1. Preview Section 1 Temperature and Thermal Equilibrium
Section 2 Defining Heat
Section 3 Changes in Temperature and Phase

2. What do you think?
Suppose you have two cups of water. One is hot and the other is cold.
How is the cold water different from the hot water?
Describe the motion of the molecules in each.
What changes would occur if the hot water was changed into steam?
What are the common scales used to measure temperature?
When is each scale generally used?
All scales use degrees to measure temperature. Which scale has the largest degrees? Explain.
When asking students to express their ideas, you might try one of the following methods.
(1) You could ask them to write their answers in their notebook and then discuss them.
(2) You could ask them to first write their ideas and then share them with a small group of 3 or 4 students. At that time you can have each group present their consensus idea. This can be facilitated with the use of whiteboards for the groups.
The most important aspect of eliciting student’s ideas is the acceptance of all ideas as valid. Do not correct or judge them. You might want to ask questions to help clarify their answers. You do not want to discourage students from thinking about these questions and just waiting for the correct answer from the teacher. Thank them for sharing their ideas. Misconceptions are common and can be dealt with if they are first expressed in writing and orally.
Probe students’ ideas on the arrangement and motion of molecules. How do they change with temperature? What happens to molecular motion and arrangement when a change of state occurs?
See if students really understand the uses and values for different temperature scales.

3. Temperature
Temperature measures the average kinetic energy of the particles.
Average speed is used because all particles do not have the same speed, and speeds change as the particles collide.
Internal energy is the energy a substance has due to the motion of the particles (kinetic energy) and the position of the particles (potential energy).

4. Temperature
Do thermometers change the temperature of the substance being measured?
If so, how?
How can you minimize the problem?
Temperature must be measured when thermal equilibrium is reached.
Always read a thermometer after it has stopped rising or falling.
At this point, equilibrium has been reached between the thermometer and the substance.
A thermometer placed in a hot substance will cool it slightly. Generally, this is not significant, but it can lead to errors in measurement.

5. Thermometers
The expansion of mercury changes the reading in this thermometer.
How does such a small change in the volume of the mercury (see circled segments) result in such a large rise inside the thermometer?
Students should realize that the bulb filled with mercury undergoes very small changes in volume. However, the opening above the bulb is very narrow, so a small volume change produces a significant rise in the tube.
Students may be familiar with thermometers used to measure body temperature that need to be “shaken down” to return to the normal setting. Laboratory thermometers do not require this because the mercury or alcohol is free to move up and down the tube without restrictions.

6. Thermometers Calibration depends on fixed temperatures.
Three common temperature scales used:
Fahrenheit for weather and medicine (U.S.)
Celsius for work in science
Kelvin or absolute for many scientific laws

7. Comparison of Temperature Scales
Point out to students that a degree on the Kelvin scale is the same as a degree on the Celsius scale, so a change of 10°C is the same as a change of 10 K. A degree on the Celsius scale is 1.8 times larger than a degree on the Fahrenheit scale because there are 180° between freezing water (32°) and boiling water(212°) on the Fahrenheit scale, and only a 100° difference on the Celsius and Kelvin scales.
An interesting video from the mechanical universe ( explains the basis of Daniel Fahrenheit’s temperature scale ( ). The professor explains this near the end of the 30 minute video (#45 - Temperature and the Gas Laws). For an interesting assignment, have students find out how he came up with his scale. The videos are available on demand through the web site shown above. They are part of the Annenberg CPB project videos.

8. The Absolute Temperature Scale
What is meant by absolute zero?
Absolute zero = 0 K
Suppose the pressure and temperature of a gas are plotted as shown, and the graph is extrapolated to 0 K. What does this suggest about P at absolute zero ( °C or 0 K)?
P = 0 at absolute zero
Students may be familiar with the Kelvin scale from the ideal gas law in chemistry. The P-T graphs and V-T graphs for gases are straight lines that extrapolate back to (0,0) if you use the Kelvin scale, thus making the relationships simple direct proportions.
Point out to students that there is no temperature lower than absolute zero, that is why it is “absolute.” Students may ask if absolute zero has ever been attained. Laboratory experiments have never reached absolute zero, although they have reached temperatures of just a half-billionth of a degree above absolute zero.
One of the Classroom Practice problems asks students to find the Fahrenheit temperature for absolute zero (slide 11). If students ask, the Rankine scale is the absolute scale corresponding to the Fahrenheit scale.

