Thermodynamics

Heat & Thermodynamics!
Kinetic-molecular theory

Specific Heat

Heat & Internal Energy

Calorimetry

Thermal Equilibrium

Heat Transfer Processes

Temperature Scales

Phase changes

Laws of Thermodynamics

Thermal Expansion

Entropy

Heat Engines & Refrigerators
Latent Heat of Vaporization

Latent Heat of Fusion

Kinetic-Molecular Theory
It was once common belief that heat was an invisible substance.
It even had a name--“caloric,” and it was believed that it could be
transferred between objects but neither created nor destroyed. To
heat up an object this caloric had to flow into it. This, they
thought, explained why objects expanded when heated. But this
theory could not explain, for example, how heat could emanate
from a cold piece of wood once it is set on fire? Where did the
caloric come from? If it had been in the wood in the first place,
the wood should have been hot all along.
The caloric theory was abandoned in the 19th century and replaced
with the kinetic-molecular theory. This new theory stated that all
matter is made up of atoms/molecules in constant motion. The
faster they move, the hotter an object will be.

Internal Energy
Internal energy (also called thermal energy) is the energy an
object or substance is due to the kinetic and potential energies
associated with the random motions of all the particles that make
it up. The kinetic energy is, of course, due to the motion of the
particles. To understand the potential energy, imagine a solid in
which all of its molecules are bound to its neighbors by springs.
As the molecules vibrate, the springs are compressed and
stretched. (Liquids and gases are not locked in a lattice structure
like this.)
The hotter something is, the faster its
molecules are moving or vibrating, and the
higher its temperature. Temperature is
proportional to the average kinetic energy
of the atoms or molecules that make up a
substance.

Units for Heat
Like any type of energy, the SI unit for heat is the Joule.
Another common unit is the calorie, which is approximately
the amount of heat energy needed to raise one gram one degree
Celsius. 1000 calories are in a Calorie, which is used to
measure the energy in foods (that the human body can make
use of). The British thermal unit (BTU) is approximately the
energy needed to raise one pound of water one degree
Fahrenheit.

1 cal = 4.186 J

1 BTU = 1055 J = 252 cal

Temperature vs. Internal Energy
Temperature and internal energy are related but not the same thing.
Temperature is directly proportional to the average molecular
kinetic energy*. Note the word average is used, not total.
Consider a bucket of hot water and a swimming pool full of cold
water. The hot water is at a higher temperature, but the pool water
actually has more internal energy! This is because, even though the
average kinetic energy of the water molecules in the bucket is much
greater than that of the pool, there are thousands of times more
molecules in the pool, so their total energy is greater.
It’s analogous to this: A swarm of 1000 slow moving bees could
have more total kinetic energy than a dozen fast moving,
hyperactive bees buzzing around like crazy. One fast bee has more
kinetic energy than a slow one, but there are a lot more slow ones.
*

true for gases, approximately true for solids and liquids whose
molecules interact with each other more.
contintued on next slide

Temperature Scales
Fahrenheit: water freezes at 32 °F; boils at 212 °F
Celsius: water freezes at 0 °C; boils at 100 °C
Kelvin: water freezes at 273.15 K; boils at 373.15 K
A change of 100 °C corresponds to a change of 180 °F. This means
5 C° = 9 F° or 1 C° = 1.8 F° Note that the degree symbol is on the
opposite side of the letter, indicating that we’re talking about
temperature differences. In other words, five steps on the Celsius
scale is equivalent to nine steps on the Fahrenheit scale, but 5 °C is
certainly not equal to 9 °F. Since these scales are linear, and they’re
offset by 32 °F, we get the conversion formula: F = 1.8C + 32
One step on the Kelvin scale is the same as one step on the Celsius
scale. These scales are off by 273.15 K, so: K = C + 273.15
Room temperature is around 293 kelvins, which is 20 °C, or 68 °F.

