Chapter 12 Thermal Energy

1. Chapter 12 Thermal Energy Quiz 12

2. Chapter 12 Objectives • Describe the nature of thermal energy • Define temperature and distinguish it from thermal energy • Use the Celsius and Kelvin temperature scales and convert on to the other

3. Chapter 12 Objectives • Define specific heat and calculate heat transfer • Define heats of fusion and vaporization • State the first and second laws of thermodynamics

4. Chapter 12 Objectives • Define heat engine, refrigerator, and heat pump • Define entropy

5. Temperature • Simply stated, it can be thought of as average kinetic energy of the particles • Why some water evaporates before 100 C is reached

6. Temperature • To be more precise, we can talk about Thermal Equilibrium. • Energy always goes from hot to cold. Energy is always transferred from hot to cold. • Heat is the transfer of thermal energy • Particle Collisions: More energetic particle gives energy to the less energetic particle

7. Temperature • If two objects are at different temperatures, heat which is the transfer energy, will flow until this equilibrium is reached. • Temperature is a quantity that determines when objects are in thermal equilibrium. • When Temperature stops changing, equilibrium has been reached • This is why absolute zero can never be touched, equilibrium is not reached.

8. Temperature Scales • Three Scales: Kelvins, Celsius, Fahrenheit • Kelvins Absolute Temperature Scale, it has a value of 0 at absolute 0 • Important to use an absolute temperature scale in many calculations otherwise you would find negative volumes of air as well as other unfortunate mistypes in the calculator

9. Temperature Scales • Celsius  Has same slope as Kelvins, just starts below it 273 degrees. • Water freezes at 0 C and boils at 100 C, convenient • Fahrenheit  The American scale • Below equation is green

10. Quick Questions • If the temperature of the room is at equilibrium, which state of matter in this room has the most kinetic average per particle? • Can an individual molecule touch absolute zero? Defend your rationale.

11. Energy Transfer • Conduction: Direct contact • Molecules come in direct contact, the fast moving particle makes the slow moving particle speed up • You touch a hot stove • Convection: Mass movement • Weather front, currents in rivers and oceans • Radiation: Waves • Light, in all forms (radio  Gamma)

12. Factors which influence conduction (rate of heat transfer) are • Temperature difference • Length of material • Cross sectional area of material • Nature of the material • Metals are good, nonmetallic bad (electrons) • Type of matter • Solid > Liquid > Gas (Due to distance between)

13. Don’t write this slide down • This is all summarized by Fourier’s Law of heat conduction • P= • Where k is constant dependent on material, A is cross sectional area, delta T is temperature difference, and d is thickness or length of material.

14. Convection • Fluid currents carrying heat from one place to another is convection • Convection: The material moves from one place to another • Conduction: The energy flows, but the material does not. • Convection only occurs in fluids.

16. Double Paned Windows • Is it possible to have too large of a gap between the two windows?

17. Convection Factors • Rate of convection of heat flow is much more difficult, but factors that influence it are the convection coefficient, surface area exposed, and temperature difference. • Wind is also a factor when dealing with convective heat loss off a person

18. Heat Capacity (Big C) • Capacity, or how much something can hold. • In the thermodynamics world, this is referring to how much energy an object can hold before increasing in temperature. • Heat Capacity can be defined as • C = Q / T (C being heat capacity, Q being Joules, T being change in T)

19. Positive or Negative Energy • Remember positive and negative work? You can have positive and negative energy changes as well. Think of objects as having a bank account. • If the object gains energy, it is a deposit and the balance goes up. • If the object loses energy, it is a withdrawal and the energy lost is negative.

20. Question • An object absorbs 1,500 J of energy from the surroundings and increases in temperature by 50 K. What is the heat capacity of the object?

21. Specific Heat Capacity (little c) • Heat capacity becomes more ‘specific’ when you include the mass of the object. The specific heat capacity of something is • Heat Capacity / mass  C / m Q / mT • Specific heat is also written as Cp

22. Questions • Gold has a specific heat of 0.128 kJ/kg*k. How much energy will a bar of gold with a mass of 248g need to absorb to increase its temperature by 25 K? • Rank the following objects according to temperature change if each had the same mass and were given the same amount of energy. • Lead = 0.13 J/g*K Iron = 0.44 J/g*K • Silver = 0.235 J/g*K Copper = 0.385 J/g*K

23. Calorimetry • The study of how much energy is in something. • Often associated with combustion reactions, most relevant to you as how much energy is in the food you eat. It is also useful for identifying unknown materials.

24. Question • While looking through a menu at a fast food restaurant, you notice one of the burger has 500 Calories in it. If you were to burn the burger in a calorimeter (which has a Heat capacity of 742 J/K), by how many degrees will the calorimeter’s temperature rise?

25. Question • Water has a specific heat capacity of 4.186 J/g. A beaker holds 400 mL of water at a temperature of 80.0 C. If a block of marble (0.86 J/g*K) with a mass of 60 g with an initial temperature of 20.0 C is placed into the water, what will be the final temperature.

