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Chapter 13

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Chapter 13

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  1. Chapter 13 Heat Energy and Its Effects

  2. Heat Transfer • When two objects A and B of different temperatures (TA >TB ) are placed in thermal contact, the temperature of the warmer decreases and the temperature of the cooler increases • Heat flow (transfer) from object A to Object B. • The heat transfer ceases when the objects reach thermal equilibrium • Heat flow is really energy transfer

  3. Heat • Heat is the transfer of energy between a system and its environment because of a temperature difference between them • The symbol Q is used to represent the amount of heat between a system and its environment

  4. Gas as Example • In a monatomic gas, the KE is the only type of energy the molecules can have • root-mean-square (rms) speed

  5. Units of Heat • Calorie • An historical unit, before the connection between thermodynamics and mechanics was recognized • A calorie is the amount of energy necessary to raise the temperature of 1 g of water from 14.5° C to 15.5° C . • A Calorie (food calorie) is 1000 cal

  6. Units of Heat, cont. • US Customary Unit – BTU • BTU stands for British Thermal Unit • A BTU is the amount of energy necessary to raise the temperature of 1 lb of water from 63° F to 64° F • 1 cal = 4.186 J • This is called the Mechanical Equivalent of Heat 1 kcal = 1 Cal = 4186 J 1 BTU = 1054 J =778 ft-lb

  7. James Prescott Joule • 1818 – 1889 • British physicist • Conservation of Energy • Relationship between heat and other forms of energy transfer

  8. Examples • Work done against friction in a certain bearing is 0.4 J/s. How much heat energy develops? • How high does a person (70kg) have to walk up a hill to burn off 1 Cal?

  9. Specific Heat (Capacity) • Every substance requires a unique amount of energy per unit mass to change the temperature of that substance by 1° C • The specific heat, c, of a substance is a measure of this amount

  10. Units of Specific Heat • SI units • J / kg °C • Historical units • cal / g °C Water: c= 1 cal/(gram·°C)=4.186 J /(gram·°C) = 1BTU/(lb ·°F) Copper: c=0.093 cal/(gram·°C) Human body: 0.83 Alcohol: 0.58 Glass: 0.20 Aluminum: 0.22

  11. Heat and Specific Heat • Q = m c ΔT • ΔT is always the final temperature minus the initial temperature • When the temperature increases, ΔT and ΔQ are considered to be positive and energy flows into the system • When the temperature decreases, ΔT and ΔQ are considered to be negative and energy flows out of the system

  12. A Consequence of Different Specific Heats • Water has a high specific heat compared to land • On a hot day, the air above the land warms faster • The warmer air flows upward and cooler air moves toward the beach

  13. Phase Changes • A phase change occurs when the physical characteristics of the substance change from one form to another • Common phases changes are • Solid to liquid – melting • Liquid to gas – boiling • Phases changes involve a change in the states, but no change in temperature

  14. Latent Heat • During a phase change, the amount of heat is given as • Q = ±m H • H is the latent heat of the substance • Latent means hidden • H depends on the substance and the nature of the phase change • Choose a positive sign if you are adding energy to the system and a negative sign if energy is being removed from the system

  15. Latent Heat, cont. • SI units of latent heat are J / kg • Latent heat of fusion, HF, is used for melting or freezing • Latent heat of vaporization, Hv, is used for boiling or condensing • Look in Handbook for the latent heats for various substances

  16. Sublimation • Some substances will go directly from solid to gaseous phase • Without passing through the liquid phase • This process is called sublimation • There will be a latent heat of sublimation associated with this phase change

  17. Graph of Ice to Steam

  18. Warming Ice (c=0.5 cal/g·°C) • Start with one gram of ice at –30.0º C • During A, the temperature of the ice changes from –30.0º C to 0º C • Use Q = m c ΔT • Will add 62.7 J of energy

  19. Melting Ice • Once at 0º C, the phase change (melting) starts • The temperature stays the same although energy is still being added • Use Q = m HF • Needs 333 J of heat • HF =333 J/g = 80 cal/g

  20. Warming Water (c=1 cal/g·°C) • Between 0º C and 100º C, the material is liquid and no phase changes take place • Energy added increases the temperature • Use Q = m c ΔT • 419 J of energy are added

  21. Boiling Water • At 100º C, a phase change occurs (boiling) • Temperature does not change • Use Q = m Hv • 2 260 J of energy are needed • Hv=2260J/g =539 cal/g

  22. Heating Steam (c=0.48 cal/g·°C) • After all the water is converted to steam, the steam will heat up • No phase change occurs • The added energy goes to increasing the temperature • Use Q = m c ΔT • To raise the temperature of the steam to 120°, 40.2 J of energy are needed

  23. Remarks • Process is reversible • Similar curve for other substances • Vaporization is a cooling process, e.g. sweat evaporates • Heat of combustion, e.g. gasoline 11000cal /g.

  24. All about water • Ice: c=0.5 cal/g·°C • Water: c=1 cal/g·°C • Steam: c=0.48 cal/g·°C • HF =333 J/g = 80 cal/g • Hv=2260J/g=539 cal/g

  25. Example A restaurant serves coffee at 70 °C in copper mugs initially at 20 °C. Find the the final temperature after coffee and mug reach thermal equilibrium. (mmug=0.1 kg, mcoffee=0.2 kg)

  26. Example • How many grams of ice at –14 °C are needed to cool 200cm^3 of coke from 25 °C to 10 °C?

  27. Example • How many grams of hot steam at 115 °C are needed to heat 1000 grams of water at 20 °C to 90 °C?

  28. Example • How much heat from burning 50 grams of gasoline?