•Heat and Temperature
Heat and modern technology are inseparable. These
glowing steel slabs, at over 1,100OC (about 2,000OF),
are cut by an automatic flame torch. The slab caster
converts 300 tons of molten steel into slabs in about
45 minutes. The slabs are converted to sheet steel for
use in the automotive, appliance and building
• The Kinetic Molecular Theory.
• Introduction.
– Ancient Greeks knew that matter was made up of
very small particles.
– Democritus wrote that matter was made up of tiny
indivisible particles he called atoms.
• We now know that atoms are not indivisible,
but are themselves made up of even smaller
particles.
• We have identified more that 200 smaller
particles that make up atoms.
• Molecules.
– Basic assumption is that matter is made up of tiny
units of structure called atoms.
– Atoms are neither created or destroyed during any
type of chemical or physical change.
– Arrangements of atoms determines type of entity
of matter.
• Elements are pure substances made up of only
one type of atom.
• Compounds are made up of one type of atom,
but have more complex structures.
• Pure substances are composed of 2 or more
elements in defined proportions.
• A molecule is the smallest particle of a
compound in which all of the atoms maintain
their identity.
–maintains all of the chemical and physical
properties of the compound.
–Some atoms naturally form molecules called
diatomic molecules: O2, F2, Cl2, N2, Br2,
Metal atoms appear in the micrograph of a crystal of
titanium niobium oxide, magnified 7,800,000 times by
an electron microscope.
• Molecules Interact.
– Cohesion.
• Some solids and liquids attract each other and
cling to each other.
• Cohesion is when this attractive force is
between like molecules.
– Adhesion.
• Some molecules are attracted to other
molecules.
• Phases of Matter.
– Solids.
• Defined shapes.
• Defines volumes.
• Molecules are fixed distances apart and have
strong cohesive forces.
– Liquids.
• Close together.
• Cohesive forces not as strong as in a solid.
• Defined volume, but not a defined shape.
– Gases
• Weak cohesive forces.
• High kinetic energy.
• Molecules far apart and move in random
motion
• No fixed shape or volume.
• Vapor is a gas that is above a liquid phase.
(A) In a solid, molecules vibrate around a fixed
equilibrium position and are held in place by strong
molecular forces. (B) In a liquid, molecules can rotate
and roll over each other because the molecular forces
are not as strong. (C) In a gas, the molecules move
rapidly in random free paths.
• Molecules Move.
– All molecules have kinetic energy due to
movements.
– This kinetic energy can be in the form of:
• Vibrational energy.
• Rotational energy.
• Translational energy where the entire molecule
has motion.
The basic forms of kinetic energy of molecules. (A)
Translational motion is the motion of a molecule as a
whole moving from place to place. (B) Rotational
motion is the motion of a turning molecule. (C)
Vibrational motion is the back-and-forth movement of
a vibrating molecule.
– The kinetic energy of a substance is
measured as the temperature of that
substance.
• Temperature is actually a measure of the
average kinetic energy and has nothing to do
with heat until there is a transfer of energy.
The number of oxygen molecules with certain
velocities that you might find a sample of air at room
at temperature. Notice that a few are barely moving
and some have velocities over 1,000 m/s at a given
time, but the average velocity is somewhere around
500 m/s.
• Temperature.
• Thermometers.
– Conceptually a thermometer is used to measure
the hotness or coldness of an object.
– What a thermometer really measures is the
average kinetic energy of an object.
– There is a physical transfer of kinetic energy to
the thermometer which them responds due to the
increase in its kinetic energy.
• Mercury.
• Ethylene glycol.
(A) A bimetallic strip is
two different metals, such
as iron and brass, bonded
together as a single unit,
shown here at room
temperatures. (B) Since
one metal expands more
than the other, the strip
will bend when it is
heated. In this example,
the brass expands more
than the iron, so the
bimetallic strip bends
away from the brass.
This thermostat has a coiled bimetallic strip that
expands and contracts with changes in the room
temperature. The attached vial of mercury is tilted one
way or the other, and the mercury completes or breaks
an electric circuit that turns the heating or cooling
system on or off.