26. Reminder • Heat Capacity is different then specific heat capacity by the addition of mass. • The human body has a certain heat capacity. • Larger bodies have larger heat capacities • The human body also has a rough specific heat capacity • Doesn’t matter what the size is, it is per gram of mass

27. Phase Transition • A phase transition is when a substance changes phases (phases being Solid, Liquid, and Gas). • To fully understand a phase transition, we need to understand what is happening at a molecular level and the bonds that hold atoms together. • Solid: Atoms are in fixed position • Liquid: Atoms are mobile but held together • Gas: Atoms are mobile and independent of one another

28. Enthalpy of Fusion • In order for a solid to become a liquid, it takes a certain amount of energy for the atoms, that were in a fixed position, to become ‘free’ of the bond they had. • The strength of the fixed bond compared to the “mobile” bond is the Enthalpy of Fusion (Solid to Liquid). Enthalpy of Fusion is the same number for freezing and melting, just reversed signs (+ or -)

29. Enthalpy of Vaporization • In order for a molecule or atom to completely break the bond they had with the nearby molecules is the Enthalpy of Vaporization (Gas to Liquid). • Again, this is the same number but opposite sign for boiling and condensing. • Mobile Bond  No Bond

30. Reminder • Take note, the temperature of the molecules don’t change during a phase transition. • The Kinetic energy of the molecules is not changing, just the bonds that hold them in place.

31. How much energy for the phase change? • The energy for the phase change is called Latent Heat • Energy required = Specific Heat (for temperature change) + Latent Heat (for phase change)

33. Questions • What values are larger, enthalpies of fusion or enthalpies of vaporization? Why? • On a molecular basis, how does the strength of IMF’s (intermolecular forces) affect the value of enthalpy of vaporization?

34. Questions • An athlete playing __________ is working hard and works up a sweat. How much energy does the body lose when 10g of water evaporates off of his skin? • The enthalpy of vaporization of water is 2,256 J/g

35. Questions • Oranges are kept safe by spraying a mist of water on them when it is cold out. How much energy does an orange absorb when 7.5g of water at a temperature of 16.5 C is cooled down to 0.0 C ice??

36. Questions • A hot cup of cocoa is way too warm to drink. The temperature of the cocoa is 85.0 C and has a mass of 320g. What will be the final temperature of the drink if 6.2g of 0.0 C ice cubes are added to the drink? The enthalpy of fusion of water is 333.7 J/g.

37. Question • A hot cup of soup is too hot to eat. The temperature of the soup is 90.0 C and has a mass of 400g. What will be the final temperature of the drink if 6.4g of 0.0 C ice cubes are added to the soup? The enthalpy of fusion of water is 333.7 J/g. Assume the soup to be water.

38. First Law of Thermodynamics • Energy is neither created or destroyed • Total increase in thermal energy of a system is equal to the work done on it and the heat added to it

39. Heat Engine • Turns heat energy into mechanical energy

40. Reverse • Fridge and air-conditioner

41. 1st Law of Thermodynamics • 1st Law: The change in internal energy of a system is equal to the heat flow into the system plus the work done on the system • U = Q + W • U = Change in Internal Energy • Q = Joules of Heat • W = Work

42. 2nd Law of Thermodynamics • Processes tend to increase the entropy (the disorder) of a system • Heat flows increasing the disorder the heated object

43. 1st Law of Thermodynamics • The system can be whatever you want to define it to be (at your discretion). The energy of the system is not only dependent upon heat flow, but on work as well. • Keeping the signs straight is important for solving the change in internal energy.

44. Meaning of + sign Q: Heat flows into the system W: Work done on the system U: Internal energy increases Meaning of – sign Q: Heat flows out of the system W: Work done by the system U: Internal energy decreases Meaning of Signs (Know for quiz)

45. Question • A Gas expands and cools down. • What is the sign of work for the gas as it expands? • What is the sign of Q for the gas? • Is it possible to tell whether the value of U is positive or negative?

46. 2nd Law of Thermodynamics • 2nd Law: The entropy of the universe never decreases • The entropy (Unit S and units of J/K) of a system is a measure of the disorder of the system. • Heat flows spontaneously from hot to cold objects because of the increase in entropy. • The entropy increase of the cold material > the entropy decrease of the hotter material and entropy of the universe increases.

47. 2nd Law of Thermodynamics • It is impossible for any process to decrease the total entropy of the universe. • It is possible to decrease the entropy of one system, but it is at the greater expense of another system.

48. 2nd Law of Thermodynamics • A generalization of the change in entropy is Q/T, where energy flowing into the system divided by the change in temperature of the metal is the change in entropy. • Energy into system = Increased Entropy • A positive entropy means increase in disorder, a negative s means a decrease in disorder.

49. 2nd Law of Thermodynamics • Entropy is also described statistically by probability of the molecules being at certain energies and locations. • How does this explain why it takes no energy for gas to expand to fill its container?

50. Questions • Heat spontaneously is flowing out of a system to its surroundings. The system has no change in internal energy though. What are the signs of W and S of the system? • System A and System B are in contact with one another. System A is warmer than System B. What are the signs of W and S for both systems? Which system has a larger change in S (absolute value)?