Thermodynamics

1. Thermodynamics Chapter 17 in Prentice Hall Text and Ch. 13.4 from Previous Unit

2. Thermodynamics • The study of energy • Changes that occur during chemical reactions and changes in states • For example: • Physical State change • How it is transformed from one form to another • How it is used to get things done

3. I. ENERGY • The ability to do work or to produce heat. • The ability to move or change something. • There are two types of energy: • Potential Energy • Kinetic Energy

4. Phase Diagram • A to B • Solid • KE • B to C • Phase Δ • Solid & Liquid • PE • C to D • Liquid • KE • D to E • Phase Δ • Solid & Liquid • PE • E to critical point • Gas • KE

5. A. Potential Energy • Potential energy: Stored energy that’s waiting for its chance to get moving. (i.e. objects that are waiting to fall off of a shelf, energy stored in chemical bonds, etc). • Chemical potential energy: • The energy that’s stored in chemical bonds. • Phase Changes

6. B. Kinetic Energy • Kinetic energy: The energy something has when it moves. (i.e. moving objects, moving particles, vibrating molecules, etc.) • Temperature is a measure of the particles in an object. • The kinetic molecular theory, which says that the amount of energy is proportional to the temperature (in K). • The more the particles in an object move around, the higher the temperature.

8. 1. Kinetic Molecular Theory of Heat • Atoms held by electromagnetic forces • Atoms vibrate • Thermal energy is a measure of this motion. • More motion  more heat • Add heat  increase motion and thermal energy. • Think atoms: When heat or pressure is applied to a gas what happens to the molecules • More energy -> more collisions and more volume

10. 2. Temperature Definitions • A measure of average particle kinetic energy • A measure of thermal energy. • Not dependent on mass. • Example Definitions: • The degree of hotness or coldness of a body or environment. • A measure of the warmth or coldness of an object or substance with reference to some standard value. • A measure of the average kinetic energy of the particles in a sample of matter, expressed in terms of units or degrees designated on a standard scale. • A measure of the ability of a substance, or more generally of any physical system, to transfer heat energy to another physical system. • Any of various standardized numerical measures of this ability, such as the Kelvin, Fahrenheit, and Celsius scale.

11. 3. A look at thermometers and temperature scales

12. i. How does a thermometer work? • Liquid (usually an alcohol) encased in narrow glass tube • Liquid thermometers are based on the principal of thermal expansion • When a substance gets hotter, it expands to a greater volume. • As the temperature of the liquid in a thermometer increases, its volume increases.

13. ii. Temperature Scales • Celsius: freezing point = 0 boiling point = 100 • Kelvin: absolute zero = 0 (atoms are motionless) • Kelvin = SI unit of temperature.

14. Compare Scales K = ºC + 273 °C = (°F - 32°)/1.8

15. Which is the smaller temperature increment - a degree Celsius or a degree Fahrenheit? Explain. Answer: degree Fahrenheit The degree Fahrenheit is the smaller increment. After all, there are 180 of these Fahrenheit divisions between the normal freezing point and the normal boiling point of water; there are only 100 of the Celsius divisions between these two temperatures. If more Fahrenheit divisions can be fit between these two divisions than Celsius divisions, then the Fahrenheit divisions must be smaller.

16. Perform the appropriate temperature conversions in order to fill in the blanks in the table below.

17. Answers:

18. iii. Why important? • Only a few countries use Fahrenheit • Planning a trip to ski in Canada or Swiss Alps? • To understand physics concepts and to communicate scientifically

19. 4. Quantifying Energy • Measuring energy: • The traditional unit of energy is the calorie (cal), which is the amount of energy you need to add to 1 gram of water to heat it by 10 C. • Food is measured in units of 1000 calories called kilocalories (kcal), which is more commonly known as the Calorie (Cal). • The metric unit (SI unit) of energy is the joule (J). There are 4.184 J/cal or 1 calorie = 4.184 J. • Because a joule isn’t very much energy, we usually measure energy in units of 1000 joules called kilojoules (kJ). • Fun Fact: John Dalton (credited for the atomic theory) tutored James Prescott Joule.

