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Thermal Physics

Thermal Physics. Change of State (Phase) of Matter. There are 3 states (phases) of matter: Solids, liquids and gases When the temperature of a solid is increased enough, it will become a liquid When temperature is increased even more, it becomes a gas. Phase Changes.

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Thermal Physics

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  1. Thermal Physics

  2. Change of State (Phase) of Matter • There are 3 states (phases) of matter: • Solids, liquids and gases • When the temperature of a solid is increased enough, it will become a liquid • When temperature is increased even more, it becomes a gas.

  3. Phase Changes • To change phase, we must add or remove ENERGY in the form of heat • When adding heat energy… • If phase DOES NOT change, temperature and KE increase • If phase DOES change, temperature remains constant and PE increases

  4. Phase Change and Temperature • Heat transferred during a change of state doesn't change the temperature.

  5. Concept Check • When adding heat energy… • When does temperature increase? • When does temperature remain constant?

  6. Thermal Energy

  7. Kinetic Theory What is the kinetic theory of matter? • Kinetic theory explains the behavior of solids, liquids, and gases. • ALL particles of matter are in constant random motion. • We don’t observe the movement directly, but we can measure it indirectly • Pressure and temperature (macroscopic) are indicators of particle movement (microscopic)

  8. Kinetic Theory: “big indicates little” Macroscopic properties: big, easy to observe (like temperature and pressure) Microscopic properties: small, difficult to observe (like motion, velocity, and momentum of atoms) http://home.roadrunner.com/~enloephysics/macro_micro.gif

  9. Temperature • The measure of an object’s average kinetic energy. • how hot or cold something is, with respect to a standard (temperature scale – ⁰F, ⁰C, or K)) • Measured with a thermometer

  10. Temperature and Molecular Kinetic Energy • All the particles in a given object have different KE’s. • Temperature measures the average KE of all the particles in an object. • So… • Higher KE = Higher Temperature • Lower KE = Lower Temperature

  11. Temperature Scales and Water • Celsius Scale - • Freeze at 0⁰ • Boil at 100⁰ • Fahrenheit Scale – • Freeze at 32⁰ • Boil at at 212o • Kelvin Scale – • Freeze at 273 K • Boil at 373 K Freezing is freezing is freezing… 0 ⁰C = 32 ⁰F = 273 K 100 ⁰C = 212 ⁰F = 373 K

  12. Concept Check • What is temperature?

  13. Pressure • A measure of the force of particle collisions over the surface of a container • Pressure = Force/Area • P = F/A (units: N/m2)

  14. Pressure and Molecular Kinetic Energy • Pressure increases when the average kinetic energy (temperature) of the particles increase • More KE means… • More velocity, and therefore more momentum • Particles hit with more force and are moving faster

  15. Observations… • As the volume changes: • Describe how pressure changes • Describe how the motion of the molecules change

  16. Concept Check • How are pressure and molecular motion related?

  17. Heat • Heat is the amount of energy transferred between two objects as a result of differences in temperature • represented by “Q” • Measured in Joules (J) or calories (c) • Direction of energy flow is always from hot to cold • +Q – heat has been absorbed (gained) • - Q – heat has been lost

  18. Thermal Equilibrium • Thermal energy transfers between two objects until the reach the same temperature • Thermal equilibrium occurs when the average kinetic energy of the atoms and molecules is the same

  19. Specific Heat • When heat flows into an object, thermal energy increases, increasing the temperature • Increase depends on the material and its mass • Different materials require different amounts of heat to change temperature

  20. Specific Heat • The amount of energy that must be added to a material to raise the temperature of a unit mass of the material by one unit temperature • Measured in J/g●K or J/g●ºC • Kelvin or Celsius, doesn’t matter… the magnitude of each degree is the same.

  21. Concept Check • What is specific heat?

  22. Thermal Expansion Increase in the size due to an increase in temperature If temp then size Happens in most solids, liquids & gases Water is an exception – it expands as it becomes a solid!

  23. Thermal Expansion • Materials expand when heated • Materials contract when cooled • Expansion joints • Ring sizes

  24. Discuss & write… • Describe Use the kinetic theory of matter to describe the effect of cold winter weather on the air pressure inside a car tire. • Describe Cake batter and bread dough can “rise” dramatically during baking. Use the kinetic theory of matter to describe one way in which the heat of an oven can help to produce this increase in size. • Describe Use the kinetic theory of matter to describe why the high specific heat of water makes it an ideal substance for thawing frozen food or for cooling down overheating machinery.