Absolute Zero & the Kelvin Scale
The Kelvin scale is setup so that its zero point is the coldest possible
temperature--absolute zero, at which point a substance would have
zero internal energy. This is -273.15 °C, or -459.69 °F. Absolute zero
can never be reached, but there is no limit to how close we can get to
it. Scientists have cooled substances to within 10-5 kelvins of absolute
zero. How do we know how cold absolute zero is, if nothing has ever
been at that temperature? The answer is by graphing Pressure vs.
Temperature for a variety of gases and extrapolating.
P
A gas exerts no
pressure when at
absolute zero.

as A
g

as B
g

gas C

T (°C)
-273.15 °C

0 °C

Thermal Equilibrium
Two bodies are said to be at thermal equilibrium if they are at the
same temperature. This means there is no net exchange of thermal
energy between the two bodies. The top pair of objects are in
contact, but since they are at different temps, they are not in thermal
equilibrium, and energy is flowing from the hot side to the cold side.
heat

hot

26 °C

cold

26 °C

No net heat flow
The two purple objects are at the same temp and, therefore are in
thermal equilibrium. There is no net flow of heat energy here.

Heat Transfer Processes
Heat energy can be transferred from one body to another in three
different ways. Upcoming slides will give an example of each.
• Conduction: Energy is transferred when two objects are in direct
contact. Molecules of the hotter object bump into molecules of the
colder object and cause them to speed up, warming the colder object.
• Convection: Energy is transferred from one body to a cooler one
via currents in a fluid (a gas or liquid).
• Radiation: All objects, at any temperature, radiate electromagnetic
radiation (light of visible and invisible wavelengths). Unlike
conduction & convection, no medium (matter of any type) is
necessary for heat transfer through radiation. Objects absorb
radiation as well. At thermal equilibrium it will absorb as much as it
radiates.

Conduction

Schmedrick decides to become a blacksmith. In order to forge a horseshoe for his horse, Bucephalus, Scmedrick heats up the shoe in a fire,
pounds on it with a mallet to shape it, and then cools it by dipping it in
a bucket of water. Because the water is colder, heat flows from the
shoe to the water--quickly at first, and more slowly as the shoe cools.
The water molecules, with little kinetic energy, are in direct contact
with the iron atoms, which are jiggling rapidly and have lots of kinetic
energy. When an iron atom bumps into a water molecule, the iron
atom slows down a bit, while the water molecule speeds up (an elastic
collision). In this way water gains the heat energy that the iron loses.
water molecule
iron atom
zoomed in view

Convection
The water near the hot horseshoe is warmer than the water further
from the shoe. This warm water is lower in density than the cooler
water, since its molecules are moving faster and taking up more space.
With lower density, the warm water begins to float to the surface,
carrying its heat energy with it. As it rises to the surface it cools and
becomes denser. Then it begins to sink, warmer water from below
taking its place. These convection currents transfer heat from the
horseshoe to the air via the water, which is the convection medium.
If the water were surrounded by
something solid or too viscous to flow,
heat could only be transferred to the air
via conduction, and it would take much
longer. Convection plays a big role in
determining global weather patterns.

Radiation
The molecules of warm water cooling the horseshoe at the surface of
Schmedrick’s bucket bump into air molecules and transfer heat to the
air via conduction. The water can also transfer energy to the air by
emitting electromagnetic radiation. This is simply light, but usually
it’s light of a wavelength that is too long for us to see--infrared.
Bodies also continually absorb radiation, but when a body is warmer
than its surroundings, it emits more than it
absorbs. Night vision technology takes advantage of this fact by detecting infrared light in
order to “see in the dark.” Radiation can cool
or warm objects even if they are surrounded
by a vacuum. (Even a perfect Thermos bottle
full of hot chocolate will eventually cool
down.) When Schmed’s bucket cools long
enough, it will be in thermal equilibrium with
the air, and the net radiation (emission absorption) will be zero.

Radiation: Power & Temperature
The rate at which a hot object emits radiation is its power output.
Recall, power, P, is the rate at which work is done or energy is
expended or absorbed. P depends on the body’s temp (in kelvin) and
on the amount of surface area it has. Power is directly proportional to
the surface area and proportional to the 4th power of absolute
temperature:

P ∝ AT4

Note that the closer the radiating body gets to absolute zero, the
lower its power output of electromagnetic radiation, meaning the
amount of internal energy it is radiating out in a unit of time is low.
Also, an object with lots of surface area will radiate at a greater rate.
Don’t forget that bodies radiate and absorb energy at the same time.
The same equation describes absorption, except we use the temp of
the surroundings. Pnet = 0 when a body is in thermal equilibrium.

