the energy that is transmitted by
electromagnetic waves and requires no material medium
for passage.
All objects radiate electromagnetic waves with the higher
the temperature of an object the shorter the predominating
wavelength of its radiation.
Example: see glass lined thermos bottle in heat packet in
class.
Transfer of Heat
Conduction -- place iron rod in fire -- the end you are
holding becomes warm due to conduction
Convection -- stove heats room by convection
Radiation -- heat the earth receives from the sun is
radiation
Natural direction of heat flow is from hot bodies to cold
ones.
Conduction -- conduction is a consequence of kinetic
behavior of matter
faster vibrating particle collide with less energetic
neighbor and transfer some of their kinetic energy the
slower moving particle

example: place hand on wood and metal sample at same
temperature. The metal will seem cooler because it
conducts the heat away much faster than the wood
Convection -- actual motion of hot fluid from one place
to another, displacing cold fluid in its path setting up a
convection current = chief mechanism of heat transfer in
fluids in most instances
natural convection - the buoyancy of heated fluids leads
to motion - heated fluid (gas or liquid) expands and
becomes less dense than surrounding cooler fluids and
rises
Radiation -- energy that is transmitted by
electromagnetic waves and requires no material medium
for passage
all objects radiate electromagnetic waves but the higher
the temperature of an object the shorter the predominate
wavelength of its radiation
There are many examples you can use to demonstrate these
three ideas but discussing a glass lined thermos bottle will
allow you discuss them as well as
High specific heat capacity material demonstrates relatively
small change in temperature for a given change in internal
energy content
add 1 calorie of heat to 1 gram of water, helium, ice, gold
temperature rises: water 1 °C

helium 1.3 °C
ice 2 °C
gold 33 °C

Calorimetry:
1) Use 2 polystyrene cups (one within the other) -- the
polystyrene will not absorb very much heat.
2) You may find if difficult for the students to be patient
when working with the calorimeters. Your step-by-step
procedure must be very simple and clear. This type lab is
also a very dangerous time for your thermometers.
3) Several labs have been included in the packet involving
the use of the calorimeters and thermometers: a) the heat of
reactions and heat of solutions labs deal with endothermic
and exothermic reactions (and you can incorporate the use
of the mole as review) b) the specific heat capacity lab is an
excellent demonstration of the principle of heat exchange
as well as specific heat capacity of different metals and
liquids. 1) for better results with this lab try to use as much
metal as possible and as little fluid as possible and still
cover the metal in the cups. Drain the metal samples
quickly after removing them from the boiling water. This is
a good time to have the students watch the boiling process
which will be one of the final topics in this unit. 2) the math
may be a bit difficult for you at first.

Thermal Expansion of water
1) From 0° C to 4° C the volume of water in a sample
decreases (the greatest density is at 4° C)
2) We know that ice floats (less dense) so that a body of
water in winter freezes from the top down. The ice is a poor
conductor of heat so that the initial layer of ice that freezes
impedes further freezing allowing fish and plant life to live
through the winter.
3) The spaces between molecules in ice are greater than
the same spaces in liquids.
4) Ice has what is called an Open Structure --> each
water molecule can participate in 4 bonds with other water
molecules, while other solid molecules can have as many as
a dozen bonds with surrounding molecules resulting in a
more compact substance.
5) As stated the density of the water increases from 0 °C
to 4 °C. Large clusters of water molecules break into
smaller clusters that occupy less space in the aggregate as
the temperature rises to 4 °C. Only above 4 °C does the
normal thermal expansion show a decreasing density
with increasing temperature. Above 4° C , the normal
thermal expansion of materials is seen. Here as the
temperature rises the density decreases.

Heat and Temperature

HEAT Heat is a form of internal energy which is
transferred from one object to another due to a difference in
temperature between the objects. Heat is the total energy of
motion of all particles (the total kinetic energies of all the
particles.)
TEMPERATURE The temperature of a body of matter
is a measure of the average kinetic energy of the random
motion of its particles. Temperature is the kinetic energy
divided by the number of particles. Temperature is that
property of a substance which determines whether it is in
thermal equilibrium with another object.
Thermal Conductivity - a measure of the ability of a
substance to conduct heat
THERMAL EQUILIBRIUM This is the situation in
which no heat moves from one object to another.
CALORIE A 15° Calorie is the amount of heat energy
needed to change the temperature of of 1 gram of water by
1° C (from 14.5° C to 15.5° C at 1 atmosphere of pressure).
1 calorie = 4.185 Joules and 1 kilocalorie = 1000 calories.
SPECIFIC HEAT CAPACITY This is the amount of
heat (in calories or Joules) that must be added or removed
from a unit mass of that substance to change its temperature
by one degree. Different substances have different
capacities because they absorb and release heat at different
rates.

WATER Water has a specific heat capacity of 1.00 cal/
g °C or 4.185 Joules/g °C. The SI unit would be 4185 J/kg
°C.
PRINCIPLE OF HEAT EXCHANGE The heat lost
by an object must equal the heat gained by the object to
which the heat is transferred. There must be a temperature
difference for heat to be transferred.
Q (heat energy) = m (mass) x At (temp.) x cp (specific heat
capacity)
(cal/Joules) (g) (°C) (cal/g °C)
or (Joule/g °C)
Problems:
1) How much heat energy is needed to raise the
temperature of 100 grams of water from 0 degrees to 30°
C?
2) A calorimeter contains 300 grams of water at 10° C.
After a food sample is burned in the calorimeter the water
temperature changes to 15° C. How much heat was given
off by the food sample?
LATENT HEAT Latent heat is the heat required to
bring about a change in state.
HEAT OF FUSION The heat of fusion is the amount of
heat that must be supplied to change a unit mass of the
substance at its melting point from solid to liquid. The heat
of fusion of water is 80 calories per gram (80 kcal/kg).

