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Heat & Thermodynamics

Heat & Thermodynamics. What is the Difference Between Heat and Temperature?. Both are related to energy but there’s a big difference. Temperature. Measure of speed of particles (kinetic energy) Measured by thermometers Work by expansion of a liquid Other types use bimetallic strip.

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Heat & Thermodynamics

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  1. Heat & Thermodynamics

  2. What is the Difference Between Heat and Temperature? • Both are related to energy but there’s a big difference

  3. Temperature • Measure of speed of particles (kinetic energy) • Measured by thermometers • Work by expansion of a liquid • Other types use bimetallic strip

  4. Digital Thermometers Use “thermistors” - temperature dependant semiconductor resistors

  5. Temperature Scales • Fahrenheit T(0F) = (1.8x 0C) + 32 • Celsius (centigrade) T(0C ) = [T(0F) –32]/1.8 • Kelvin (Celsius + 273)

  6. Examples • Zero degrees Celsius is what Kelvin? Answer: 273o • What is the boiling point of water in degrees Kelvin? Answer: 373o • 200 degrees Celsius is what in Kelvin? Answer: 473o

  7. Absolute Zero • 0 degrees Kelvin = -273 Celsius • Lowest possible temperature • Molecular motion ceases Courtesy Michigan State University

  8. Kinetic Theory of Heat • All matter is made of tiny atoms and molecules, constantly in motion • Faster is hotter solid gas

  9. Temperature and Kinetic Energy • In ideal gas temperature is proportional to average kinetic energy per molecule • Closely related in liquids and gases

  10. Temperature does not depend on the amount

  11. Heat does depend on the amount There is twice as much kinetic energy of moving molecules in two liters of water as in one liter. Analogy: Heat is like the total height of students in this room, temperature is like their average height.

  12. Which has more heat? • A swimming pool full of ice water? • A cup full of boiling water? Answer: the swimming pool, because it has so much more water.

  13. Heat – Energy Transferred • Definition: Energy that transfers because of temperature difference • Heat flows governed by average molecular kinetic energy difference • Heat always flows from high energy to low • Cold- Absence of Heat * does not exist!

  14. Thermal Equilibrium • Objects at same temperature are at thermal equilibrium – no heat flows.

  15. Measuring Heat • One calorie is the amount of heat needed to raise the temperature of one gram of water by one degree Celsius. • Kilocalorie raises the temperature of one kg of water by 10 C (also called Calorie or food calorie) • One calorie = 4.186 joules • One kilocalorie = 4186 joules • One kilocalorie= 1 Food Calorie

  16. Density of Water • Density is mass per unit volume • D = M/V • One gram per cubic centimeter • One kilogram per liter • One thousand kg per cubic meter

  17. Specific Heat Capacity • Different materials change their temperature by different amounts when they absorb the same amount of heat. • Some have more ways of storing energy than others • Water has very high specific heat (capacity) • Metals have much less

  18. Q = mcDT • Q = mcDT expresses how heat absorption works. C is specific heat • Question: A certain rock has a specific heat of 0.25 (water is 1.0) How much heat will be required to heat 5.0 kg rock from 20 to 800C? • Q = 5.0kg x 1000g/kg x 0.25 Cal/g 0C x 60 0C • Q = 75,000 C = 7.5 x 104 Calories

  19. Calories • 1 calorie raises the temperature of 1 gram of water by 1 Celsius degree. • 1 kilocalorie (kcal or Calorie) raises temperature of 1 kg of water by 1 degree Celsius • 1 British Thermal Unit (BTU) raises temperature of 1 pound of water by one degree Fahrenheit

  20. Application • If 10 calories of heat go into a gram of water, how much will the temperature increase? Answer : 10 degrees C • How much heat is needed to raise the temperature of 10 grams of water by one degree C? Answer: 10 calories • How much heat is needed to raise the temperature of 10 grams of water by 10 degrees C? • Answer: 100 calories

  21. Specific Heat • How much does temperature rise when heat is put into something? • It depends on the material as well as the mass and the quantity of heat: Q = m c Dt c is specific heat in calories/g oC • Water has the highest specific heat of any common material, 1 cal/g oC • Metals generally have low specific heats, which makes them easy to cool or heat.

  22. Temperature Change • Exothermic Reactions, are where heat is released, this will cause the temperature of the surrounding to increase. (think a roaring fire to warm a cold room) ∆T = Tf-Ti • Endothermic Reactions, are where heat is absorbed, this will cause the temperature of the surroundings to decrease. (think ice melting in warm drink) ∆T = Ti-Tf

  23. Example • How much heat is required to raise the temperature of 1000g water from room temperature (20oC) to boiling (100oC)? • Q = m c Dt = 1000g x 1 cal/g oC x 80 oC = 80,000 calories (or 80 kilocalories) Fact: It would take about a tenth as much heat to raise the temperature of an equal amount of iron this much

  24. Other Examples • If 2000 J of energy are added to .25 kg of water at an initial temperature of 35⁰C what will the final temperature be? • If 5000 cal are added to an unknown substance and 200 g of it change temperature by 34⁰C what is the specific heat of that substance?

  25. Mixtures • 100 g of water at 50oC is added to 100g water at 70oC. What will be the final temperature? • You guessed it: 60oC • Mix a liter of 20oC water with two liters of 30oC water. What is the final temperature of the system?