Black Body
A black body is an ideal absorber. It absorbs any radiation that is
incident upon it (any light that hits it). It exists only in theory.
Suppose Schmedrick has Bucephalus is all shoed up and ready to
run. Schmed hops on the back of his trusty steed, and with a mighty
“Hi ho Bucephalus! Away!” he heads off into the sunset. Before
falling off, Schmedrick ponders the sunlight streaming through the
atmosphere from 93 million miles away. Not all of the light that
reaches Earth makes it to the surface. The atmosphere reflects some
of it back into space and absorbs some of it. (It scatters away more
of the blue light than the red, which is why sunsets look red.) It is
the same story for the light hitting Bucephalus: his coat absorbs
some of it (and warms him); and some is reflected (otherwise he
would be called Bucephalus the Invisible Horse).
All real-world objects interact this way with light. Only a black
body would absorb all light, including wavelengths we can’t see.

Thermal Conductivity, k
Heat transfer via conduction was described a few slides back. Thermal
conductivity, k, refers how easily heat can move through a material.
Metals have high thermal conductivity, meaning heat passes through
them readily. Wood is a fairly good insulated of heat, and styrofoam is
even better. These materials have low thermal conductivities. k is
very low for air as well. (Attic insulation and styrofoam cups trap air,
making them good insulators.) Heat from a boiler passes through all
sides of its metal enclosure. The rate at which heat is transferred is
given by:

k A (T - T )
H=
2
1
L

T2

T1
heat

A = area of side wall
L = thickness of wall
k = thermal conductivity of the metal
T2 - T1 = temperature difference
H is simply power, and its SI unit is the Watt.

SI Units for Thermal Conductivity
k A (T - T )
H=
2
1
L
k must have units that cancel out all the units on the right,
leaving only the units for H. The units are:

W
m·K

or equivalently,

W
m · °C

Since one kelvin is as big a change in temp as one degree
Celsius, these units are equivalent.
Note: k for thermal conductivity is not the same as the k
in Hooke’s Law in which it represents the spring
constant!

Thermopane Windows

heat

In a house we often want to prevent heat from getting
in or getting out. Windows can be problematic.
Thermopane windows have two or more panes of
glass with air or some other gas between the panes.
Which type of window, a double pane or a thick
single pane, is better for minimizing heat transfer, if
the total thickness is the same?

heat

There is more glass in the single pane window to
block the heat, but the air in between the panes of the
double pane window has thermal conductivity that is
about 35 times lower than that of the glass itself. So
much more heat would be transferred through the
single pane.

Laws of Thermodynamics
• Zeroth Law: If object A is in thermal equilibrium with object B,
and if object B is in thermal equilibrium with object C, then objects
A and C are also in equilibrium. This is sort of a “transitive
property of heat.”
• First Law: Energy is always conserved. It can change forms:
kinetic, potential, internal etc., but the total energy is a constant.
Another way to say it is that the change in thermal energy of a
system is equal to the sum of the work done on it and the amount
of heat energy transferred to it.
• Second Law: During any natural process the total amount of
entropy in the universe always increases. Entropy can be defined
informally as a measure of the randomness or disorder in a system.
Heat flows naturally from a hot to cooler surroundings as a
consequence of the second law.

Entropy:
Statistical Approach
Entropy is related to probability. Let’s
look at the possible outcomes of flipping
four coins, of which there are sixteen
(2 4 = 16). The outcomes are grouped into
macrostates according to the number of
heads. Each macrostate is made up a
microstates. For example, the 3-heads
macrostate is comprised of 4 microstates,
because there are 4 combinations that
yield 3 heads. One microstate in the 3heads macrostate is H H T H. The number
of microstates in a macrostate determines
how likely that state is to exist.
continued on next slide

H H H H

4 heads

T
H
H
H

H
T
H
H

H
H
T
H

H
H
H
T

3 heads

T
T
T
H
H
H

T
H
H
T
T
H

H
T
H
T
H
T

H
H
T
H
T
T

2 heads

T
T
T
H

T
T
H
T

T
H
T
T

H
T
T
T

1 head

T

T T T 0 heads

Entropy
Macrostate
0
1
2
3
4

(cont.)