HEAT OF VAPORIZATION The heat of vaporization
is the amount of heat that must be supplied to change a unit
mass of the substance at its boiling point from liquid to gas
or vapor state. For water this is 540 cal/g or 540 kcal/kg.
HEAT OF SUBLIMATION The heat of sublimation
is the heat needed to change a solid to a gas.
HEAT OF CONDENSATION The heat of
condensation is the reverse of the heat of vaporization, it is
the heat given off when a gas condenses to a liquid.

Four States of Matter
Matter is defined as any material that has mass, occupies
volume, and exhibits inertia (resistance to movement).
Solids definite shape and volume, resist deformation
very close spacing of particles that make up the solid
these particles appear to vibrate about fixed points
particles vibrate faster at higher temperatures, slower at
lower temperatures
Crystalline solids particles are arranged in regular,
repeated patterns - said to have "long-range order" to their
structure -example would be NaCl (table salt)
Amorphous solids solids that lack the definite
arrangement in crystals are `amorphous' which means
`without form'. Can be thought of as liquids whose
stiffness is due to exaggerated viscosity. These solids are

said to have "short range order". Examples are pitch,
glass, plastics. Polymers are flexible and some will change
their structure when undergoing a physical change
(examples are rubber bands, Saran Warp, Lucite, DNA,
fats, cellulose, glycogen. Jello is a natural glucose
polymer.
Liquids definite volume, resist compression, will flow,
takes the shape of its container
greater spacing between molecules, liquid particles
appear to travel in straight line paths between collisions
but appear to rotate or vibrate about moving points
Gases Have no definite shape or volume, takes the shape
and volume of its container
can be compressed or dispersed, the particles vibrate very
rapidly, are relatively far far apart, and there are no forces
holding them together
Plasma very high temperature ionized gas (as high as
100 million degrees in some fusion reactors). These
plasmas have no fixed volume or shape, most are mixtures
that are not easily containable. They all have particles that
are electrically charged and of low density. The Milky
Way is a huge plasma.
Energy Definitions
Energy: having the ability to do work (move matter)
Work: a push or pull over some distance (force x distance)

Force: a push or a pull
Potential Energy: the energy a body possesses by virtue of
its position, composition, and or condition
stored energy or energy of position
P.E. = mass x gravity x height
examples: water behind a dam,
stretched/compressed spring, explosives
Kinetic Energy: the energy of motion (conserved in
every elastic collision)
K.E. = 1/2 mass x velocity2
heat energy flows from hot objects to
cooler ones through transfer of K.E. when particles collide
Momentum: mass time velocity (momentum is
conserved in every collision where there is no friction)
Linear momentum of a moving body is a measure of its
tendency to continue in motion at a constant velocity. The
conservation of linear momentum states that in the absence
of forces from outside the system the total momentum of
colliding particles cannot change but the distribution of the
total momentum may change. Momentum is redistributed
in a collision.
Intermolecular Forces:
potential energy forces that hold molecules together and
in correct position in solids
potential energy forces that hold molecules together in
liquids
the kinetic energy of the molecules in solids and liquids

cannot overcome intermolecular forces holding the
molecules together (so they do not fly apart)
gas molecules have enough kinetic energy to break free
from intermolecular forces or to keep such forces from
forming
Kinetic – Molecular Theory of Gases
gases are made up of molecules that are in continuous
motion
an increase in the temperature increases the speed of the
molecules, thus increasing the kinetic energy of the
substance
All gases are compressible
Gases display diffusion (random movement of molecules
from one area to another with a net change in
concentration – rate varies with temperature and
molecular mass)
Gases can be liquefied (called liquefaction)
Closed System Criteria
In using the above information we look at pressure,
temperature, and volume in a closed system.
1. In a closed system nothing escapes or is allowed in
(unless we choose to allow it)
2. all molecules are in motion (have K.E.)
3. molecules exert a uniform pressure on all surface areas
of the walls of the container
4. Pressure = force/area (see examples given in class)

5. Atmospheric pressure is the cumulative effect of the
force generated by the weight of the
atmosphere. Given values that must be used in
problems include: 14.7 lb/in2, 101.3 kPa, 1
atmosphere, 760 mm of Hg, 1 033.6 g/cm2
6. Molecules exert pressure on other molecules inside
container as they collide, push, and bounce off other
molecules
7. The pressure a gas exerts on the walls of its container
is the sum of the forces acting on the walls (equals the
frequency of collisions with the walls of the container
plus the force of each molecule as it pushes against the
wall) due to the random collision of limitless numbers
of these moving molecules.
Collisions that occur between molecules are perfectly
elastic, the particles bounce off each other and exchange
energy, but there is no loss of energy
* elastic atomic collisions: atoms (molecules) bounce
back as far/fast as it would have had it not collided (no
change in the total kinetic energy of the two particles
before and after the collision)
* inelastic collisions: the normal order in which the
objects lose energy and slow down
Momentum is conserved in every collision where there is
no friction, energy is conserved only in elastic collisions.
Gas Laws

1. J.L. Gay-Lussac’s Law If the volume remains
constant, the pressure is directly proportional to the
absolute temperature:
P ~ T P1 / T1 = P2 / T2
2. Boyle’s Law If the temperature remains constant,
the volume of a gas varies inversely with the pressure:
V ~ 1/P P1 V1 = P2 V2
3. Charles’ Law If the pressure is kept constant, the
volume of a gas is directly proportional to its absolute
temperature:
V ~ T V1 / T 1 = V2 / T 2
For each degree increase in temperature, the volume
increases 1/273 of its original volume
4. Combined gas law:
P1 V1 / T1 = P2 V2 / T2
5. Ideal Gas Law:
PV = nRT
Overall conclusions:
The temperature of a gas increases when it is compressed
because the average energy of its molecules
increases. The molecules rebound from the inward
moving piston, traveling faster than before hitting the
piston.