• Thermometer Scales.
– Fahrenheit scale
• Sets boiling point of water at 212 OF and
freezing point of water at 32 OF.
• 180 divisions between these two.
• Like most English measures is quite
cumbersome.
– Celsius scale.
• Sets boiling point of water at 100 OC and
freezing point of water at 0 OC.
• 100 divisions between these two points.
– Kelvin or absolute scale.
• Begins at absolute zero, the temperature at
which all kinetic energy is changed into
potential energy.
–ie, all molecular motion ceases.
• Boiling point of water is 373 K and freezing
point of water is 273 K.
• Divisions are same as for Celsius scale
The Fahrenheit, Celsius, and Kelvin temperature
scales
• Conversions.
– TF = 1.8 TC + 32 OC
– TC = (TF - 32 OF)
•
1.8
• 1.8 accounts for the divisions between freezing
point of water and boiling point of water.
–There are 1.8 divisions in the F scale for
every 1 division in the C scale.
– TK = TC + 273.
– Example.
• The temperature of Lake Superior in August
averages 34 OF. What is the temperature in OC.
• Use: TC = (TF - 32 OF)
•
1.8
• TC = (34 OF - 32 OF)
•
1.8
• TC = (2 OF)
•
1.8
– Example:
• What is the equivalent Celsius temperature of
400.0 K? The equivalent Fahrenheit
temperature?
• Use: TK = TC + 273
• Rearrange to : TC = TK - 273
• TC = 400.0 K - 273 = 127.0 OC
• TF = 1.8 (127.0 OC) + 32 OC =
• Heat.
• Internal and External Energy.
– External energy is total potential and kinetic
energy of everyday sized objects.
– Internal energy is the total kinetic and potential
energy of an object molecules.
One theory about how friction results in increased
temperatures: Molecules on one moving surface will
catch on another surface, stretching the molecular
forces that are holding it. They are pulled back to their
home position with a snap, resulting in a gain of
vibrational kinetic energy.
External energy is the kinetic and potential energy that
you can see. Internal energy is the total kinetic and
potential energy of molecules. When you push a table
across a floor, you do work against friction. Some of
the external mechanical energy goes into internal
kinetic and potential energy, and the bottom surface of
the legs becomes warmer.
• Heat as Energy Transfer.
– Temperature is a measure of the average kinetic
energy of an object.
– Heat is a measure of the internal energy that has
been absorbed or transferred from one body to
another.
• Increasing the internal energy is called heating.
• Decreasing the internal energy is called
cooling.
– Two ways to increase temperature:
• From a temperature difference, with energy
moving from a region of higher temperature to
a region of lower temperature.
• From an object gaining energy by way of a
temperature conversion.
Heat and temperature are different concepts, as shown
by a liter of water (1,000 mL) and a 250 mL cup of
water, both at the same temperature. You know the
liter of water contains more heat since it will require
more ice cube to cool it, say, 25OC than will be
required for the cup of water. In fact, you will have to
remove 48,750 additional calories to cool the liter of
water
• Measures of Heat.
– The metric unit of measuring work, energy, or
heat is the joule.
– The metric unit of heat is the calorie.
• A calorie is the amount of energy needed to
increase the temperature of 1 gram of water 1
OC (from 14.5 OC to 15.5 OC.
• A kilocalorie is the amount of energy needed to
increase the temperature of 1 kg of water 1 OC.
The Calorie value of food
is determined by
measuring the heat
released from burning the
food. If there is 10.0 kg
of water and the
temperature increased
from 10OC to 20OC the
food contained 100
Calories (100,000
calories). The food
illustrated here would
release much more
energy than this.
Joule worked with the English system of measurement
used during his time. When a 100 lb object falls 7.78
ft, it can do 778 fl?lb of work. If the work is done
against friction, as by stirring 1 lb of water, the heat
produced by the wok raises the temperature 1OF.
– The English unit of heating is the BTU.
• A BTU is the amount of energy needed to
increase the temperature of 1 lb of water 1 OF.
• A Quad is 1 quadrillion BTU 1 X 1015 BTU.