  26. Calorimeter Energy device used to eliminate the loss of energy to the surroundings. Thermos and styrofoam cup work on same principles.

  27. High Specific Heat of Water • Makes it a good coolant (water also has high conductivity although this is not the same) • Large bodies of water such as oceans moderate climate • Gives coastal communities relatively mild summers and winters • Another peculiar fact about water. It’s highest density (and smallest volume) is at 4oC. • Water at bottom of frozen lake is always 4oC

  28. Change of Phase

  29. Heat of Fusion • Energy needed to melt 1 gram of a solid into its liquid form. • Hf ice= 80 cal/g • Hf= Q/m • Q= Hf x m • m= Q/Hf • This is endothermic, absorbed into the system, the same is true of the releasing, just called Heat of Crystallization.

  30. Heat of Vaporization • Energy needed to vaporize 1 g of some liquid into a gas. • Hv water= 540 cal/g • Hv= Q/m • Q= Hv x m • m = Q/Hv • This is endothermic as well, again being absorbed, the same is true of releasing just called Heat of Condensation.

  31. Phase Change Example • How many calories are required to raise the temperature of 85 grams of – 20⁰C ice to 150⁰C? (Think about all the different steps you must take to solve this problem) ∆T of ice Phase change (melt the ice) ∆T of water Phase change (boil the water) ∆T of steam

  32. Heat of Fusion Lab (Prove Hf= 80 cal/g) • Cook 125 ml of water to exactly 50⁰C. Record Ti. • Pour exactly 100 ml of water into sty. cup. Record Vi. (DO NOT DESTROY STY. CUPS) • Add ice chips 2-3 at a time, carefully stirring with thermometer, until you reach 0⁰C. Record Tf. • Carefully remove any ice chips. Record Vf. • Remember Dwater 1g= 1 ml • Calculate Q (released by water)= m x ∆T x c • Calculate the mass of melted ice= Vf-Vi • Calculate the Hf of ice = Q/m • Answer this question: Why did we start at 50⁰C? Hint: Think about heat exchange and room temperature.

  33. Thermal Expansion • Most materials expand when heated • Only exception is water between 00C and 400C • Expansion joints in bridges, cracks in sidewalks allow for expansion

  34. Bimetallic Strip • How your thermostat works

  35. Don’t Let Your Car’s Engine Overheat • Aluminum expands more than iron • Pistons made of aluminum • Cylinder made of iron

  36. Mechanical Equivalent of Heat • Discovered by James Joule • Falling weight makes paddle turn • 4.186 x 103 J = 1 kcal • Interpretation: HEAT IS ENERGY TRANSFER Courtesy W. Bauer http://lecture.lite.msu.edu/~mmp/kap11/cd295.htm

  37. Joule’s Apparatus Link to Joule’s original article

  38. Example • When digested a slice of bread yields 100 kcal. How high a hill would a 60 kg student need to climb to “work off” this slice of bread? 100 kcal x 4.186 x 103 J/kcal = 4.2 x 105 J W = mgh h = W/mg = 4.2 x 105 / (60 kg)(9.80 m/s2) = 714m = 7.1 x 102 m If the body is only 20 percent efficient in transforming the bread, how high need they climb?

  39. Bullet in Block • When a 10 g bullet traveling 500 m/s is stopped inside a 1kg wood block nearly all its KE is transformed to heat. How many kcal are released? KE = ½ mv2 = 0.5 x 0.010 kg x (500)2 = 1250 J 1250 J x 1 kcal/4186 J = 0.30 kcal

  40. Thermodynamics • Study of heat and its transformation into mechanical energy • Based on conservation of energy • Explains how engines like car motors work

  41. First Law of Thermodynamics • Generally, when you add heat to a system it changes into an equal amount of some other form of energy • Heat added = increase in internal energy + external work done by the system

  42. Work Done On and By • Compressing a gas by pushing down on a piston = work done on • A gas expands by pushing a piston up = work done by

  43. Questions • 20 J of heat is added to a system that does no work. What is the change in internal energy? Answer +20 J • 20 J of heat is added to a system that does 10 J of work. What is the change in internal energy? Answer +10 J

  44. 20 J of heat is added to a system that does 30 J of work. What is the change of internal energy? Answer -10 J • 20 J of heat is added to a system that has 10 J of work done on it. What is the change of internal energy? Answer +30 J

  45. Bicycle Pump • What do you think happens when you operate the pump. Where does the work you do go? • It goes to heat, some through friction, some to adiabatic compression of the air inside the pump • What does “adiabatic” mean? • Answer: No heat enters or leaves Q=0

  46. Adiabatic Processes • Compression or expansion of a gas so that no heat enters or leaves • Example: gas in cylinder of car or diesel engine • Why adiabatic? Because it happens too fast for much heat to enter or leave. • In adiabatic compression, temperature rises. • In diesel engine, enough to ignite gas without spark plug A process can also be adiabatic if it happens inside a well insulated conatiner.

  47. Courtesy “How Stuff Works”

  48. Courtesy Shell Canada

  49. Adiabatic Expansion • Produces cooling • Example: blow on your hand first with wide open mouth, then with puckered lips • How do you explain the results?

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