Molecules rebounding from fixed walls have unchanged
speeds.
The temperature of a gas decreases when its volume is
expanded because the average energy of its molecules
decreases. The molecules rebounding from outward
moving piston move slower than before.

Gas Law Problems:
1) An insulated system is known to have a temperature of
100.0° C at a pressure of 4.00 atm. If the absolute
temperature is cut in half, what will be the new:
___________ atm, __________kPa, ______________° C,
_______________ K
2) The volume is given as 27.0 L. If the pressure goes
from 3.00 atm. to 9.00 atm., what is the new:
___________L, _____________ kPa
3) The volume is given as 5.00 L. If the absolute
temperature goes from 273 to 819 K, what is the the new
__________L (if new temperature was 800. K, what is the
new volume in liters?)
4) The temperature is given as 25.0° C. If the volume is
decreased from 100. mL to 10.0 mL, what is the
new: ____________K, ____________° C

Thought Work -- Heat & Gas Laws
1. Devise a way to remove carbon dioxide (carbonation)
from soda pop. This must be done quantitatively. How does
temperature affect the solubility of gases being dissolved in
liquids under pressure?
2. Explain what happens to a marshmallow when it is
toasted. Why does this happen? What would happen to a
frozen marshmallow? Why does a marshmallow float?
3. Find the pressure you exert when standing on both feet,
on one foot, and lying flat on your back.
4. Explain why a toy balloon filled with hydrogen
partially deflates overnight.
5. Using a steel ball and pieces of old pottery or modeling
clay, devise an experiment that would demonstrate
potential energy, kinetic energy, and momentum (all of
which are involved with mass and velocity).
6. Suppose you had two identical sections of glass plate
before you, one heated above body temperature and the
other cooled by ice. What happens when you breathe on the
two of them and why? Maybe try this at home first.
7. Devise an experiment to show the concept of diffusion,
another to show cohesion, adhesion, and surface tension,
and one to show buoyancy and Pascal's Law.
8. Changing ice to water requires 80 calories per gram of
ice, but changing water to steam requires about 540 calories

per gram of water. What does this tell you about the
intermolecular forces in ice and water, both qualitatively
and quantitatively? Also explain why the chemical change
of splitting or forming water requires about 5 times as
many calories as the physical change of state.
Chemical Properties of Matter
Chemical properties are those properties of a substance
that can be determined by a chemical test. Chemical
properties are seen by the material's tendency to change,
either alone or by interaction with other substances, and in
doing so form different materials.
1.does the substance support combustion: examples are
O2 and Cl2
2. does the substance burn (combustibility)
3. how does the substance react with acids (does it dissolve,
evolve gases, explode, do nothing)
4. how does the substance react with oxygen (burn, form
new compounds)
5. what is its reaction with electricity (usually it will be
separated into simpler components)
examples: alcohol burns, iron rust, wood decays, sodium
explodes in water

Physical Properties of matter

Physical properties are those properties used in identifying
substances when we use our senses. These do not require
chemical analysis.
1. color - reaction of eye and brain in recognizing
combinations of certain wavelengths of visible light.
2. hardness - a measure of the ability of a substance to
resist abrasion (see Mows Scale of Hardness)
3. density - the mass divided by its volume (often
reported as specific gravity which is a unitless relationship
between the density of the substance and the density of
water)
4. texture - how object feels to touch; usually rough or
smooth
5. magnetic attraction - is the material attracted to a
magnet or can it be magnetized (must contain Fe, Co. Ni, or
steel)
6. solubility - the amount of a substance which will
dissolve in a known amount of solvent at a given
temperature
7. taste - reaction of taste buds to stimuli along with the
brain's recognition of the pattern
8. light transmission - is the substance transparent,
translucent, or opaque

9. viscosity - a measure of the internal resistance
(friction) to flow in a liquid (molasses and tar would have
high viscosities)
10. refractive index amount a ray of light is bent as it
passes through a substance (technically the ratio of the
speed of light in that substance to the speed of light in a
vacuum)
11. specific heat capacity - the amount of heat energy
(calories or Joules) required to change the temperature of 1
gram of a substance by 1 °C
12. atomic radius - the distance from the center of an
atom's nucleus to the outermost orbital electron
13. boiling point - the temperature at which the liquid's
vapor pressure equals atmospheric pressure during the
boiling of a pure substance the temperature remains
constant as long as both liquid and vapor are present
14. melting-freezing point - temperature at which solid-
liquid phase is in equilibrium - during melting of pure solid
the temperature remains constant; when all solid is melted
and only liquid is present, further heating results in a steady
increase in temperature to the boiling point
15. odor - olfactory nerves are stimulated by certain
molecule and send messages to the brain which remembers
the pattern
16. expansion - contraction coefficients - materials
expand or contract a known amount when heated or cooled

Collapsing Can Demo
area of surface of Coke can = 0.031 m2
pressure on this area = 3.1 E 3 N (about 680 lb)
1 atm = 1.0 E 5 N/m2

if we can reduce pressure by as little as 75% there would be
a 500 lb difference between pressure inside and outside the
can
1) when can is inverted in water bath - the water
seals the opening and cools the can
2) as can cools vapor condenses - reduction of
pressure inside can
3) can is sufficiently weak and water sufficiently
viscous that can collapses before it fills with water
4) must use all aluminum can

Physical Changes in State
a change in the physical properties of a substance without
a change in the chemical composition
the arrangement of molecules may be changed but the
molecular make-up remains the same

these changes deal with intermolecular forces which
increase or decrease during the change.
Problems:
1. 72 grams of ice + 51 840 calories yield 72 grams of
water vapor. How many calories must be removed from
water vapor to condense it back to ice?
2. If 1 gram of water at room temperature evaporates,
about 600 calories are taken from the surroundings to
convert the liquid to a gas. How many calories are `needed'
to change 1001 grams of gas to a liquid?
3. If 50 grams of water vapor loses 36 000 calories in
turning to ice, how many calories would 1 gram of water
vapor have to lose to be turned back to ice?