• 778 ftlb = 1 BTU
• 4.184 ftlb = 1 calorie
• 4,184 J = 1 kcalorie
– Example: a 2,200.0 kg automobile is moving at
90.0 km/hr (25.0 m/s). How many kilocalories
are generated when the car brakes to a stop?
• KE = 1/2mv2
• KE = 1/2(2,200 kg)(25.0m/s)2
• KE = (1,100 kg)(625.0 m2/s2)
• KE = 687,500 m2/s2
• KE = 687,500 J
• Kcal = 687,500 J X 1 kcal/4,184J = 164
kcalories
• Specific Heat.
– Three variables that influence energy transfer.
• The temperature change.
• The mass of the substance.
• The nature of the material being heated.
– The amount of heat (Q) needed to increase the
temperature (Ti) of a pot of water from the initial
temperature to a final temperature (Tf) is
proportional to (Tf-Ti).
• Q  (Tf-Ti).
• Q  T.
– The quantity of heat (Q) absorbed or given off
during a certain change in temperature is also
proportional to the mass (m) of the substance.
•Qm
– Putting this all together we get:
• Q  mcT
• c is the specific heat of the substance.
• Specific heat is the energy needed to increase
the temperature of 1 gram of a substance 1 OC.
– When two materials of different
temperatures are involved in heat transfer
and are perfectly insulated from the
surroundings, the heat lost by one is equal to
the heat gained by the other.
• heat lost = heat gained.
• Qlost = Qgained
• (mcT)lost = (mcT)gained
Of these three
metals, aluminum
needs the most
heat per gram per
degree when
warmed, and
releases the most
heat when cooled.
– Example: How much heat must be supplied to a 500.0
g pan to increase its temperature from 20.0 OC to 100.0
OC if the pan is made of a) iron and b) aluminum.
• Iron from table 5.2 has a specific heat of 0.11
cal/gOC.
• Q = mcT
• Q = (500.0g)(0.11 cal/gOC)(80.0OC)
• Q = 4,400 cal or 4.40 kcalories
• Aluminum from table 5.2 has a specific heat of 0.22
cal/gOC
• Q = mcT
• Q = (500.0g)(0.22 cal/gOC)(80.0OC)
• 8,800 calories or 8.80 kcalories
• Heat Flow.
– Conduction.
• Anytime there is a temperature difference; there
is a natural tendency for temperature to flow
from the area of higher temperature to the area
of lower temperature.
• Conduction is the transfer of energy from
molecule to molecule.
• The rate depends on the temperature difference,
the area and thickness of the substance, and the
nature of the material.
Thermometers place
in holes drilled in a
metal rod will show
that heat is conducted
from a region of
higher temperature to
a region of lower
temperature. The
increased molecular
activity is passed
from molecule to
molecule in the
process of
conduction.
• Some materials are good conductors while
others are good insulators.
–Conductors transfer energy very efficiently.
–Insulators transfer energy very inefficiently,
–The best conductors are usually metals
which have very little air space between
molecules.
–The best insulators have a great deal of air
space between molecules.
–The absolute best insulator is a vacuum as
there are no molecules to pass on energy.
Fiberglass insulation is rated in terms of R-value, a
ratio of the conductivity of the material to its
thickness.
– Convection.
• Large scale transfer of heat by a large scale
displacement of groups of molecules with
relatively higher kinetic energy.
• Molecules with higher kinetic energy are
moved from one place to another place.
• Happens only in liquids and gases where fluid
motion can carry molecules with higher kinetic
energy over a distance.
(A) Two identical volumes of air are balanced, since
they have the same number of molecules and the same
mass. (B) Increased temperature causes one volume to
expand from the increased kinetic energy of the gas
molecules. (C) The same volume of the expanded air
now contains fewer gas molecules and is less dense,
and it is buoyed up by the cooler, more dense air.
Convection
currents move
warm air
throughout a
room as the air
over the heater
becomes warmed,
expands, and is
moved upwards
by cooler air.
– Radiation.
• Radiation involves the form of energy called
radiant energy that moves through space.