  25. Calculating Heat Changes • Q = mc∆T • Q = heat gained(+) or lost(-) • m = mass (g) • c = specific heat of material (J/g●K or J/g●ºC) • ∆T = temperature change (K or ºC) • IMPORTANT: Notice that mass is measured in GRAMS!!!!!

  26. Example 1 • How much heat is absorbed by 60 g of copper (c = 0.386 J/gºC) when its temperature is raised from 20ºC to 80ºC ? ANS: 1389.6 J

  27. Example 2 • A child is given a 175 g silver spoon which she promptly puts in her mouth. The spoon was initially at room temperature (20°C) and the child’s mouth is 37°C. If 684 Joules of energy is gained by the spoon, what is the specific heat of silver?

  28. Law of Heat Exchange • Remember the Law of Conservation of Energy • The sum of heat loss and heat gain in a closed system is zero. • When 2 bodies of unequal temp. are mixed, the cold body absorbs heat from the warm body (loses heat) until an equilibrium temperature is reached. • Qloss + Qgain = O

  29. Example 3 • If 30 grams of water (c = 1 cal/g°C) at 12⁰C is mixed with 80 grams of water at 88 ⁰C , what will the final temperature be?

  30. Heat Transfer & Thermodynamics Chapters 22 & 24

  31. Conduction Conduction is explained by collisions between atoms or molecules, and the actions of loosely bound electrons. • When the end of an iron rod is held in a flame, the atoms at the heated end vibrate more rapidly. • These atoms vibrate against neighboring atoms. • Free electrons that can drift through the metal jostle and transfer energy by colliding with atoms and other electrons.

  32. Convection Convection occurs in all fluids, liquid or gas. When the fluid is heated, it expands, becomes less dense, and rises. Cooler fluid then moves to the bottom, and the process continues. In this way, convection currents keep a fluid stirred up as it heats.

  33. Radiation How does the sun warm Earth’s surface? It can’t be through conduction or convection, because there is nothing between Earth and the sun. The sun’s heat is transmitted by another process. Radiation is energy transmitted by electromagnetic waves. Radiation from the sun is primarily light.

  34. Heat Transfer

  35. First Law of Thermodynamics • The net heat put into a system is equal to the change in internal energy of the system plus the work done by the system. • Basically, the Law of Conservation of Energy restated.

  36. First Law of Thermodynamics If we add heat energy to a system, the added energy does one or both of two things: • increases the internal energy of the system if it remains in the system • does external work if it leaves the system So, the first law of thermodynamics states: Heat added = increase in internal energy + external work done by the system

  37. First Law of Thermodynamics James Joule’s Paddle Wheel Apparatus As the weights fall, they give up potential energy and warm the water accordingly. This was first demonstrated by James Joule, for whom the unit of energy is named.

  38. First Law of Thermodynamics • Q = U + W • Q = Heat added to a system • U = Increase in internal energy • W = Work done by (+) or on (-) the system • Always measured in Joules

  39. Example • A quantity of gas in a cylinder receives 1200J of heat from a hot plate. At the same time 600J of work are done on the gas by outside forces pressing down on a piston. Calculate the change in internal energy of the gas.

  40. Second Law of Thermodynamics • Natural processes go in a direction that maintains or increases the total entropy of the universe. • Recall: Entropy – • Measure of disorder; the more entropy, the higher the temperature. • Simply, heat flows from high to low temperature

  41. Heat Engines When heat energy flows in any heat engine from a high-temperature place to a low-temperature place, part of this energy is transformed into work output.

  42. Entropy Entropy is the measure of the amount of disorder in a system. Disorder increases; entropy increases.

  43. Entropy • Gas molecules escaping from a bottle move from a relatively orderly state to a disorderly state. • Organized structures in time become disorganized messes. Things left to themselves run down. • When a physical system can distribute its energy freely, entropy increases and energy of the system available for work decreases.

  44. Entropy The laws of thermodynamics are sometimes put this way: • You can’t win (because you can’t get any more energy out of a system than you put in). • You can’t break even (because you can’t even get as much energy out as you put in). • You can’t get out of the game (entropy in the universe is always increasing).

  45. Absolute Zero • A theoretical temperature at which no further thermal energy can be removed from an object • Usually shown as -273ºC or 0 K • Kelvin Scale is based on Absolute zero • “No system can reach absolute zero” – Third Law of Thermodynamics

  46. Absolute Zero The absolute temperatures of various objects and phenomena.

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