ice (0° C) + heat à water vapor (100° C)
36 g 25 920 cal 36 g
water vapor (100° C) à ice (0° C) + heat
36 g 36 g 25 920 cal
2 H2 + O2 à 2H2O + heat energy
released
4g 32 g 36 g 136 600 cal
2H2O + energy à 2H2 + O2

36 g 136 600 cal 4g 32 g

Chemical Changes in State (Phase)
The molecular make-up (the specific arrangement of
atoms) is changed, resulting in new substances being
formed and energy changes occurring.
EXOTHERMIC - any chemical change that releases energy
is exothermic
the amount of heat released is greater than the amount
of heat used to start the reaction
bond making is exothermic (energy is released into
surroundings)
example: oxidation à wooden splint burning ( heat, light,
gases like CO2 and H2O being given off with carbon and
ashes left over)
other examples: burning H2 in O2, body reactions,
dissolving metals in strong acids, mixing acid and water,
homogenization, plaster of Paris in water, sugar
dehydration
ENDOTHERMIC - any chemical change that absorbs
energy is endothermic
energy continues to be absorbed as long as the
reaction continues
bond breaking is endothermic (energy is absorbed
from surroundings)
example: electrolysis à splitting some compound

(usually water) by running an electric current through it
other examples: photosynthesis, pasteurization, canning
vegetables
Sugar dehydration demo here
The chemical change involving splitting or forming
water takes about 5 times as many calories as the physical
change of state. The reason is that atoms (or molecules) are
bonded together in a compound; the stronger the bond the
more energy holding the parts together, thus more energy
required to break these bonds. A physical change needs far
less energy to overcome intermolecular forces holding
groups of molecules together. Much more energy is needed
to break bonds within molecules than to overcome the
forces between molecules.
physical change -- strength of intermolecular forces
increased or decreased
chemical change -- bonds formed or broken
energy absorbed -- bonds broken or intermolecular forces
overcome
energy released -- bonds formed or intermolecular forces
strengthened
Problems:
Tell whether each of the following is a chemical or physical
change and further describe each chemical change as

endothermic or exothermic and the physical changes as
absorbing or releasing energy.
dry ice sublimates
CO2 + H2O + sunlight à glucose
air in heated tire expands
burning coal
water frozen into ice
acid dissolves metal

Endothermic vs Exothermic Reactions
All chemical reactions involve bond breaking and bond
making.
Bond breaking is endothermic (energy is absorbed from
surroundings)
Bond making is exothermic (energy is released into
surroundings)
Imagine stretching a rubber band until it breaks. You must
do work to stretch the band because the tension in the band
opposes your efforts. You lose energy; the band gains
it. Something similar happens when bonds break in a
chemical reaction. The energy required to break the bonds
is absorbed from the surroundings.
Energy is absorbed or released when the heat capacities of
the products and reactants differ. Usually this is

small. Remember that heat capacity is best thought of with
a penny and specific heat best thought of as copper metal.
Neutralization reactions are usually exothermic but when
you add baking soda to vinegar it is slightly
endothermic. The neutralization reaction actually does
release heat:
HC2H3O2 + NaHCO3 à CO2 + NaC2H3O2 (aq) + H2O
This is because there is net bond formation. The products
collectively have lower energy than the reactants. But
evaporation of the liquid occurs as the carbon dioxide
escapes from solution. Evaporation absorbs heat, cooling
the liquid. (The expansion of the carbon dioxide gas
bubbles as they are released also helps to cool the
surroundings by Joule-Thomson cooling). The net result is
an endothermic reaction.
Mixing a strong acid with water is exothermic. Breaking a
chemical bond requires energy (remember that stretching a
spring until it breaks requires energy). Forming a chemical
bond will release energy. So in a reaction that releases heat
(exothermic) there must be net bond formation. Lets looks
at HCl dissolved in water:
HCl à H+ (aq) + Cl1- (aq)
You would think at first this would be a heat absorbing
(endothermic) process, because it looks like the bond
between H and Cl is broken. But there is another reaction
hiding here. The hydrogen ion reacts with water to form a
complex of the form: H3O·(H2O)+n where n is a number

between 1 and 9. It is much easier just to write H+
(aq). Because the hydrogen ion is so tiny, a large amount
of charge is concentrated in a very small area, and the polar
water molecules are strongly attracted to it. This
"hydration" of the hydrogen ion involves the formation of a
covalent bond to one of the waters and a large number of
strong hydrogen bonds, so it’s a strongly exothermic
process. This causes the mixing of a strong acid with water
to be strongly exothermic overall.
Exothermic processes Endothermic processes
making ice cubes melting ice cubes
formation of snow in clouds conversion of frost to water
vapor
condensation of rain from evaporation of water
water vapor
a candle flame forming a cation from an atom
in the gas phase
mixing sodium sulfite and baking bread
bleach
rusting iron cooking an egg
burning sugar producing sugar by
photosynthesis
forming ion pairs separating ion pairs
combining atoms to make a splitting a gas molecule apart
gas molecule
mixing strong acids and mixing water and ammonium
water nitrate
nuclear fission melting solid salts

Heating Curve Information: This graph will aid in
understanding the following information.
Phase Change Diagram
Solid - Gas Phase Change
This change involves sublimation which is the direct
change of a solid to a gas (deposition is the opposite).
Examples include: moth balls (naphthalene),
paradichlorobenzene, camphor, iodine crystals, and
CO2 fire extinguishers (advantages: does not conduct
electricity, colder than water, replaces O2 since CO2 is
heavier and settles on ground area and the CO2 does not
combust), will sublime away reducing cleanup -
disadvantages include difficulty in keeping container
pressurized over time, fact that you cannot use on living
things due to extreme cold, and cost).