# of Microstates
1
4
6
4
1

Probability
1 / 16
1/4
3/8
1/4
1 / 16

Macrostate 3 (the group w/ 3 heads) is the most
probable since it contains the most microstates (combinations). Macrostate 2 has 6 microstates, so its
probability is 6 / 16 = 3 / 8. This macrostate is the most
random, or disordered, since there are so many ways 2
heads can come up in 4 flips. Entropy is a measure of
disorder, and for this system it’s at a max when in
macrostate 2. Minimum entropy occurs when the coins
are in macrostate 0 or 4, since there is a high degree of
order in these states--only one microstate each. These
are the least likely microstates to occur.
continued

H H H H
T
H
H
H

H
T
H
H

H
H
T
H

H
H
H
T

T
T
T
H
H
H

T
H
H
T
T
H

H
T
H
T
H
T

H
H
T
H
T
T

T
T
T
H

T
T
H
T

T
H
T
T

H
T
T
T

T

T T T

Entropy

(cont.)

Suppose our coin system is in macrostate 4 (all heads). This
represents maximum order, minimum entropy. Every so often one of
the coins is chosen at random and flipped. With each flip there is a
50-50 chance that the macrostate will change. With time (after
enough flips), it is doubtful that the system will still be in the
minimum entropy state. It is much more likely to be in macrostate 2,
the state with the most entropy.
The 2nd Law states that during any process the universe moves toward
more probably states--states with more entropy. It is possible to
decrease the entropy of our coin system by physically turning all tails
over so that there are all heads, but in doing this we must expend
energy. This energy expenditure increases the entropy of our
surroundings more than it decreases the entropy of the system. Thus
the entropy of the universe is increased.

Entropy

(cont.)

In our coin example we only dealt with four coins. In real life even a
quadrillion atoms or molecules might not be very much. (A single
bacterium contains about 100 billion atoms.) How much more likely
is it for a system to be in its highest entropy state than in its lowest?
It depends on how big the system is:
Number of
Coins
4
10
20
50
100

Ratio of
Probabilities
6:1
252 : 1
184,756 :1
~1014 : 1
~1029 : 1

This means that if 100 coins were
dumped on the floor it is about 100
billion billion billion times more
likely for half the coins to come up
heads than for all of them to be
heads!

Entropy: Statistics Formula
We’ve seen that there are six ways to get exactly two heads in four flips.
There were only sixteen combinations of four heads and tails, so we just
listed them and counted how many had exactly two heads. But you
wouldn’t want to have to list all the combinations in fifty flips, since
there are 250 combos—over a quadrillion lines of 50 H’s and T’s! So
we’ll use some math instead. The number of ways to place 50 H’s in
100 spots is “100 choose 50,” which is written like this: 100
50
In general,

( )
n
r

n!
=
r ! (n – r ) !

( )

Let’s try out the formula with 2 heads in 4 flips:

( )
4
2

=

4!
4·3·2·1
= 6, as we showed by
=
2 ! (4 – 2) !
(2 · 1) (2 · 1)
listing combinations

Entropy & Fluids

Suppose a beaker of very hot water is poured into an aquarium of cool
water. Conservation of energy would not be violated if all the hot
water remained right at the spot where it was poured. But the 2 nd Law
demands that the thermal energy eventually become evenly
distributed. The cool water has molecules moving at a wide range of
speeds (red = fast; blue = slow). Since the water is cool, there are
more blues than reds. The hot water poured in has mostly red. The
aquarium has less disorder (entropy) when all the fast molecules are in
one spot than when they are mixed in. With time a much more likely
situation exists, with a much higher entropy.
continued

time

Entropy & Fluids

(cont.)
Imagine how many different ways you could take 100 blue balls and
paint 8 of them red. There are about 1.86 · 1011 ways to do this.
Many, many more of those ways look like the picture on the right
than on the left. The diffusion of perfume from an open bottle
throughout a room is also a consequence of the 2nd Law. Unlike
diffusion, though, the “hot” water molecules don’t necessarily have
to move so that they are spread out evenly. Convection currents will
allow some to move, but it is really the heat energy rather than the
molecules themselves that must distribute itself equally throughout
the aquarium.

