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Physics 1
Thermal Physics
Chapters 21-23
Links: http://homepage.mac.com/phyzman/phyz/BOP/2-06ADHT/index.html
and http://www.colorado.edu/physics/2000/index.pl
1
What do you think? know?
1. Why does popcorn pop?
2. On a camping trip, your friend
tells you that fluffing up a down
sleeping bag before you go to
bed will keep you warmer than
sleeping in the same bag when it
is still crushed. Why?
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3. Why is it difficult to
build a fire with damp
wood?
4. Why does steam at
o
100 C cause more
severe burns than liquid
o
water does at 100 C?
3
5. Until refrigerators were invented,
many people stored fruits and
vegetables in underground cellars.
Why was this more effective than
keeping them in the open air?
6. In the past, when a baby had a
high fever, the doctor might have
suggested gently sponging off the
baby with rubbing alcohol. Why
would this help?
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7. Why does water expand when it freezes?
8. Why, during the final construction of the
St. Louis arch, was water sprayed on the
previous sections as the last section was
put in place?
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Objectives
• 1. Describe thermal energy and compare it
to potential and kinetic energies.
• 2. Describe changes in temperatures of
two objects reaching thermal equilibrium
• 3. Identify various temperature scales, and
convert between them
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Objectives
• 4. Explain heat as energy transferred
between substances at different
temperatures in one of 3 ways
(conduction, convection, radiation)
• 5. Relate heat and temperature
• 6. Apply principle of energy conservation
to calculate changes in potential, kinetic, &
internal energy
7
Objectives
• 7. Perform calculations with specific heat
capacity
• 8. Interpret the various sections of a
heating curve
8
Objectives
• 9. Recognize that a system can absorb or
release energy as heat in order for work to
be done on or by that system
• 10. Compute work done during
thermodynamic process
• 11. Distinguish between isovolumetric,
isothermal, and adiabatic thermodynamic
processes
9
Objectives
• 12. Illustrate how the first law of
thermodynamics is a statement of
energy conservation
• 13. Calculate heat, work, and change in
internal energy using lst law of T-D
• 14. Apply 1st law of T-D to describe
cyclic processes
• 15. Recognize why 2nd law of T-D
requires 2 bodies at different temps.
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1.Relate temperature to the kinetic
energy of atoms and molecules
• Temperature scales
• In the USA, the Fahrenheit temperature scale is used.
Most of the rest of the world uses Celsius, and in science
it is often most convenient to use the Kelvin scale.
• The Celsius scale is based on the temperatures at which
water freezes and boils. 0°C is the freezing point of
water, and 100° C is the boiling point. Room temperature
is about 20° C, a hot summer day might be 40° C, and a
cold winter day would be around -20° C.
• To convert between Fahrenheit and Celsius, use these
equations:
•
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21.1 Temperature
Temperature and Kinetic Energy
Temperature is related to the random motions of the
molecules in a substance.
In the simplest case of an ideal gas, temperature is
proportional to the average kinetic energy of
molecular translational motion.
• Convert 72 oF to oC
• C = 5/9 (F-32)
• C = 5/9 (72-32) = 22oC
• Convert -10 oC to oF
• F = 9/5 C + 32
• F = 9/5(-10) + 32 = 14oF
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Objective 3:
Temperature Scales
Temperature degree scales comparison
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Celsius to Kelvin: T = Tc + 273.15
Problem:
1. The lowest outdoor temperature ever
recorded on Earth is -128.6 o F.,
recorded at Vostok Station, Antarctica,
in 1983. What is this temperature on
the Celsius and Kelvin scales?
Answers: -89.22oC, 183.93 K
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• Obj. #2 - Hotter temperature means
• – more heat present in a substance
• – the faster the molecules of the
substance move.
• Obj #5 – Relate heat and temperature
• Heat units: calorie or joule (amount of
heat energy present in a substance)
• Temperature units: degree
(proportional to heat energy present in
a substance)
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21.2 Heat
If you touch a hot stove, energy enters your hand
from the stove because the stove is warmer than
your hand.
If you touch ice, energy passes from your hand
into the colder ice.
The direction of spontaneous energy transfer is
always from a warmer to a cooler substance.
The energy that transfers from one object to
another because of a temperature difference
between them is called heat.
Heat unit: calories or Joules (4.186 J/cal)
2.Describe changes in temperatures of
two objects reaching thermal equilibrium
• The temperature of the hotter substance will
decrease. The temperature of the colder
substance will increase. Each change will stop
when the temperatures are the same –
thermal equilibrium. In other words, thermal
energy travels from hot to cold.
• Obj. #4 - Heat energy can be transferred
by
– Convection (motion of fluid), conduction
(touching), or radiation (electromagnetic
waves)
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Convection
Heat transfer in fluids generally takes place via convection.