• All objects with a temperature above absolute
zero give off radiant energy.
• The absolute temperature of the object
determines the rate, intensity, and kinds of
radiant energy emitted.
• Energy, Heat, and Molecular
Theory.
• Phase Change.
– The motion of a molecule can be increased by:
• Adding heat through a temperature difference.
• The absorption of one of the five forms of
energy.
• Temperature increases according to the specific
heat of the substance.
– When a substance changes from one state to
another, the transition is called a phase
change.
• A phase change always absorbs of releases
energy, a quantity of heat that is not associated
with a temperature change.
• Latent heat is the hidden energy of a phase
change, which is energy that goes in or comes
out of internal potential energy.
– Three major types of phase change.
• Solid-liquid.
• Liquid-gas.
• Solid-gas
– Solid-liquid.
• The temperature at which a substance changes
from a liquid to a solid is called the freezing
point.
• The temperature at which a solid changes to a
liquid is the melting point.
• Both of these occur at the same temperature.
– Liquid-gas.
• The temperature at which a liquid changes from
the liquid phase to the gaseous phase is the
boiling point.
• The temperature at which a gas or vapor
changes to the liquid phase is the condensation
point.
• Both of these occur at the same temperature.
– Solid-gas.
• A phase change directly from a solid to a gas or
vapor is called sublimation.
Each phase
change absorbs
or releases a
quantity of
latent heat,
which goes
into or is
released from
molecular
potential
energy.
This graph shows three
warming sequences and two
phase changes with a
constant input of heat. The
ice warms to the melting
point, then absorbs heat
during the phase change as the temperature remains
constant. When all the ice has melted, the now liquid
water warms to the boiling point, where the
temperature again remains constant as heat is absorbed
during the second phase change from liquid to gas.
After all the liquid has changed to gas, continued
warming increases the temperature of the water vapor.
(A)Work is done against gravity to lift an object,
giving the object more gravitational potential energy.
(B) Work is done against intermolecular forces in
separating a molecule from a solid, giving the
molecule more potential energy.
Compare this graph
to the one in Figure
5.20. This graph
shows the
relationships between
the quantity of heat
absorbed during
warming and phase changes as water is warmed from
ice at -20OC to water vapor at some temperature above
100OC. Note that the specific heat for ice, liquid water,
and water vapor (steam) have different values.
– Latent heat of fusion.
• The latent heat of fusion is the heat involved in
a solid-liquid phase change in melting or
freezing.
• A melting solid absorbs energy and a freezing
liquid releases this same amount of energy,
warming the surroundings.
• The total heat involved in a solid-liquid phase
change depends on the mass of the substance
involved.
–Q = mLf
–Where Lf is the latent heat of fusion for the
substance involved
– Latent heat of vaporization.
• The amount of heat involved during a phase
change from a liquid to a gas or vapor is called
the latent heat of vaporization.
• The latent heat of vaporization is the heat
involved in a liquid-gas phase change where
there is evaporation or condensation.
• The escaping molecules absorb energy from the
surroundings, and a condensing gas releases
this exact same amount of energy.
• The total heating depends on the amount of
water vapor condenses so that:
–Q = mLV
–Where LV is the latent heat of vaporization
for the substance involved.
• Example:
– How much energy does a refrigerator remove
from 100.0 g of water at 20.0 OC to make ice at 10.0 OC
– Three steps.
– Q1 = mcT to cool from 20.0 OC to 0.0 OC
– Q1 = (100.0g)(1.00cal/gOC)(0.0OC-20.0OC)
– = 2,000 cal = 2.00 X 10 3 cal.
– Q2 = mLf to remove latent heat of fusion.
– (100.0g)(80.0cal/g)
– 8,000 cal = 8.00 X 103 cal
– Q3 = mcT to go from 0.0 OC to -10 OC
– (100.0g)(0.500cal/g)(10.0OC-0.0OC)
– 500 cal = 5.00 X 102 cal
– Qtotal = Q1 + Q2 + Q3
– = 2.00 X 103 cal + 8.00 X 103 cal + 5.00 X 102 cal
– = 1.05 X 104 cal
• Evaporation and Condensation.