Liquid - Solid Phase Change
This discussion deals with melting-freezing point. A
complete discussion of this concept using ice and heat units
will be completed in class.
See class discussion of ice cube. The addition of 1 calorie
of heat to the ice cube at 0° C does not cause a change in
the temperature of the ice cube though 1 calorie would
change the temperature of 1 gram of water at 0° C.

It will take 80 calories just to melt the ice cube. That heat
that is consumed in melting the solid is converted into
potential energy. Freely moving molecules in liquids, with
respect to intermolecular attraction, possess more energy
than similar molecules bound rigidly in solids at the same
temperature.
Remember that temperature is a measure of the average
kinetic energy only while heat content is a measure of the
total kinetic energy plus potential energy possessed by that
body.
See class examples of the heat energy needed to change
ice at any temperature to steam at any temperature.
The melting-freezing point is defined as the temperature at
which the solid and liquid phases are in equilibrium. This is
the temperature at which a change of state between the
solid and liquid phase can occur. Some of the solid will be
melting and some of the liquid will be freezing.
When a solid is heated to its melting point, its atoms or
molecules acquire enough energy to shift the bonds holding
them together so they form separate clusters. This
clustering in liquids is confirmed with X-ray studies but the
clusters are constantly shifting their arrangement unlike the
permanent arrangement in solids.
When heat is added to a solid the temperature of the solid
will increase till it reaches the melting-freezing point. It
will remain there until all the solid has melted and only

then can the temperature of the liquid rise according to its
specific heat.
Water molecules at 0° C contain more energy than the ice
molecules at 0° C , not in the form of a faster more rapid
motion but in the form of an ability to resist the attractive
forces tending to pull them together.
Melting points also depend on pressure (though not as
much as boiling points.) Ice is strange in that its melting
point decreases with increasing pressure. Almost all other
materials show increasing melting points with greater
pressures. The pressure an ice skater exerts on the ice due
to the small area of the skate blade is usually enough to
melt the ice creating a thin film of water that acts as a
lubricant. On unusually cold days the pressure may not be
enough to melt the ice and thus skating would be
impossible.
Liquid - Gas Phase Change
The change from a gas to la liquid is condensation. This is
due to cooling and/or a pressure change.
In liquids, the energy of the particles is raised by adding
heat. When some molecules have enough K.E. they break
away from the liquid surface and become vapor.
If the temperature falls, there is a decrease in the energy of
the moving molecules and the liquid may eventually freeze
to the solid phase.

Process of EVAPORATION: Molecules that have enough
energy of motion (K.E.) break free from intermolecular
forces and escape into the air as vapor. Some may return to
the liquid is their energy is lost to other atoms.
The liquid surface left behind is cooled. In evaporation the
molecules that escape are the ones with the greatest
velocity (heat) thus the average velocity and K.E. of the
remaining particles is reduced. This results in cooling
effects. Heat must be absorbed from the surroundings to
continue the evaporation process.
Adding heat increases evaporation because the VAPOR
PRESSURE is increased. This is the pressure exerted by
the vapor (gas) of a substance when it is in equilibrium
with liquid or solid phase. The system is in equilibrium
when the rate of evaporation equals the rate of
condensation.
The temperature at which the liquid's vapor pressure is
equal to outside (atmospheric) pressure is that
liquid's BOILING POINT. At this temperature the
pressure of the vapor escaping from liquid equals the
outside pressure.
When the vapor pressure equals outside pressure bubbles
of vapor form and push through to the surface. As they
move into the gas phase we say this is boiling. Conduction
of heat creates the gas, which rises because it is less dense
than the liquid, as it strives for equilibrium.

The boiling point varies with atmospheric pressure. In
mountains, the boiling point is below 100°C because the
pressure of the atmosphere is less.
Cooking requires longer times at high altitudes because of
low boiling point
Pressure cookers make food cook more rapidly because the
foods can be heated above the normal boiling point without
actually boiling.
Intermolecular Forces and Latent Heat
if we heat a mixture of ice and water, we find that no
matter how much heat is transferred to the mixture, the
temperature remains at 0° C until the last of the ice is
melted. Only after all the ice is melted is heat converted
into kinetic energy, and only then can the temperature of
the water begin to rise. Experiment shows that 80 calories
of heat must be absorbed from the outside world in order
that 1 gram of ice might be melted, and that no temperature
rise takes place in the process. The ice at 0° C is changed to
water at 0° C.
But if the heat gained by the ice is not converted into
molecular kinetic energy, what does happen to it? If the
Law of Conservation of Energy is valid, we know it cannot
simply disappear.
The water molecules in ice are bound together by strong
attractive forces that keep the substance a rigid solid. In
order to convert the ice to liquid water (in which the
molecules, as in all liquids, are free of mutual bonds to the

extent of being able to slip and slide over, under, and
beside each other) those forces must be countered. As the
ice melts, the energy of heat is consumed in countering
those intermolecular forces. The water molecules contain
more energy than the ice molecules at the same
temperature, not in the form of a more rapid motion or
vibration but in the form of an ability to resist the attractive
forces tending to pull them rigidly together.
The Law of Conservation of Energy requires that the
energy change in freezing be the reverse of the energy
change in melting. If liquid water at 0° C is allowed to lose
heat to the outside world, the capacity to resist the
attractive forces is lost, little by little. More and more of the
molecules lock rigidly into place, and the water freezes.
The amount of heat lost to the outside world in this process
of freezing is 80 calories for each gram of ice formed.
In short, 1 gram of ice at 0° C, absorbing 80 calories,
melts to 1 gram of water at 0° C; and 1 gram of water at 0°
C giving off 80 calories, freezes to 1 gram of ice at 0° C.
The heat consumed in melting ice or any solid, is
converted into a sort of potential energy of molecules. Just
as a rock at the top of a cliff has, by virtue of its position
with respect to gravitational attraction more energy than a
similar rock at the bottom of the cliff, so do freely moving
molecules in liquids, by virtue of their position with respect
to intermolecular attraction, possess more energy than
similar molecules bound rigidly in solids.

