1. Lecture Outline Chapter 11 College Physics, 7 th Edition Wilson / Buffa / Lou © 2010 Pearson Education, Inc.

2. Chapter 11 Heat © 2010 Pearson Education, Inc.

3. Units of Chapter 11 Definition and Units of Heat Specific Heat and Calorimetry Phase Changes and Latent Heat Heat Transfer © 2010 Pearson Education, Inc.

4. 11.1 Definition and Units of Heat Heat is a form of energy. It is a type of energy transfer. Energy can be added to or removed from the total internal energy of a system. S.I. Unit of Heat: Joules There are other units of heat, though; the most common one is the kilocalorie: One kilocalorie (kcal) is defined as the amount of heat needed to raise the temperature of 1 kg of water by 1 C° © 2010 Pearson Education, Inc.

5. 11.1 Definition and Units of Heat Confusingly, the calories listed on nutrition labels in the U.S. are really kilocalories (sometimes called Calories). Some other labels are more accurate (left, Australia; right, Germany). © 2010 Pearson Education, Inc.

6. 11.1 Definition and Units of Heat This figure illustrates the three most common units of heat. © 2010 Pearson Education, Inc.

7. 11.1 Definition and Units of Heat Careful experimentation using the apparatus below compared kilocalories and Joules. This relationship is called the mechanical equivalent of heat. © 2010 Pearson Education, Inc.

8. 11.1 Definition and Units of Heat At a birthday party, a student eats a piece of cake (food energy value of 200 Cal). To prevent this energy from being stored as fat, she takes a stationary bicycle workout class right after the party. This exercise requires the body to do work at an average rate of 200 Watts. How long must the student bicycle to achieve her goal of “working off” the cake’s energy?

9. 11.2 Specific Heat and Calorimetry The amount of heat needed to increase the temperature of a solid or liquid depends on the amount of the substance, the temperature change, and the properties of the substance itself. The constant c is called the specific heat of the substance. © 2010 Pearson Education, Inc.

10. 11.2 Specific Heat and Calorimetry Specific Heat Capacity (Specific Heat) is defined as the heat required to raise (or lower) the temperature of 1 kg of a substance by 1 degree Celsius. Units: J/(kg x K) or J/(kg x C) What does this really mean though??

11. 11.2 Specific Heat and Calorimetry

12. To prepare pasta, you bring a pot of water from room temperature (20 degrees Celsius) to its boiling point. The pot itself has a mass of 0.900 kg, is made of steel, and holds 3.00 kg of water. Which of the following is true: –a.) The pot requires more heat than the water –b.) The water requires more heat than the pot –c.) The require the same amount of heat Determine the required heat for both the water and the pot, and the ratio of the heat of water vs. the heat of the pot.

13. 11.2 Specific Heat and Calorimetry Calorimetry is the quantitative measurement of heat exchange; it is done using a calorimeter. How does a calorimeter work? © 2010 Pearson Education, Inc.

14. 11.2 Specific Heat and Calorimetry Students in a physics lab are to determine the specific heat of copper. They place 0.150 kg of copper into boiling water and let it stay there for awhile, so as to reach 100 degrees Celsius. Then they carefully pour the hot copper into an aluminum calorimeter cup containing 0.200 kg of water at 20 degrees Celsius. The final temperature of the mixture in the cup is 25 degrees Celsius. If the aluminum cup has a mass of 0.0450 kg, what is the specific heat of copper? (Assume no heat exchange with the environment)

15. 11.2 Specific Heat and Calorimetry Specific heat can be defined for gases as well, but gases do not have constant volume or pressure. We therefore define two specific heats for gases—one at constant volume ( c V ), and one at constant pressure ( c P ). What happens to a gas at constant volume? What happens to a gas at constant pressure? © 2010 Pearson Education, Inc.