9. Temperature Conversions
Have students solve the first equation for TC to get an equation for the conversion of Fahrenheit into Celsius.
TC = (5/9)(TF - 32)

10. Classroom Practice Problems
One day it was -40°C at the top of Mont Blanc and -40°F at the top of Mount Whitney. Which place was colder?
Answer: Neither (-40°C = -40°F)
What is the Fahrenheit temperature equivalent to absolute zero?
Answer: °F
What is the Celsius temperature on a hot summer day when the temperature is 100. °F?
Answer: 37.8°C
For problems, it is a good idea to go through the steps on the overhead projector or board so students can see the process instead of just seeing the solution. Allow students some time to work on problems and then show them the proper solutions. Do not rush through the solutions. Discuss the importance of units at every step. Problem solving is a developed skill and good examples are very helpful.
To solve each problem, simply use the conversion equations on the previous slide.

11. Now what do you think?
Suppose you have two cups of water. One is hot and the other is cold.
How is the cold water different from the hot water?
Describe the motion of the molecules in each.
What changes would occur if the hot water was changed into steam?
What are the common scales used to measure temperature?
When is each scale generally used?
All scales use degrees to measure temperature. Which scale has the largest degrees? Explain.
The molecules move faster in the hot cup, on the average. Some of the molecules in the hot cup may be moving slower than some of those in the cold cup, so we must talk about average molecular speed. The particles are farther apart when the water changes into a gas.
Scientists use Celsius and Kelvin temperature scales, depending on the circumstances. Fahrenheit is used in the U.S. for meteorology. The Celsius and Kelvin degrees, which are equal in size, are both 1.8 times larger than a Fahrenheit degree. Have students count the number of degrees between freezing water and boiling water.

12. What do you think?
Internal energy is the energy due to the kinetic and potential energy of the particles.
Does the ice water or an equal quantity of hot chocolate have greater internal energy? Why?
Which has more internal energy, a gallon of cold water or a drop of hot chocolate?
How will the internal energy of the water and hot chocolate change over time?
How will this change occur?
When asking students to express their ideas, you might try one of the following methods.
(1) You could ask them to write their answers in their notebook and then discuss them.
(2) You could ask them to first write their ideas and then share them with a small group of 3 or 4 students. At that time you can have each group present their consensus idea. This can be facilitated with the use of whiteboards for the groups.
The most important aspect of eliciting student’s ideas is the acceptance of all ideas as valid. Do not correct or judge them. You might want to ask questions to help clarify their answers. You do not want to discourage students from thinking about these questions and just waiting for the correct answer from the teacher. Thank them for sharing their ideas. Misconceptions are common and can be dealt with if they are first expressed in writing and orally
Students will probably agree that the hot chocolate has the greater amount of internal energy. It is important that they eventually realize that the internal energy does depend on the quantity. If students are unclear about this now, be sure to return to this point on the concluding slide.
Students may not have a clear understanding of the process that causes one substance to lose energy while the other substance gains energy. Try to get them to explain the process on a molecular level. That is what we are measuring with internal energy.

13. Internal Energy and Heat
Internal Energy (U) is the energy contained within the particles of a substance.
Heat (Q) is the internal energy transferred between objects.
Heat always moves from a higher-temperature object to a lower-temperature object.
The rate of transfer depends on the difference in temperature.
The greater the temperature difference, the greater the rate of energy transfer (if other factors are equal).
Point out to students that heat is not the same as internal energy; heat is internal energy that moves from a hotter substance to a colder substance.

14. Internal Energy and Heat
Which way does heat flow if you place a warm canned beverage in cold water?
How does this occur, on a molecular level?
Answers:
from the beverage into the water
Molecules in the beverage are moving faster, so they strike the molecules in the can. This causes the molecules in the can to move faster. The molecules in the beverage slow down as they lose energy. The molecules in the can then strike the molecules of the liquid, causing them to move faster. Eventually, the beverage, can, and water will reach the same temperature. The can and beverage will be colder, and the water will be warmer.

15. Heat has Units of Energy (joules)
Historically:
The calorie is the heat required to raise the temperature of 1.0 g of water by 1.0°C at a specific temperature (but nearly true no matter what the temperature).
The BTU is the heat required to raise the temperature of 1.0 pound of water by 1.0°F at a specific temperature (but nearly true no matter what the temperature).
Also note that kcal are the same as dietary Calories. Therefore, the energy contained in 1.0 Calories from food will increase the temperature of 1000 g of water by 1°C.
Some texts still refer to calories and kcal as units of heat. This text will use joules for all forms of energy.