It is the kinetic and potential energies of the molecules that
together make up the internal energy that represents the
heat content. It is kinetic energy only that is measured by
the temperature. By changing the potential energy only, as
in melting or freezing, the total heat content is changed
without changing the temperature.
In converting a gram of liquid water at 100° C to a gram of
steam at 100° C what remains of the intermolecular
attractions must be completely neutralized. Only then are
molecules capable of displaying the typical properties of
gases -- that is, virtually independent motion. In the earlier
process of melting, only a minor portion of the
intermolecular attractive force was countered, and the
major portion remains to be dealt with. The latent heat of
vaporization of water (the amount of heat required to
convert 1 gram of water at 100° C to 1 gram of steam at
100° C) is 540 calories, almost seven times the earlier 80
calories needed in changing ice to water.
The energy content of steam is thus surprisingly high. 100
grams of water at 100° C can be made to yield 10 000
calories as it cools to the freezing point. 100 grams of
steam at 100° C can be made to give up 54 000 calories
merely by condensing it to water. The water produced can
then give up another 10 000 calories if it is cooled to the
freezing point. It is for this reason that steam engines are so
useful and hot water engines would never do as a
substitute.
If we boil water in a kettle its temperature remains at 100°
C, no matter how fast we boil it, but we have to keep

adding heat to keep it boiling. Heat is absorbed by the
molecules as they escape their liquid state and become a
gas. The amount of heat needed to pull apart liquid
molecules is called heat of vaporization (calories/gram).
The heat of vaporization which a molecule must absorb
before it can become a gas molecule is released by it when
it cools again to liquid, as heat of condensation. Liquids
with low boiling points, such as alcohol or ether, chill the
hand as the molecules pick up their heat of vaporization
and become a gas. The same is true for a glass of water, it
will be cooler than room temperature.
The kinetic molecular theory states that the kinetic energy
depends on heat energy, which can be measured as
temperature. A thermometer in boiling water and a
thermometer in the vapor just above the boiling surface will
read the same; 100° C at sea level. Therefore the average
kinetic energy of the liquid molecules must be the same as
the average kinetic energy of the gas molecules above it.
An average molecule in the liquid state will be moving as
fast as an average molecule in the gaseous state.
Gas particles move in a straight line until they collide with
another bit of matter, then they bound away in some other
direction but always in a straight line, and without losing
any of their energy to friction in the collision. The particles
have perfect resilience. However, they will change their
kinetic energies in the familiar way of all normal matter, as,
for instance, do billiard balls. A slow-moving particle hit
from behind by a fast one is speeded up, while the fast one
is slowed down, but the sum total of their kinetic energies

remains the same. In the world of normal matter, perfect
elasticity is unknown, as there is friction between surfaces.
Two billiard balls when they collide will change each
other's speed and direction of motion, and they will also
roll to a stop in a short time. The ultimate particles of
matter lose not a bit of their energies in collisions. They
simply exchange speeds. If two particles collide, their total
heat before and after is the same, but the originally slower
particle after the collision is traveling faster and is therefore
hotter, while the formerly speedier particle is now cooler
and moving more slowly that it was. Heat and molecular
motion, according to the theory, are two ways of speaking
about the same thing.

Heat/Temperature Problems:
1. Suppose a piece of iron (mass = 21.50 g at a
temperature of 100.° C) is dropped into an insulated
container of water (mass of water = 132 g and the
temperature before adding iron was 20.0° C). What will be
the final temperature of the system (at thermal
equilibrium)? The specific heat of iron is 0.113 cal/g° C.
2. If 200. grams of water is to be heated from 24.0° C to
100.0° C to make a cup of tea, how much heat must be
added?
3. Which is more effective in cooling a drink, 10 grams of
water at 0° C or 10 grams of ice at 0° C? Explain your
answer quantitatively.

4. A 3.00 kg lead bar at 100.0° C is placed in 4.00 kg of
water at 20.0° C. The final temperature of the lead bar
would be ___________. (cp of lead is 0.0305 cal/g° C)
5. A 0.60 kg copper kettle holds 1.70 kg of water at
30.0° C. A 0.10 kg iron ball at 120.0° C is dropped into the
water. What is the final temperature of the water? (cp of
copper = 0.377 J/g° C and iron is 0.448 J/g° C)
6. A piece of iron with a mass of 20.50 grams at a
temperature of 100.0° C is dropped into 140.00 grams of
water at 40.0° C. What will be the final temperature of the
system. The cp of iron is 0.45 J/g° C)
7. A sample of mercury metal is heated from 25.5° C to
52.5° C. In the process, 187 cal of heat are absorbed. What
mass of mercury was in the sample? The specific heat of
mercury is 0.033 cal/ g° C
8. A block of aluminum weighing 140. g is cooled from
98.4° C to 62.2° C with the release of 1080 cal of heat.
From these data, calculate the specific heat of aluminum.
9. A cube of gold weighing 192.4 g is heated from 30.0°
C to some higher temperature, with the absorption of 226
cal of heat. The specific heat of gold is 0.030 cal/ g
° C. What was the final temperature of the gold?
10. A total of 54.0 cal of heat are absorbed as 58.3 g of
lead is heated from 12.0° C to 42.0° C. From these data,
what is the specific heat of lead?