Convection currents are set up in the fluid because the
hotter part of the fluid is not as dense as the cooler part,
so there is an upward buoyant force on the hotter fluid,
making it rise while the cooler, denser, fluid sinks.
Conduction
When heat is transferred via conduction, the
substance itself does not flow; rather, heat is
transferred internally, by vibrations of atoms and
molecules.
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Radiation:
energy is transferred in the
form of electromagnetic waves.
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22.1 Conduction
Touch a piece of metal and a piece of wood in your
immediate vicinity. Which one feels colder? Which is
really colder?
• If the materials are in the same vicinity, they
should have the same temperature, room
temperature.
• The metal feels colder because it is a better
conductor.
• Heat easily moves out of your warmer hand into
the cooler metal.
• Wood, on the other hand, is a poor conductor.
• Little heat moves out of your hand into the wood,
so your hand does not sense that it is touching
something cooler.
22.1 Conduction
The good insulating properties of materials such as
wool, wood, straw, paper, cork, polystyrene, fur, and
feathers are largely due to the air spaces they
contain.
Birds fluff their feathers to create air spaces for
insulation.
Snowflakes imprison a lot of air in their crystals and
are good insulators. Snow is not a source of heat; it
simply prevents any heat from escaping too rapidly.
22.1 Conduction
Strictly speaking, there is no “cold” that passes
through a conductor or an insulator.
Only heat is transferred. We don’t insulate a home to
keep the cold out; we insulate to keep the heat in.
No insulator can totally prevent heat from getting
through it. Insulation slows down heat transfer.
22.2 Convection
Convection occurs in all fluids, liquid or
gas.
When the fluid is heated, it expands,
becomes less dense, and rises.
Cooler fluid then moves to the bottom,
and the process continues.
In this way, convection currents keep a
fluid stirred up as it heats.
22.5 Absorption of Radiant Energy
A blacktop pavement and dark automobile
body may remain hotter than their
surroundings on a hot day.
At nightfall these dark objects cool faster!
Sooner or later, all objects in thermal contact
come to thermal equilibrium.
So a dark object that absorbs radiant energy
well emits radiation equally well.
22.3 Radiation
Radiant energy is any energy that is
transmitted by radiation.
From the longest wavelength to the shortest,
this includes:
• radio waves,
• microwaves,
• infrared radiation,
• visible light,
• ultraviolet radiation,
• X-rays,
• and gamma rays.
22.3 Radiation
a. Radio waves send signals through the air.
b. You feel infrared waves as heat.
c. A visible form of radiant energy is light
waves.
22.3 Radiation
Most of the heat from a fireplace goes up the
chimney by convection. The heat that warms
us comes to us by radiation.
22.5 Absorption of Radiant Energy
Good emitters of radiant energy are also good
absorbers; poor emitters are poor absorbers.
22.5 Absorption of Radiant Energy
Because the sun shines on it, the book absorbs
more energy than it radiates.
• Its temperature increases.
• As the book gets hotter, it radiates more energy.
• Eventually it reaches a new thermal equilibrium
and it radiates as much energy as it receives.
• In the sunshine the book remains at this new
higher temperature.
22.6 Newton’s Law of Cooling
An object hotter than its surroundings
eventually cools to match the surrounding
temperature.
Its rate of cooling is how many degrees its
temperature changes per unit of time.
The rate of cooling of an object depends on
how much hotter the object is than the
surroundings.
Specific heat capacity
The amount of energy that must be added to
raise the temperature of a unit mass of a
substance by one temperature unit.
• Units: cal/goC
• For Water: 4.186 J/goC
• For Aluminum: 0.900 J/goC
• Which one heats faster?
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21.6 Specific Heat Capacity
think!
Which has a higher specific heat capacity—
water or sand? Explain.
21.6 Specific Heat Capacity
think!
Which has a higher specific heat capacity—
water or sand? Explain.
Answer:
Water has a greater heat capacity than sand.
Water is much slower to warm in the hot sun
and slower to cool at night. Sand’s low heat
capacity, shown by how quickly it warms in
the morning and how quickly it cools at night,
affects local climates.
Table 21.1
Page 413
Specific Heat Capacities
Material
J/goC
cal/goC
Water
4.187
1.00
Aluminum
0.900
0.215
Clay
1.4
0.33
Copper
0.387
0.092
Olive Oil
1.97
0.471
Steel (iron)
0.448
0.107
A table similar to this is on page 413 of the
Lead
0.128
0.031
Hewitt book.
