02-20-08 - Thermodynamics
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02-20-08 - Thermodynamics

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02-20-08 - Thermodynamics 02-20-08 - Thermodynamics Presentation Transcript

  • Thermodynamics Chapter 12
  • Review of Energy
    • Kinetic Energy
      • Energy associated with motion
    • Potential Energy
      • Many types of potential energy (gravitational, chemical, etc)
      • Energy at rest
      • Stored energy
    • Units of Energy
      • joules (J)
  • Temperature and Heat
    • The temperature of a hot cup of coffee left sitting on the table will fall until it also reaches thermal equilibrium with the air temperature in the room.
    • When a soda can is taken out of the refrigerator and left on the kitchen table, its temperature will rise – rapidly at first but then more slowly – until the temperature of the soda equals that of the air in the room. At this point, the soda and the air temperature in the room are in thermal equilibrium .
    • The change in temperature is due to the transfer of energy between object and the environment.
  • Temperature and Heat
    • Thermal energy : the total potential and kinetic energy associated with the random motion and arrangement of the particles of a material.
    • Heat, Q , is thermal energy that is absorbed, given up, or transferred from one body to another.
      • Heat is thermal energy in motion.
      • Heat is used when the transfer of thermal energy from one body to another body at a different temperature is involved.
    • (a) Q is negative when heat energy is transferred to the environment from the system.
    • (b) Q = 0 J when the transfer of heat energy between the system and the environment is equal.
    • (c) Q is positive when heat energy is transferred to a system from the environment.
  • Temperature
    • Central concept of thermodynamics is temperature.
      • Our “temperature sense” is often unreliable.
      • On a cold winter day, an iron railing seems much colder to the touch than a wooden fence post, even though both are at the same temperature.
      • This error in perception results because the iron removes energy from our fingers more quickly than the wood does.
  • Temperature
    • Temperature is a measure of the average kinetic energy of all the particles within an object.
  • Temperature & Kinetic Energy
    • The temperature of a substance will increase if the average kinetic energy of its particles is increased .
    • If the average kinetic energy of particles decreases , so does the temperature of the substance.
    • Specific heat Every substance gains or loses heat based on it’s identity. This physical property of the substance is called the specific heat capacity of the object.
    • The specific heat capacity, C , of a solid or liquid is defined as the heat required to raise a unit mass of the substance by one degree of temperature.
  • Define: Specific Heat
    • Amount of energy required to raise the temperature of 1kg by 1 o C
    Heat Energy =(mass)x specific x change heat in temp OR Q = m c  t
  • Heat Change
    • To determine the amount of thermal energy gained or lost by a mass:
    • Heat energy is gained if Q is positive.
    • Heat energy is lost if Q is negative.
  • Law of Heat Exchange
    • For a closed system in which heat energy cannot enter or leave, the heat lost by objects at a higher temperature is equal to the heat gained by objects at lower temperature until thermal equilibrium is reached (at which point the final temperature of both objects is the same).
    • The final temperature will be somewhere between the initial low temperature and the initial high temperature.
  • Law of Heat Exchange
    • Conservation of Energy :
    • Q lost = Q gained
    • To avoid problems with signs, for
    • Q lost = Q gained problems,
    • it is best to make  T = T hi – T lo
  • Heats of Transformation
    • When energy is absorbed as heat by a solid or liquid, the temperature of the object does not necessarily rise.
    • The thermal energy may cause the mass to change from one phase, or state, to another.
    • The amount of energy per unit mass that must be transferred as heat when a mass undergoes a phase change is called the heat of transformation, L.
  • Phase Changes
  •  
  • Thermal Expansion of Solids
    • Solids expand when heated and contract when cooled (with a few exceptions).
      • Heated solids increase or decrease in all dimensions (length, width, and thickness).
      • When a solid is heated, the increase in thermal energy increases the average distance between the atoms and molecules of the solid and it expands.
  • Thermal Expansion of Solids
    • Thermal expansion can be explained on a molecular basis.
    • Picture the interatomic forces in a solid as springs, as shown in the picture on the right.
    • Each atom vibrates about its equilibrium position. When the temperature increases, the amplitude and associated energy of the vibration also increase.
  • Examples of Uses of Thermal Expansion
    • Dental materials used for fillings must be matched in their thermal expansion properties to those of tooth enamel, otherwise consuming hot drinks or cold ice cream would be painful.
    • In aircraft manufacturing, rivets and other fasteners are often cooled using dry ice before insertion and then allowed to expand to a tight fit.
    • You can loosen a tight metal jar lid by holding it under a stream of hot water. Both the metal of the lid and the glass of the jar expand as the hot water adds energy to their atoms. With the added energy, the atoms can move a bit farther from each other than usual, against the interatomic forces that hold every solid together. However, because the atoms in the metal move farther apart than those in the glass, the lid expands more than the jar and is loosened.
    • Expansions slots are often placed in bridges to accommodate roadway expansion on hot days. This prevents buckling of the roadway. Driveways and sidewalks have expansion slots for the same reason.
    • Ex. A 4.0 kg sample of glass heated from 1 o C to 41 o C, and was found to have absorbed 32 J of energy. What is the specific heat of the glass?
    • Q = C x m x Δ T
    • Δ T = 41 o C – 1 o C = 40 o C
    • 32 J = C (4.0 kg)(40 o C)
    • 32 J = C (160 kg o C)
    • C = 0.2 J/kg o C
  • Your Turn!
    • Determine the specific heat of a material if a 35 g sample absorbed 48J as it was heated from 293K to 313K.
    • If 980 kJ of energy are added to 6.2 L of water at 291 K, what will the final temperature of the water be?
      • 1 kJ = 1000 J
      • 1 L of water has 1 kg of mass
      • C water = 4180 J/kg °K