11. A piece of erbium metal weighing 100.0 g and heated
to 95.0° C is dropped into 200.0 g of water initially at 20.0°
C. The final temperature of the mixture is 21.5° C. What is
the specific heat of erbium metal?
12. A block of rhenium metal (specific heat = 0.0329 cal/
g° C) is heated to 88.2° C and then dropped into 100.0 g of
water initially at 26.4° C. The final temperature of the
mixture is 32.4° C. What was the mass of the block of
rhenium?
13. When 258.6 g of benzene vapor is condensed to a
liquid at its boiling point, 33 875 cal of heat are released.
What is the heat of vaporization for benzene?
14. A sample of ethyl alcohol is converted from a liquid to
a vapor with no temperature change. In the process 30 640
cal of heat are absorbed. What mass of ethyl alcohol was in
the sample? The heat of vaporization of ethyl alcohol is
210. cal/g.
15. The heat of combustion of methane is 212.8 kcal per
mole. How much heat will be produced in the combustion
of 100.0 g of methane?
16. The heat of combustion of toluene is 934.2 kcal per
mole. How much heat will be released during the
combustion of 250.0 g of toluene? The formula for toluene
is C6H5CH3
17. Copper has a density of 8.94 g/cm3 and a specific heat
of 0.090 cal/g° C. A cube of copper is heated from 10.5° C

to 214° C. The cube of copper has dimensions of 5.00
cm. How much heat would the copper cube absorb?
18. The specific heat of water is 4.185 J/g° C (1.00 cal/g
° C). A piece of a pure metal with a mass of 24.0 g at a
temperature of 45.0° C is added to 55 mL of water at
60.0° C. The final equilibrium temperature of the mixture
is 95.4° C. Find the specific heat of the pure metal in both
cal/g° C and J/g° C.
19. The specific heat of ice is 2.03 J/g° C. How much
heat is needed to convert 550. g of ice at –15.0° C to
10.0° C?
20. What is the total amount of heat needed (in calories
and joules) to convert 2.25 kg of ice at 0.0° C to steam at
200.0° C.
21. The specific heat of silicon is 0.057 cal/g° C and the
density of silicon is 4.4 g/cm3. The volume of a cylinder
formula is given as p r2L. The addition of 6000. calories
raises the temperature of the silicon cylinder 55.5° C. Find
the radius of the cylinder. The length of the cylinder is
given as 6.00 cm.
22. A piece of metal with a mass of 75.5 g is heated to
84.5° C and added to 100.0 mL of water at 5.0° C. The
final temperature of the mixture is 75.0° C. Find the
specific heat of the metal.
23. Granite has a specific heat of 800. J/g° C. What mass
of granite is needed to store 1.50 E 6 J of heat if the
temperature of the granite is to be increased by 15.5° C?

24. A 55 kg block of granite has an original temperature
of 15.0° C. What will be the final temperature of this
granite if 4.5 E 4 kJ of heat energy are added to the granite?

The above curve attempts to demonstrate the addition of
100 calories of heat energy per minute to 1 gram of water.
A thorough class discussion will attempt to identify the
amount of time needed for each change to occur.

Knowledge of specific heat capacity, latent heat of fusion
and vaporization is needed to determine these values.

The above chart shows a graph of the contraction of an air
bubble in a capillary tube filled with oil. As the tube was
cooled the length of the air bubble was measured and
plotted as dots. When it could no longer be cooled to a
lower temperature the graph line was extrapolated to find
Absolute Zero (a temperature unobtainable in the actual
world).

Short Answer Questions
Directions: Answer each question fully. Use complete
sentences. Skip two lines of notebook paper (double
double space if typed) between each question. Use only the
front of paper.

1. Explain latent heat (use water as an example). Give
quantitative examples and clearly explain how
intermolecular forces are involved during these
changes.
2. Discuss how specific heat capacity might be used to
identify an unknown metal sample.
3. Compare and contrast endothermic and exothermic
chemical reactions.
4. Compare and contrast kinetic and potential
energy, using as many practical examples as possible.
5. Using examples, explain fully elastic and inelastic
collisions. Try to include some of the problems these
concepts cause in science
classes. Includemomentum in your discussion.
6. Describe how a liquid thermometer is made and how
it could be calibrated without another thermometer
7. Explain how chemical and physical properties might
be used to identify an unknown substance. Pick one
common substance and give physical/chemical details
for it.
8. Discuss the way in which gas exerts pressure on the
walls of its container.
9. Explain the Kinetic-Molecular Theory of Gases as it
relates to ordinary life. Include examples of Boyle’s
Law and Charles’ Law in your explanation.
10. Compare intermolecular forces with chemical
bonds within molecules.
11. Discuss atmospheric pressure and how it relates to
gas laws, boiling points, and vapor pressures.
12. Explain how conduction, convection, and
radiation might occur when using our

calorimeters. Compare our calorimeters with good
Thermos bottles.
13. Explain fully melting-freezing point theory.
14. Explain fully boiling point theory and related
phenomena such as elevation changes, pressure
cookers, super heating, and vapor pressure.
Molecular Motion Demonstrator
This demonstration tool allows us to model a variety of
atomic behavior. The following are notes (these will save
the student having to make them during the demonstration
and will allow them to study the modeling better.
General observations
1) some molecules move faster than others, with
constant random motion
2) molecules collide with each other and the walls
3) near elastic collisions
4) rarity of 3 body collisions
5) generally there is a large amount of space between
the molecules
6) random motion in straight lines between collisions
7) pressure exerted on whatever they hit
Diffusion - as the particles move across the
barrier they demonstrate diffusion: (movement of
molecules from one area to another with a net change in
concentration)
Temperature

1) related to average speed of molecules
2) would 2 different gases at same temperature have
same average speeds?
3) observe different speeds of different gas particles -
> temperature relates to average K.E. of molecules
4) increase vibration rate: average speed increases,
frequency of collisions increases, mean free path
decreases
Similarities with real world
1) model and real gases involve particles in
continual, random motion, with straight line paths
between collisions
2) particles in model occupy only a small fraction of
total volume
3) changes in amount of movement is associated with
change in temperature
4) rebounding forces of particles produce pressure in
both model and real gases
Dissimilarities with real world
1) model uses particles approximately ten million
times the diameter of the particles in real gases
2) distance between collisions in real gases much
greater than the distance traveled by plastic balls in
relation to size
3) the speeds of the balls in the model are at most a
few miles per hour while gas molecules travel at speeds
of hundreds of miles per hour
4) real gas molecules collide elastically while our

model will run down due to friction is energy is not
applied
5) real gases are three dimensional, model is two
dimensional
6) most gas molecules are not spherical
7) spaces between particles in model taken up by air,
in gases empty spaces occupies the space between
molecules
Liquids and Solids
attractive forces between gas molecules are
called van der Waal forces - positive charge of nucleus
is not completely shielded by electron cloud so at
short distances the nucleus may be attracted to
electron cloud of another atom - these gas molecules
may come together, slow down, and allow attractive
forces to form
Boyle’s Law
Charles’ Law