21. Temperature vs Heat • Temperature is a measure of the average kinetic energy and does not depend upon the amount of matter in the sample. • Heat is the total kinetic energy that flows because of a difference in temperature and does depend on the amount of matter. • Net Worth is to Temperature as Heat is to Currency • Think if you give Beyoncé $5 vs Mrs. Heaton $5 • Consider that one drop of heat (food coloring) has been added to each. Which would be the hottest?

22. Thermometers are Speedometers • Temperature depends on amount of matter

23. Consider two samples of different gases. One sample consists of helium atoms and the other sample consists of diatomic oxygen molecules. If the samples are at the same temperature, will the particles within the sample have the same average speed? • Answer: No • Temperature is a measure of the average kinetic energy of the samples. Translational kinetic energy depends upon both the mass of the particles and the average speed at which the particles move. In comparing two samples of different gases at the same temperature, the gas with the more massive particles has the slowest particle speeds. So in comparing the speeds of helium atoms and diatomic oxygen molecules, one must be conscious of the relative masses of the two particles. Helium particles, being roughly one-eighth the mass of diatomic oxygen molecules, will move with a considerably faster speed.

24. The particles in a sample of table salt (sodium chloride) are not free to move about. They are locked in place in a structure known as a crystal lattice. Can the particles of sodium chloride possess kinetic energy? • Answer: Yes • Even though they do no possess any translational kinetic energy, they still possess some vibrational kinetic energy. The sodium and chloride ions can wiggle about their fixed lattice positions. The back and forth vibrational motion of the particles is what gives them vibrational kinetic energy. This explains why a thermometer will register a temperature when placed in the sample of matter.

25. C. Heat Transfer

26. Heat and Heat Transfer • Heat: energy flows between objects due to temperature difference • Flows always from hot to cold • Heat Transfer: • Conduction • Convection • Radiation

27. To describe heat transfer we need to define what the heat is being transferred between • System: Whatever is being studied or observed. • Surroundings: Everything outside the system. • Universe: The system + the surroundings.

28. Heat Transfer and Temperature Heat lost by one must equal heat gained by other Both end up at same final temperature • ENDOTHERMIC: System to Surroundings [bonds breaking] • Feels cold and WORK is done BY the SYSTEM • EXOTHERMIC: Surroundings to System [bonds forming] • Feels hot and WORK is done ON the SYSTEM hot object cold object Heat insulation

29. 1. Thermal equilibrium • The energy transfers back and forth between two objects are equal (two sides) • Biology: Concentration inside/outside cell • Chemistry: reactants to products • Physics: systems vs surroundings • OR net zero energy transfer • Scenarios • Ice on a table • Where is the heat? • Air • What is making the ice melt? • A warm can in a refrigerator • Can loosing heat, warming fridge

31. For each of the following designations of a system and a surroundings, identify the direction of heat flow as being from the system to the surroundings or from the surroundings to the system.

32. For each of the following designations of a system and a surroundings, identify the direction of heat flow as being from the system to the surroundings or from the surroundings to the system.

33. 2. ComparingSpecific Heat vs Heat Capacity

34. a. Specific Heat • Specific Heat (C): • the amount of energy required to raise 1 kg by 1 K or 1 g by 1 ◦C • Depends on amount of heat, change in temperature and material • Unit = J/kg∙K In chemistry we usually see J/g∙K • Examples: Water C = 4180 J/kg∙K Iron C = 450 J/kg∙K Copper C = 385 J/kg∙K • Each phase has a specific heat: ex) Specific Heat of Water is greater than that of steam

35. Check Points • 1) What does a thermometer measure? Average kinetic energy • 2) Does a thermometer measure heat? No • 3) What work and energy is done by your car as you drive to school? The car moves and carries you to school (work) and the energy change occurs as fuel is burned (heat released). • 4) Why do farmers or gardeners spray plants with water or flood farms with an inch of water before a winter freeze?

36. b. Heat Capacity • Amount of heat (q) needed to increase exactly 1◦C of an object • Depends on: • Mass • Chemical Composition • Increase mass, increase heat capacity

37. Check Point • 5) Consider and explain Mark Twain's quote in regards to heat capacity and specific heat: • “The coldest winter I ever spent was the summer I spent in San Francisco.”