– Evaporation occurs when enough energy is
inputed into a system to cause liquid molecules to
overcome attractive forces near the surface,
escape, and become a gas or vapor.
– In evaporation, more molecules are leaving the
liquid state than are returning.
– In condensation, more molecules are returning to
the liquid state than are leaving.
– When the condensation rate is equal to the
evaporation rate, the air above the liquid is
saturated (holds all the vapor that it is capable of
holding).
Temperature is associated
with the average energy
of the molecules of a
substance. These numbered
circles represent arbitrary
levels of molecular kinetic energy that, in turn,
represent temperature. The two molecules with the
higher kinetic values [25 in (A)] escape, which lowers
the average value from 11.5 to 8.1 (B). Thus
evaporation of water molecules with more kinetic
energy contributes to the cooling effect of evaporation
in addition to the absorption of latent heat.
– Four ways to increase the rate of evaporation.
• An increase in temperature of the liquid will
increase the average kinetic energy of the
molecules and thus increase the number of high
energy molecules capable of escaping from the
liquid state.
• Increase the surface area of the liquid in contact
with the air.
• Removal of water vapor from near the surface
will prevent the return of molecules to the
liquid phase.
• Reducing atmospheric pressure will reduce one
of the forces holding molecules in a liquid.
• Relative Humidity.
– The ratio of how much water vapor is in air to
how much water vapor it could hold at a certain
temperature is the relative humidity
– Usually expressed as a percent.
The inside of this closed
bottle is isolated from the
environment so the space
above the liquid becomes
saturated. While it is
saturated, the evaporation
rate equals the condensation
rate. When the bottle is
cooled, condensation
exceeds evaporation and
droplets of liquid form on
the inside surfaces.
The curve shows the
maximum amount of
water vapor in g/m3
that can be in the air
at various
temperatures.
• Thermodynamics.
• Introduction.
– The laws of thermodynamics describe what
happens to energy as it is transformed into work
and to other forms.
– Thermodynamics is concerned with internal
energy, which is the total internal kinetic and
potential energy of a system.
– The system is the component we want to describe.
– The state of the system are the variable under
which it exists, temperature, pressure, volume,
heat, etc…
– Everything outside of the system is the
surroundings.
A very simple heat
engine. The air in (B)
has been heated,
increasing the molecular
motion and thus the
pressure. Some of the
heat is transferred to the
increased gravitational
potential energy of the
weight as it is converted
to mechanical energy.
• The First Law of Thermodynamics.
– The energy supplied to a system is equal to the
change in internal energy
• The Second Law of Thermodynamics.
– Heat flows from objects with a higher temperature
to objects with a cooler temperature.
• The Second Law and Natural Processes.
– Energy can be viewed from two considerations of
scale:
• The observable external energy of an object.
• The internal energy of the molecules, or
particles that make up an object.
– Two kinds of motion that the particles of an
object can have.
• A coherent motion where they move together.
• An incoherent, chaotic motion of individual
particles.
– Work on an object is associated with coherent
motion, while heating an object is associated with
its internal incoherent motion.
The heat supplied (QH)
to a heat engine goes into
the mechanical work (W)
and the remainder is
expelled in the exhaust
(QL). The work
accomplished us
therefore the difference
in the heat input and
output (QH) - (QL), so the
work accomplished
represents the heat used,
W = J(QH - QL)
A heat pump uses work
(W) to move heat from a
low temperature region
(QL) to a high temperature
region (QH). The heat
moved (QL) requires work
(W), so J QL = W.
– Entropy.
• Energy is always degrading toward a more
disorderly state.
• The total entropy of the universe is continually
increasing.
• The natural process is for the sate of order to
degrade into a state of disorder with a
corresponding increase in entropy.
– Eventually all of the useable energy in the
universe will diminish to unusable forms.
• The universe will at some time reach a limit of
disorder called the heat death of the universe.
• The heat death of the universe is the theoretical
limit of disorder, with all molecules spread far,
far apart, vibrating slowly with a uniform low
temperature.