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Chapter 11

Chapter 11. Heat. Heat. Heat is a transfer of internal energy Units of Joules calorie—amount of heat needed to raise the temperature of 1 g of water by 1 °C 1 cal = 7.186 J Calorie (kilocalorie)—amount of heat needed to raise the temperature of 1 kg of water by 1 °C

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Chapter 11

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  1. Chapter 11 Heat

  2. Heat • Heat is a transfer of internal energy • Units of Joules • calorie—amount of heat needed to raise the temperature of 1 g of water by 1°C • 1 cal = 7.186 J • Calorie (kilocalorie)—amount of heat needed to raise the temperature of 1 kg of water by 1°C • 1 kcal = 1000 cal = 4186 J • British Thermal Units (Btu)—amount of heat needed to raise the temperature of 1 lb of water by 1°F Mechanical Equivalent of Heat

  3. Specific Heat • 1 molecule of water is lighter than 1 molecule of Iron • 1 kg of water contains more molecules than 1 kg of Iron • Equal amounts of heat added to both distributes among more water molecules than Iron molecules • Iron molecules have greater increase in average KE, therefore higher temperature

  4. Specific Heat • Q = Heat (J) • m = mass of substance (kg) • T = change in temperature (°C or K) • Specific heat (c)—amount of heat needed to raise the temperature of 1 kg of a substance by 1°C • Units of J/kg°C • Larger molecules tend to have smaller specific heats • “heat capacity” • Why antifreeze instead of water?

  5. Specific Heat • Calorimetry—quantitative measure of heat exchange • Needs a closed thermal system (all heat exchange is internal) • Heat lost by one part = heat gained by other part(s) • Used to determine specific heats

  6. Phase Changes • Solid Phase • Molecules are held tightly in place by attractive forces (bonds) • Liquid Phase • Molecules have enough KE to push other molecules out of the way, but not to break free of bonds • Gas Phase • Molecules have enough KE to break free from the bonds • Plasma Phase • Molecules have enough KE to “throw” all electrons off their atoms (atoms become completely ionized)

  7. Phase Changes • Melting—solid to liquid • Freezing—liquid to solid • Evaporation—liquid to gas • Condensation—gas to liquid • Sublimation—solid to gas • Deposition—gas to solid • Frost and snow

  8. Phase Changes • Latent Heat (L)—heat energy associated with phase changes • Heat energy is must be added to go from a lower KE phase to a higher KE phase • Heat energy is released when going from a higher KE phase to a lower KE phase • Latent Heat does not cause a temperature change Sweating as a cooling technique Danger of steam burns

  9. Phase Changes • Latent heat of fusion (Lf) • Melting—substance absorbs Lf • Freezing—substance releases Lf • Latent heat of vaporization (Lv) • Evaporation—substance absorbs Lv • Condensation—substance releases Lv • Latent heat of sublimation (Ls) • Sublimation—substance absorbs Ls • Deposition—substance releases Ls Table 11.2, page 368 Example 11.6 and 11.7, page 370

  10. Phase Changes

  11. Phase Changes Lower pressure, lower boiling point Higher pressure, Higher condensing point

  12. Heat Transfer • Conduction—transfer of heat energy through direct contact • Hot plate • Convection—transfer of heat energy by movement mass that contains the energy • Warm air rises, cold air descends • Radiation—transfer of heat energy via electromagnetic radiation • Sun, fire

  13. Heat Transfer • Conduction is due to molecular collisions • Thermal conductors—good conductors of heat • Metals, atoms with lots of free electrons • Thermal insulators—poor conductors of heat • Nonmetals, atoms with few free electrons

  14. Heat Transfer • In general . . . • Gases are the worst conductors • Liquids are better conductors than gases • Solids are better conductors than liquids

  15. Heat Transfer • Q/t = Rate of heat flow through an object • k = thermal conductivity of material (J/ms°C) • A = surface area (m2) • T = temperature difference between the ends (°C or K) • d = thickness of material (m) • Only works for conduction • T/d is sometimes referred to as the thermal gradient Table 11.3, page 374 Example 11.8, page 374

  16. Heat Transfer • Convection • Natural convection—due to density changes via temperature (warm air rising, cold descending) • Land breezes vs. Sea breezes • Forced convection—material is mechanically forced to move • Home furnace • Human circulatory system • Automobile cooling system

  17. Heat Transfer • Radiation • Molecule vibration causes natural emission of electromagnetic radiation • Objects radiate over a range of wavelengths • Most emission occurs in infrared portion of the spectrum • Water naturally vibrates at infrared frequencies, and will absorb this radiation • Peak emission frequency increases with temperature • Red hot, white hot

  18. Heat Transfer •  = 5.67 x 10-8 W/m2K4 • Stefan-Boltzmann constant • A = surface area (m2) • e = emissivity of material • No units, always between 0 and 1 • How good an absorber/emitter the object is • Dark surfaces are good absorbers and emitters • Ideal absorber (e = 1) referred to as a blackbody • Shiny surfaces are poor absorbers and emitters • T = absolute temperature (K)

  19. Heat Transfer • Net rate of energy loss/gain between a radiating object and its surroundings • Ts = absolute temperature of surroundings Example 11.10, page 381

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