16. Heat Transfer
In what three ways can internal energy be transferred from a hot object to a colder object?
Conduction is the transfer of heat through a substance by molecule to molecule contact.
Metals are good conductors.
Styrofoam is a good insulator.
Convection is the transfer of energy by the movement of a fluid.
Hot air in a room rises and cold air moves in to replace it.
Radiation is the transfer of energy by electromagnetic waves.
No matter is transferred, only energy.
Use the following examples to probe students about each method of energy transferring before revealing the related bullet points.
First, ask students how heat gets from their hand into the liquid in a can of soda they are holding. See if they come up with the term conduction.
Next ask students how the heat from a campfire can warm your hand if it is held high above the flame. There is no contact, so conduction is not the answer. They may say radiation. This would be true above or beside the fire. Above the flame, you also have the warm air rising toward your hand.
Finally, ask students how the sun’s energy reaches Earth. Radiation is heat transfer through electromagnetic waves. The sun’s energy is transferred to Earth in this way. The sun loses energy as the radiation is sent out into space, and Earth gains energy as it absorbs this radiation.

17. Heat and Work Work can be changed into internal energy (U).
Rub your hands together and you’ll feel the increase in internal energy produced by your work.
Pull a nail from a piece of wood and the nail is hot.
Mechanical energy (PE + KE) is conserved when there is no friction.
Total energy, including internal energy, is always conserved.

18. Energy Conservation
Any loss of one type is balanced by a gain in the other types of energy.
Predict the sign (+, -, or 0) for the change in each quantity when:
A child slides down a plastic playground slide
A car applies the brakes to stop on a level road
Example 1: KE increases (+), PE decreases (-), and U increases (+) because friction warms the slide and the child. The total change for all three is zero.
Example 2: KE decreases (-), PE is unchanged (0), and U increases (+) due to friction. The total change for all three is zero.

19. Now what do you think?
Does the ice water or an equal quantity of hot chocolate have greater internal energy? Why?
Which has more internal energy, a gallon of cold water or a drop of hot chocolate?
How will the internal energy of the water and hot chocolate change over time?
How will this change occur?
For equal quantities, the hot chocolate has greater internal energy because the particles have greater kinetic and potential energy. However, a gallon of water has more internal energy than a drop of hot chocolate, because internal energy also depends on quantity.
The hot chocolate will lose internal energy while the ice water gains internal energy. The hot chocolate will transfer energy to the air when the molecules in the liquid strike the air molecules and transfer KE to them. Similarly, the air molecules are moving faster than the ice-water molecules. Thus, collisions between the two will cause the ice-water molecules to move faster and the air molecules to slow down.

20. What do you think?
What property of water makes it so useful as a coolant in automobiles, nuclear reactors, and other machinery?
How does it differ from other liquids regarding its ability to cool substances?
Why do you feel cool when getting out of a warm swimming pool on a hot day?
How do you feel if it is windy out? Why?
How do you feel if it is an indoor pool? Why?
When asking students to express their ideas, you might try one of the following methods.
(1) You could ask them to write their answers in their notebook and then discuss them.
(2) You could ask them to first write their ideas and then share them with a small group of 3 or 4 students. At that time you can have each group present their consensus idea. This can be facilitated with the use of whiteboards for the groups.
The most important aspect of eliciting student’s ideas is the acceptance of all ideas as valid. Do not correct or judge them. You might want to ask questions to help clarify their answers. You do not want to discourage students from thinking about these questions and just waiting for the correct answer from the teacher. Thank them for sharing their ideas. Misconceptions are common and can be dealt with if they are first expressed in writing and orally.
Students may attribute water’s cooling properties to its ability “conduct” heat. Water, however, is not a very good conductor. See if students make statements regarding the ability of water to absorb heat without increasing its temperature a great deal. They may also discuss the cost of water and its availability, which are important factors as well. Try to lead the discussion back towards a discussion of heat and thermal energy.
Try to elicit as many ideas as possible regarding the cool sensation when getting out of a pool. Students may know that water evaporation is a cooling process. If they do, ask them why evaporating water would cool you off. Does water evaporating from a freshly-mopped floor cool the floor as well?

21. Specific Heat Capacity
Specifc heat capacity (cp) measures the amount of heat required to raise the temperature by 1°C for 1 kg of a substance.
It is different for every substance.
SI Units: J/kg•°C
Alternate form of this equation: Q = cpmT

22. Suppose each metal shown above absorbs 100 J of energy.
Which will show the greatest increase in temperature? the least?
How does water compare to iron with regard to heat capacity?
The temperature of aluminum would rise the least, and the temperature of lead would rise the most. Students need to realize that a high heat capacity means that the temperature will rise more slowly as energy is absorbed because the substance has a high capacity for absorbing heat.
Water has a heat capacity that is more than 9 times greater than iron. For a given amount of heat, the temperature of the iron will go up 9 times as much as the water. Point out that water therefore makes a nice coolant for cast-iron engine blocks, and for machinery made of aluminum and iron. It can absorb a large quantity of heat without a large increase in temperature.