Not shown: Clay (14,000), Olive Oil (19,700),
Silver
0.23
0.056
and Steel (448)
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• Heat Transfer Heat Transfer
Q = mCΔT = mC (Tf – Ti)
• Q = mcΔT = mc(Tf – Ti)
•
•
•
•
Q, quantity of heat in joule
m, mass of substance in g
c, specific heat for water in 4.186 J/goC
t, temperature in Celsius
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See table 12-1 on page 318
Find the amount of heat needed to change the
temperature of 5.0 g of liquid water from 8.0oC to
100oC.
Q = mcDt = 5.0g(4.186 J/goC) (92oC) = 1.9 x 103 J
Again,
specific heat is the amount of
heat necessary to change one g of
a substance 1 degree Celsius or
Kelvin.
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• 12/3 When you turn on the hot water to
wash dishes, the water pipes have to heat
up. How much heat is absorbed by a
copper water pipe with a mass of 2.3 kg
when its temperature is raised from 20.0oC
to 80.0oC?
• Q = mcDt
• Q = (2300g)(0.386J/goC)(60.0oC)
• Q = 53268 J or 5.3x104 J
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Q = mc Dt
Dt = Q
mc
Dt =
836,000 J goC
20,000g (4.186 J)
Dt = 9.98oC
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Specific Heat Capacities
Which
one is
greatest
that you
use
everyday?
OR
J/kgoC
40
Law of Heat Exchange
Q lost = Q gained
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Emily is testing her baby’s bath water and finds
that it is too cold so she adds some hot
water. If she adds 2.00 kg of water at 80.0oC
to 20.0kg of bath water at 27.0 oC, what is
the final temperature of the bath water?
Q lost = Q gain
2.00kg(1cal)(80oC-tf) = 20.0kg(1cal)(tf-27oC)
goC
goC
160 – 2tf = 20 tf – 540
-22 tf = -700
tf = 31.8oC
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Latent Heat is energy transferred
during phase changes
Crystalline materials change phase -- melt and
freeze or vaporize and condense -- at a single,
fixed temperature. Energy is required for a
change of phase of a substance. It is called
latent heat because there is no change or
difference in temperature.
Latent heat of fusion Lf describes the heat necessary
to melt (or freeze) a unit mass of a substance.
Q = m Lf
Latent heat of vaporization Lv describes the heat
necessary to vaporize (or condense) a unit mass of a
substance.
Q = m Lv
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Temperature vs Heat
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Formulas
•
•
•
•
•
•
Temperature change use: Q = mc D t
Melting or freezing: Q = m Lf
Evaporation or Condensation: Q = mLv
Lf is latent heat of fusion, 80 cal/g
Lv is latent heat of vaporization, 540 cal/g
These values are for water.
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How much heat must be gained by 0.100 kg
of ice at 0 oC in order for all of it to melt?
Q = m Lf
Q = 100g (80cal)
g
Q = 8000 or 8.00 x 103 g
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Q = mcDt + mLf
Specific Heat of Ice = 0.49 cal
g oC
Q = 100g(0.49cal)(20.0 oC) + 100g(80cal)
g oC
g
Q = 8980 = 8.98 x 103 cal
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Problem
• A jar of tea is placed in sunlight until it
reaches an equilibrium temperature of
32oC. In an attempt to cool the liquid to
0oC, which has a mass of l80 g, how much
ice at 0oC is needed? Assume the specific
heat capacity of the tea to be that of pure
liquid water.
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m tea = 180gm ice = ? which is the mass of the water that
has melted
c tea = c water = 4186 J/kgoC
H f = 3.33 x 105 J/kg and t final = 32oC
Q lost = Q gained tea loses and water gains only melting
the ice
(mcDt)tea = (mHf)ice
m ice
= (mcDt)tea
Hf ice
m ice
= (.180kg)(4186 J/kgoC)(32oC)
3.33x105 J/kg
= 7.2 x 10-2kg
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1st Law of Thermodynamics
DU = Q - W
The change in thermal energy of
an object is equal to the heat
added to the object minus the
work done by the object.
See Fig 12-11 on
Page 326
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A heat engine
Transforms heat at high
temperature into mechanical energy
and low-temperature waste heat
A heat pump (refrigerator) absorbs
heat from the cold reservoir and
gives off heat to the hot reservoir.
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2nd Law of Thermodynamics
• In the 19th Century French engineer
Sadi Carnot studied the ability of
engines to convert thermal energy into
mechanical energy.
• He developed a logical proof that even
an ideal engine would generate some
waste heat.
• Carnot’s result is best described by the
term entropy which is the measure of
the disorder in a system.
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The change is entropy, D S, is shown by the
equation:
DS = Q / T
Units: J/K
The change in entropy of an object is equal to
the heat added to the object divided by the
temperature of the object.
Natural processes occur in a direction that
increases the entropy of the universe. All
processes tend toward disorder unless some
action occurs to keep them ordered. i.e., heat
can flow only from hot to cold naturally.
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Practice Problems
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12/32 answer
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•The End
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