Change in Entropy Equation
Because most systems are many up of so many particles, calculating
entropy via probabilities would be very difficult. Fortunately, we are
normally concerned only with changes in entropy. If we have a
system in which energy is not changing forms, the change in entropy
is defined as:

∆Q
∆S =
T
∆ S = change in entropy
∆ Q = change in internal energy (heat flow)
T = absolute temperature
The 2nd Law of Thermodynamics says that during any process:

∆ Suniverse = ∆ Ssystem + ∆ Ssurroundings ≥ 0

Change in Entropy Example

(cont.)
As the glass cooled we assumed that the air temp didn’t go up
appreciably due after the heat transfer, which would have complicated the problem. Important points:
• The temps were converted to kelvins.
• The glass lost as much thermal energy as air gained, as the 1st Law
requires.
• ∆ Qglass is negative since the glass lost thermal energy so ∆ Sglass is
also negative.
• ∆ Qair is positive since the air gained thermal energy so ∆ Sair is
also positive.
• Even though the ∆ Q’s are the same size, the ∆ S’s aren’t, since the
temps are different.
• The positive ∆ S is greater than the negative ∆ S, as the 2nd Law
requires.

Second Law Consequences
• Heat will not flow from a cold body to a hot body.
• “Reverse diffusion” is a no-no (such as smoke from a fire
isolating itself in a small space).
• An object or fluid of uniform temperature (no matter how hot)
cannot do useful work. (There must be temperature difference
so that there will be a heat flow, which can be used to do work.)
• The various forms of energy tend to degrade over time to
thermal energy. This represents useful, low probability forms of
energy converting into an unusable, high probability form.
• Without input of energy, bodies tend to reach thermal
equilibrium. (We can maintain temperature differences via
refrigerators or heating units, but this requires energy.)

Second Law Consequences

(cont.)

• Any time we do something that decreases the entropy of a
system, the energy we expend in doing it increases the entropy
of the surroundings even more.
• A perpetual motion machine is impossible to make. A
perpetual motion machine is a device that would absorb thermal
energy from a hot body and do as much work as the energy it
absorbed. (See pics on next slide.)
• During any process the entropy of the universe cannot
decrease. Expending energy to decrease the entropy of a system
will lead to an increase in entropy for the surrounding by a
greater amount.

Heat Engines
A heat engine takes advantage of temp differences to produce useful
work. The amount of work done depends on the size of the reservoirs,
engine efficiency, and the temp difference (TH - TC). QH is the heat that
flows from the hot region; QC is the heat flowing into the cold region.
W is the useful work done by engine. The smaller QC is, the more
efficient the engine is. The engine on the right satisfies the 1st Law but
violates the 2nd Law, i.e., 100% efficiency is unattainable.
Hot Reservoir, T H
QH

W

Hot Reservoir, T H
QH

Engine

Engine

QC

W

QC = 0

Cold Reservoir, T C
Real engine. QH = QC + W

Cold Reservoir, T C
Impossible engine. QH = W

Refrigerators
A refrigerator forces heat from a cold region into a warmer one. It
takes work to do this, otherwise the 2nd Law would be violated. Can
a fridge be left open in the summer to provide a make shift air conditioner? Nope, since all heat pumped out of the fridge is pumped back
into the kitchen. Since QH > QC because of the work done, leaving
the refrigerator open would actually make your house hotter!
Hot Reservoir, T H
QH

W

Hot Reservoir, T H
QH

Engine

Engine

QC

W=0

QC

Cold Reservoir, T C
Real fridge. QC + W = QH

Cold Reservoir, T C
Impossible fridge. QC = QH

Specific Heat
Specific heat is defined as the amount of thermal energy needed
to raise a unit mass of substance a unit of temperature. Its
symbol is C.
For example, one way to express the specific heat of water is
one calorie per gram per degree Celsius: C = 1 cal / (g · ºC), or
4.186 J / (g · ºC). This means it would take 20 cal of thermal
energy to raise 4 grams of water 5 ºC.
Water has a very high specific heat, so it takes more energy to
heat up water than it would to heat up most other substances (of
the same mass) by the same amount. Oceans and lake act like
“heat sinks” storing thermal energy absorbed in the summer and
slowing releasing it during the winter. Large bodies of water
thereby help to make local climates less extreme in temperature
from season to season.