23. Calorimetry
Calorimetry is a problem-solving approach to heat transfer problems.
Conservation of energy
Qgained = -Qlost
A calorimeter is an insulated cup with water used for the experiment.
Ignore heat gained or lost by the cup, as it is small.
Discuss the calorimeter shown. In calorimetry experiments, the object is placed in the water, and the temperature of the water increases or decreases. Since energy is conserved, students can solve for the unknown on the other side of the equation. (The next two slides gives practice problems for this concept.)
In extremely precise experiments you must consider the heat gained by the cup, the stirrer, and the thermometer. However, the stirrer is very light and only the inner layer of the cup undergoes a change in temperature because the cup is well insulated.

24. Classroom Practice Problem
In a student’s experiment, a kg metal ball is placed in a calorimeter filled with kg of water at 21.0°C. The initial temperature of the ball is 98.5°C. After reaching equilibrium, the temperature is 27.3°C. Find the specific heat capacity of the metal and use the table to determine the type of metal.
Answer: 444 J/kg•°C , very close to iron
For problems, it is a good idea to go through the steps on the overhead projector or board so students can see the process instead of just seeing the solution. Allow students some time to work on problems and then show them the proper solutions. Do not rush through the solutions. Discuss the importance of units at every step. Problem solving is a developed skill and good examples are very helpful.
Conservation of energy
Heat gained by the water (Qw) = -Heat lost by the metal (Qm)
cwmw Tw = cm mm Tm
(4.186 x 103 J/kg•°C) x kg x (27.3° °C) = -[cm x kg x (27.3°- 98.5°C)]
Solving for cm yields 444 J/kg•°C, which is close to iron.

25. Classroom Practice Problem
A bathtub has 20.0 kg of water at 60.0°C and the bather wants the temperature to be 30.0°C. How much 20.0°C water must be added to the bath water to achieve the desired temperature?
Answer : 60.0 kg
For problems, it is a good idea to go through the steps on the overhead projector or board so students can see the process instead of just seeing the solution. Allow students some time to work on problems and then show them the proper solutions. Do not rush through the solutions. Discuss the importance of units at every step. Problem solving is a developed skill and good examples are very helpful.
This is a conservation of energy problem similar to the last one. The unknown in this case is the mass of the cold water.
Qc = -Qh
The heat capacity of water is the same on both sides so it cancels out, leaving:
m x ( )°C = -[20.0 kg x (60°C - 30°C)]
m = 60.0 kg

26. Latent Heat Latent heat is heat gained or lost during phase changes.
When substances melt, freeze, boil, condense, or sublime, the temperature does not change during the phase change.
Heat absorbed changes the potential energy of the particles.

27. Latent Heat
Heat of fusion (Lf) is the heat required to melt 1 kg of a substance.
Also equals the heat released when 1 kg freezes
Which graph segment represents this?
Heat of fusion is segment B, and heat of vaporization is segment D. The other portions show the increase in temperature (KE) as the heat is absorbed.
Segment D has a broken line because the heat of vaporization for water is so large the line would have been very long.
You can refer to Table 5 in the text to show students how the heat was evaluated for each segment.
Heat of vaporization (Lv) is the heat required to change 1 kg of a substance from a liquid to a gas.
Which graph segment represents this?

28. Latent Heats of Fusion and Vaporization
Note the high heat of vaporization for water (2.26 million J/kg). The evaporation of water requires large amounts of heat. This is why swimming pools get colder as the water evaporates from the surface. The surface molecules absorb heat from the molecules below and evaporate, leaving those molecules left behind with less energy.

29. Now what do you think?
What property of water makes it so useful as a coolant in automobiles, nuclear reactors and other machinery?
How does it differ from other liquids regarding its ability to cool substances?
Why do you feel cool when getting out of a warm swimming pool on a hot day?
How do you feel if it is windy out? Why?
How do you feel if it is an indoor pool? Why?
Water has a very high specific heat capacity. Water can absorb more heat than nearly all liquids without large increases in temperature.
Evaporation is a phase change that removes heat. The latent heat of vaporization for water is 2.26 x 106 J/kg, so every gram of water that evaporates from your skin requires 2.26 x 103 J of energy. This cools your skin. It is faster on a windy day, so you feel colder. Indoors, evaporation is slowed because of the high humidity, so you do not feel as cold as you would outdoors.