38. II. Equations

39. A. Specific Heat:Heat Transferred • Amount of heat gained or lost by an object Q = mC∆T • Q = heat transferred (unit = joules) • m = mass • C = specific heat • ∆T = change in temperature (Tfinal– Tintial) • Next slide expands on these variables and in Part III. Enthalpy (H) is expanded upon.

40. The enthalpy change that accompanies the heating/cooling of a pure substance is determined by the equation: q = ∆H = mCp∆T • ∆H = the change in enthalpy (positive for heating, negative for cooling) • m = the mass of the thing being heated (in grams) • Cp = the specific heat / heat capacity – the amount of energy needed to heat the thing by 10 C. • ∆T = the change in temperature (in degrees Celsius). • q = heat [NOTE: If there is no non-expansion work on the system and the pressure is still constant, then the change in enthalpy will equal the heat consumed or released by the system (q)- this just means When PRESSURE is constant (as it is in solids and liquids) q = ∆H]

41. Example 1 kg block of iron is heated from 37 ºC to 57 ºC. The specific heat of iron is 450 J/kg∙K. What was the heat transferred to the iron? Q = mC∆T = (1 kg)(450 J/kgC)(57 ºC – 37 ºC) = 1 x 450 x 20 = 9,000 J

42. B. Calorimetry: measuring heat of a process • -Qlost = Qgained • -mC∆Tlost = mC∆Tgained • -mC(Tf-Ti)hot object = mC(Tf-Ti)cold material • Watch your +/- signs

43. Example You dropped a 0.020 kg block of iron at 180 ºC into 0.40 kg tank of water at 10 ºC. The Ciron is 450 J/kg∙K and the Cwater is 4180 J/kg∙K. What’s the final tempurature? mC(Tf-Ti)iron + mC(Tf-Ti)water = 0 (0.02)(450)(Tf – 180) + (0.4)(4180)(Tf – 10) = 0 9(Tf -180) + 1672(Tf -10) = 0 9Tf – 1620 + 1672Tf – 16720 = 0 1681Tf – 18340 = 0 1681Tf = 18340 Tf = 18340/1681 Tf = 10.9 ºC

44. C. Phase Change Energy • When a substance is heated or cooled, you must take into account the energy at the phase changes. • To do this we use the following equations • Where Enthalpy (H) is discussed later on in the slides

45. Example • How much energy is needed to heat 55 grams of water ice from a temperature of -150 C to steam at a temperature of 1500 C? • This problem requires several steps: • Step 1: Heat the ice from -150 C to 00 C • ∆H = mCp∆T = (55 grams)(2.03 J/g0C)(150) = 1,675 J • Step 2: Undergo the phase change from a solid to liquid. • ∆H = n∆Hfus = (3.06 mol)(6.01 kJ/mol) = 18.39 kJ • Step 3: Heat the liquid from 00 C to 1000 C • ∆H = mCp∆T = (55 grams)(4.184 J/g0C)(1000) = 23,012 J • Step 4: Undergo the phase change from a liquid to gas. • ∆H = n∆Hvap = (3.06 mol)(40.7 kJ/mol) = 124.54 kJ • Step 5: Heat the steam from 1000 C to 1500 C • ∆H = mCp∆T = (55 grams)(2.01 J/g0C)(500) = 5,528 J • And when you’re done with all of this, just add up the different values you found above: ∆Htotal = 1.68 kJ + 18.39 kJ + 23.01 kJ + 124.54 kJ + 5.53 kJ ∆Htotal = 173.15 kJ

46. Example #9 from HW

47. III. Laws of Thermodynamics

48. The First Law of Thermodynamics: • Energy can’t be created or destroyed, but can change form. • Amount of energy in a system remains constant

49. A. Enthalpy (H) • Heat content of a system at constant pressure • The amount of heat that a system can potentially give to other systems (heat content of a system).

50. 1. Why do we use ΔH vs H? • This internal energy (the sum of all possible forms of energy in the system) cannot be measured. • THEREFORE, instead of talking about how much enthalpy something has, we instead talk about how much the enthalpy of a system changes (hence the triangle) when heat is taken away from it or added to it. • This term is given the symbol ∆H, where ∆ represents the change in enthalpy that occurs during a process.