Relative Humidity
When molecules from the vapor above a
liquid surface impinge on the surface, they may be trapped
there, so that a constant two-way traffic of molecules to and
from the liquid occurs. If the density of the vapor above the
liquid is sufficiently great, as many molecules return as
leave it at any time, a situation that is described by saying

that the region is saturated with the substance. The higher
the temperature, the greater the maximum vapor density: at
0° C the density of water vapor at saturation is 5 g/m3, at
20° C it is 17 g/m3, at 100° C it is 598 g/m3, and at 300° C
it is all the way up to 45.6 kg/m3.
If for any reason (such as a sudden drop in
temperature) the vapor density exceeds the saturation value,
condensation will be more rapid than evaporation until
equilibrium is reestablished. It is for this reason that on a
hot day moisture condenses on the outside of a glass that
contains a cold drink. The relative humidity of a volume
of air describes its degree of saturation with water vapor.
Relative humidities of 0, 50%, and 100% mean respectively
that no water vapor is present, that the air contains half as
much moisture as the maximum possible, and that the air is
saturated. On a hot day the evaporation of sweat from the
skin is the chief means by which the human body dissipates
heat, and a high relative humidity is uncomfortable because
it impedes the process. A low relative humidity is also
undesirable because it leads to the drying of the skin and
mucous membranes. The regulation of relative humidity is
as important a function of a heating or of an air-
conditioning system as the regulation of temperature.
We know that heating air decreases its relative
humidity and cooling air increases its relative humidity. For
instance, between 10° C and 20° C the saturated vapor
density (which corresponds to 100% relative humidity) just
about doubles. This means that if outside air at 10° C
whose relative humidity is say, 70% is taken inside a house

and heated to 20° C, the relative humidity indoors will only
be 35% since the actual vapor density stays the same. A
way to humidify heated air in winter is clearly desirable. If
the outside air is at 30° C with 70% relative humidity, then
cooling it down to 24° C is enough to bring it to saturation,
which is 100% relative humidity; further cooling will cause
water to condense out. An air-conditioning system
therefore should incorporate means to remove water vapor
from the air being cooled.
Relative humidity is water vapor density relative to
saturation density. Changing the temperature of a body of
air also changes its relative humiditiy.
Principle of Heat Exchange - The Coffee-Cream
Problem
Newton's law of cooling: The rate of heat conduction is
proportional to the temperature difference between an
object and its surroundings.
The Stefan-Boltzmann law of radiation: The rate of heat
lost by radiation is proportional to the fourth power of the
absolute temperature.
The historic problem:
Ah, you see, there is this business man who likes a
large amount of cream in his coffee, and he wants the
resultant mixture as hot as possible. (Alas, there is no
microwave oven available).

He has just prepared his boiling coffee when he is
called by the boss for a quick conference of ten minutes
duration. The boss tolerates no coffee in his presence.
What to do? To keep the coffee as hot as possible
should he add the cream now or wait until after the
conference?
Which do you think he should do?
_____________________________________
Using beakers for the coffee cups and water for the coffee
and cream, design an experiment to test the problem. Use
the computer and temperature probe to measure the
temperature over time. How would you set up a graph(s)
to help prove the best solution.
Formulas
K.E. = ½ m v2
P.E. = m g h
Q = mass x D t x cp
Q = mass x Hf
Q = mass x Hv
K = ° C + 273
° C = K - 273
1 calorie = 4.185 joules

Pressure = Force
area

Temperature Scale Comparisons

Fahrenheit Celsius Kelvin
Boiling p. of
212° 100° 373
water
Body
98.6° 37° 310
Temperature
Freezing p.
32° 0° 273
water
Coincidence
- 40° - 40° 233
p.
Absolute -273.16
- 460° 0
Zero °

Heat of Solution Reactions Lab
Chemical and physical changes are usually accompanied
by the liberation or absorption of energy. If energy is
evolved, the reaction is said to be EXOTHERMIC. If the
energy is absorbed, the reaction is ENDOTHERMIC.

Heat is a form of energy. The calorie or joule is the unit
used to express heats of reaction. The calorie is defined as
the amount of heat required to raise the temperature of 1
gram of water one Celsius degree. For conversions, 1
calorie = 4.185 joules.
Energy may be transformed from one kind to another
within an isolated (or closed) system but the total energy
does not change. If the change in energy of a system can
be measured and if this change is due solely to a chemical
reaction, then the energy change must be equal to that of
the chemical reaction itself.
In this experiment a simple calorimeter will be used and
the change in energy of the system will be measured by
observing the temperature of a given weight of water
before and after the reaction occurs. The specific heat
capacity of water (i.e., the energy required to raise the
temperature 1o C of 1 gram of the material) is very nearly
1.00 cal/goC for temperatures between 0o C and 100o C.
Thus, if a calorimeter contained 100 grams of water at
23.0o C initially and after the reaction took place there were
still 100. grams of water and the temperature was now
30.0o C, the energy liberated in the reaction would be
Q (cal) = mass (g) x Dt (oC) x cp (cal/goC)
700. cal = 100. g x 7.00o C x 1.00 cal/goC
This calculation assumes that no energy was required to
raise the temperature of the calorimeter itself and that no

The nature of heat

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