16. 11.3 Phase Changes and Latent Heat Three phases of matter: solid, liquid, and gas. Solid has a definite shape and the strongest intermolecular bonds. Liquid flows but is relatively incompressible, so it has a definite volume. Gas is compressible, and will expand to fill a container. © 2010 Pearson Education, Inc.

17. 11.3 Phase Changes and Latent Heat Phase changes: Solid → liquid: melting Liquid → gas: evaporating, boiling Gas → liquid: condensing Liquid → solid: freezing Solid → gas: sublimating Latent heat is the amount of heat absorbed or released when a substance undergoes a phase transition. © 2010 Pearson Education, Inc.

18. 11.3 Phase Changes and Latent Heat During a phase transition, the heat energy goes to changing the intermolecular bonds, and the temperature does not change. The heat needed for a phase change is: Here, L is the latent heat; L f is the latent heat of fusion (solid ↔ liquid) and L v the latent heat of vaporization (liquid ↔ gas). Q can be +/- because… © 2010 Pearson Education, Inc.

19. 11.3 Phase Changes and Latent Heat Heat added as ice becomes steam: © 2010 Pearson Education, Inc.

20. 11.3 Phase Changes and Latent Heat Heat is added to 1.00 kg of cold ice at -10 degrees Celsius. How much heat is required to change the cold ice to hot steam at 110 degrees Celsius?

21. 11.4 Heat Transfer Heat transfer takes place via three mechanisms: 1. Conduction 2. Convection 3. Radiation We will discuss each separately. © 2010 Pearson Education, Inc.

22. 11.4 Heat Transfer Conduction is the transfer of heat through a substance. Molecules at higher temperatures move rapidly and transfer energy to less energetic molecules. Ex: Hot coffee cup; heating tea on a stove. Typically, metals are good conductors of heat—they have electrons that are free to move throughout the material—and nonmetals are not. Called Thermal Conductors © 2010 Pearson Education, Inc.

23. 11.4 Heat Transfer Nonconductors of heat are also called insulators. [There is an absence of free electrons] Specifically Thermal Insulators Examples are wood and cloth

24. 11.4 Heat Transfer The heat flow rate through a slab of material is proportional to its surface area and to the temperature difference, and inversely proportional to its thickness. The constant k is called the thermal conductivity. [Table 11.3 – page 401] S.I. Unit: J/s © 2010 Pearson Education, Inc.

25. 11.4 Heat Transfer This diagram illustrates the geometry of heat transfer by conduction. © 2010 Pearson Education, Inc.

26. 11.4 Heat Transfer Let’s consider how much heat would be conducted in 1.0 hr through a wooden (pine) ceiling without insulation, if the room measures 3.0 m by 5.0 m and is 2.0 cm thick. Also, the temperature inside the room is 20 degrees Celsius, while in the attic is 8 degrees Celsius.

27. 11.4 Heat Transfer Heat transfer in fluids is mostly by convection, which is the result of mass transfer; that is, heat is transferred as warmer fluid moves to replace cooler fluid. Convection may be spontaneous (as below) or forced. © 2010 Pearson Education, Inc.

28. 11.4 Heat Transfer Many homes are heated using forced hot air; this is an example of forced convection. © 2010 Pearson Education, Inc.

29. 11.4 Heat Transfer Foam insulation is sometimes blown into the space between inner and outer walls of a house. Since air is a better thermal insulator than foam, why is the foam insulation needed? a.) to prevent loss of heat by conduction b.) to prevent loss of heat by convection c.) for fireproofing

30. 11.4 Heat Transfer Radiation is the only type of heat transfer that can take place through a vacuum. You can feel the radiation of heat when you stand near a fireplace. This radiation is in the form of electromagnetic waves, in the infrared part of the spectrum. Transfer by electromagnetic waves! Common referred to as radiant energy Example: Sunlight © 2010 Pearson Education, Inc.

31. 11.4 Heat Transfer A good emitter of radiation is also a good absorber: © 2010 Pearson Education, Inc.