Specific Heat Equation

Q = m C ∆T
Q = thermal energy
m = mass
C = specific heat
∆T = change in temp
Ex: The specific heat of silicon is 703 J / (kg · ºC). How much
energy is needed to raise a 7 kg chunk of silicon 10 ºC ?
answer:
Q = 7 kg · 703 J ·10 ºC = 49 210 J
kg · ºC
Note that the units do indeed work out to be energy units.

Calorimetry
Schmedrick takes another horseshoe out of the fire when it’s at
275 ºC, drops in his bucket of water, and this time covers the bucket.
The bucket and cover are made of an insulating material. The bucket
contains 2.5 L of water originally at 25 ºC. The 1.9 kg shoe is made
of iron, which has a specific heat of 448 J / (kg · ºC). Let’s find the
temp of the horseshoe and water once equilibrium is reached.
Let’s assume that the container allows no
heat to escape. Then the 1st Law implies
that all heat the shoe loses is gained by the
water. Since one milliliter of water has a
mass of one gram, the bucket contains
2.5 kg of water. At thermal equilibrium
the water and shoe are at the same temp.
The total thermal energy in the bucket
does not change, but it is redistributed.

Calorimetry

(cont.)

Let T = the equilibrium temperature.

Q lost by iron = Q gained by water
miron Ciron ∆Tiron = mwater Cwater ∆Twater
(1.9 kg) (448 J / kg · ºC) (275 ºC - T) = (2.5 kg) (4186 J / kg · ºC) (T - 25 ºC)
Note how the ∆T terms are written so that
each side is positive. We’ve got a simple
linear equation with T on both sides.
Solving it gives us T = 43.8 ºC. This is
the equilibrium temp--the final temp for
both the shoe and water. If T had come
out over 100 ºC, the answer would have
been invalid, since the specific heat for
steam is different than that of water.

Latent Heat
The word “latent” comes from a Latin word that means “to lie
hidden.” When a substance changes phases (liquid ↔ solid or
gas ↔ liquid) energy is transferred without a change in
temperature. This “hidden energy” is called latent heat. For
example, to turn water ice into liquid water, energy must be
added to bring the water to its melting point, 0 ºC. This is not
enough, however, since water can exist at 0 ºC in either the liquid
or solid state. Additional energy is required to change 0 ºC ice
into 0 ºC water. The energy increases the internal energy of the
water but does not raise its temp. When frozen, water molecules
are in a crystalline structure, and energy is needed to break this
structure. The energy needed is called the latent heat of fusion.
Additional energy is also needed to change water at 100 ºC to
steam at 100 ºC, and this is called the latent heat of vaporization.

Latent Heat Formula

Q = m Lf or

Q = m Lv

Q = thermal energy
m = mass
L = heat of fusion or vaporization
L is the energy per unit mass needed to change the state
of a substance from solid to liquid or from liquid to gas.
Ex: Lf (the latent heat of fusion) for gold is 6440 J / kg.
Gold melts at 1063 ºC. 5 grams of solid gold at this

temp will not become liquid until additional heat is
added. The amount of heat needed is:
(6440 J / kg) (0.005 kg) = 32 J. The liquid gold will still be
at 1063 ºC.

Latent Heat / Specific Heat Example
Superman vaporizes a 1800 kg ice monster with
his heat ray vision. The ice monster was at
-20 ºC. After being vaporized he is steam at
135 ºC. How much energy did Superman expend?
Substance
Specific Heat (in J / kg · ºC)
ice
2090
liquid water
4186
steam
1970
For water : Lf = 3.33 ·105 J / kg; Lv = 2.26 ·106 J / kg
Q = (1800 kg) (2090 J / kg · ºC) (20 ºC) heating ice to melting pt.
+ (1800 kg) (3.33 ·105 J / kg) ice to water, const. temp of 0 ºC
+(1800 kg) (4186 J / kg · ºC) (100 ºC) heating water to boiling pt.
+ (1800 kg) (2.26 ·106 J / kg) water to steam, const. temp of 100 ºC
+ (1800 kg) (1970 J / kg · ºC) (35 ºC) heating steam to 135 ºC
= 5.62 ·109 J total energy expended by Superman

Latent Heat & Entropy

Schmedrick is enjoying a cool glass of soy milk while relaxing on a
cot on a winter morning in his backyard. Suddenly his dog, Rover,
barks at a squirrel and startles Schmed, who drops his drink. A 10 g
ice cube at 0 ºC falls to the ground and melts. The temp outside is
10 ºC. Calculate the change in entropy of the universe due to the
melting of the ice only. answer:
For the cubie: Q = m Lf = (0.01 kg) (3.33 ·105 J / kg) = + 3330 J.
This is the energy absorbed by the ice from the surroundings.
∆Sice = ∆Qice / Tice = +3330 J / 273 K = +12.198 J / K.
For the surroundings: Q = -3330 J, since the surroundings lost as
much thermal energy as the cubie gained. The temperature of the
backyard does not decrease significantly, though, with such a small
energy loss. ∆Ssurr = ∆Qsurr / Tsurr = -3330 J / 283 K = -11.767 J / K.
For the universe: ∆Suniv = ∆Ssurr + ∆Sice = 12.198 J / K - 11.767 J / K
= +0.431 J / K. Thus, the 2nd Law is satisfied.

Internal Energy, Work, & Heat
The internal energy, ∆Eint, of a substance or object can be changed
in two ways:
1. by letting heat flow in or out of the substance, Q
2. by the substance doing work or having work done on it, W
In summary: ∆Eint = Q - W, which is one way to state the 1st Law.

weight

heat

gas

membrane

Q is positive if heat flows in. W is the
work done by the substance. If the gas
expands because of the added heat, it
will do work by lifting the weight up.
Then W would be positive, and the
work the gas does would decrease its
internal energy.

Internal Combustion Engine
In the carburetor of your car, air and fuel are mixed. The gaseous
mixture is injected into a cylinder, compressed by a piston, and
ignited by a spark plug. (If your car has fuel injection, which is more
efficient, there is no carburetor; instead fuel is sprayed into the
cylinders at appropriate times, where it vaporizes.) The fuel mixture
contains internal as well as chemical potential energy. After burning
most of the potential energy is released. This energy heats the gas in
the cylinder, raising its internal energy. The burning gas also does
work on the piston as it expands. The force applied to the piston
causes the crankshaft to rotate. The crankshaft is hooked up to the
transmission. The exhaust gases are expelled from the cylinder so
that the cycle can begin again. Cars are very inefficient, since most
of the chemical potential energy in the gasoline goes into heating the
exhaust gases, which pollute our atmosphere and contribute to global
warming. Only a small amount of the chemical potential energy does
useful work.

Thermal Expansion

As a material heats up its atoms/molecules move or vibrate more
vigorously, and the average separation between them increases. This
results in small increases in lengths and volumes. Buildings, railroad
tracks, bridges, and highways contain thermal expansion joints to
prevent cracking and warping due to expansion. The amount of expansion depends on the original length/volume, the type of material, and
the change in temp. L is length, V is volume, T is temp, α is the coefficient of linear expansion, and β is the coef. of volume expansion.
When a solid of a single material expands, it does so proportionally in
all directions. Since volume has 3 dimensions and length is only 1,
β = 3 α.
Length expansion:

∆L
= α ∆T
L

Volume expansion:
cold solid

hot solid

∆V
= β ∆T
V

Top view
handle

Bimetallic Strip
steel

(brass on other side)

A bimetallic strip is a strip of two different metals—often steel on one
side and brass on the other. When heated the strip curves because the
metals have different coefficients of thermal expansion. Brass’s
coefficient is higher, so for a given temperature change, it expands
more than steel. This causes the strip to bend toward the steel side.
The bending would be reversed if the strip were made very cold.
steel side
Side view

brass side

Thermostats
Bimetallic strips are used in thermostats, at least in some older
ones. When the temperature changes, the strip bends, making or
breaking an electrical circuit, which causes the furnace to turn
on or shut off. In this model when the strip bends it tilts a bulb
of mercury, which then bridges two wires and allows current to
flow.

Thermodynamics

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