2. Check scantrons Final Exam Grades: % midterm + (2 X %final) 3 MIDTERM Pd 1 Day 3 2-25 Final Exam Grades: (% you want X 3) - % midterm 2 = % you need on the final

3. Question of the Day 1. A 1.0 mole sample of ethanol (C 2 H 5 OH) has a heat capacity of 110.4 J/°C. Calculate the specific heat of ethanol (hint: What is the relationship between heat capacity and specific heat?) Day 3 2-25= 2.4 J/g°C

4. Question of the Day 2. The specific heat of ice is 2.1 J/g°C. What is the heat capacity for a 2.0 mole sample? Day 3 2-25 Period 3 only = 75.6 J/°C

5. 4 Thermodynamics Energy Enthalpy Entropy Spontaneity

6. 5 State Functions Property determined under specified conditions of T, P, location, etc Independent of HOW conditions are reached

7. 6 State Functions Examples _____________________________________ Energy d A B B A  E = difference between A & B

8. 7 State Functions: Hess’s Law For a reaction that occurs in several steps, the sum of the  ’s for each step is the  for the net reaction. Thus, it is possible to calculate  for a complex reaction by adding the  ’s of each simple reaction that leads to the complete reaction.

9. 8 State Functions: Hess’s Law The Equation  TOTAL =  step 1 +  step 2 +  step 3 + … +  step n

10. 9 Possibilities: From system to surroundings From surroundings to system How heat transfer is made: Heat = q Positive change = heat added to system Negative change = heat released from system Work = w Positive change = work done ON system Negative change = work done BY system Energy Transfers:

11. 10 Energy Transfers First Law of Thermodynamics Energy lost by the system EQUALS energy gained by the surroundings Energy gained by the system EQUALS energy lost by the surroundings

12. 11 Energy Transfers Exothermic Change Energy moves from ______________ to ________________ _________ of q Temperature of surroundings ________________ system surroundings Loss increases

13. 12 Energy Transfers Endothermic Change Energy moves from ________________ to _____________ _________ of q Temperature of surroundings ________________ surroundings system Gain decreases

14. 13 Energy Transfer Equations Internal Energy:  E = E f – E i E f = final energy E i = initial/starting energy Energy Change due to Heat & Work:  E = q + w Use + q if endo Use – q if exo Use + w if work is done ON system Use – w if work is done BY system

15. Assignment Review section 17.1 and complete #s 5-11 on page 561. Due Monday 2-25 for period 3 Due Tuesday 2-26 for period 1

16. Question of the Day 1. What factors determine the heat capacity of an object? mass and chem. Comp. Day 4 2-26

17. Question of the Day 2. The specific heat of ice is 2.1 J/g°C. What is the heat capacity for a 2.0 mole sample? Day 4 2-26 Period 1 only = 75.6 J/°C

18. 17 Enthalpy Thermodynamic property that measures/describes the changes in heat content of a system under constant pressure. Symbol: Units of measure: Equations and usage H,  H cal, kcal, J, kJ State function  H = H f - H i +  H H f > H i Heat gained & endo -  HH f < H i Heat lost & exo

19. Constant-Pressure Calorimeters The enthalpy (H) of a system accounts for the heat flow of the system at constant pressure. The heat absorbed or released by a reaction at constant pressure is the same as the change in enthalpy, symbolized as ΔH. BOOK - Calorimetry

20. Constant-Pressure Calorimeters The value of ΔH of a reaction can be determined by measuring the heat flow of the reaction at constant pressure. In this textbook, the terms heat and enthalpy change are used interchangeably. In other words, q = ΔH. BOOK - Calorimetry

21. 20 Calculating Enthalpy Change in Reactions Formula:  H RXN =  H PRODUCTS -  H REACTANTS 2 H 2(g) + O 2(g) __________________________ __________________________ 2 H 2 O (g) Higher E means Less stability Lower E means more stability + – 483.6

22. 21 Calculating Enthalpy Change in Reactions CH 4(g) + 2O 2(g) CO 2(g) + 2H 2 O (g) CO 2(g) + 2H 2 O (l) H + + – – 802 890

23. 22 Calculating Enthalpy Change in Reactions Phase matters!!!!!!! Heat In = Heat Out

24. Thermochemical Equations How can you express the enthalpy change for a reaction in a chemical equation? In a chemical equation, the enthalpy change for the reaction can be written as either a reactant or a product.

25. Thermochemical Equations In the equation describing the exo reaction of calcium oxide and water, the enthalpy change can be considered a product. CaO(s) + H 2 O(l) → Ca(OH) 2 (s) + 65.2 kJ Calcium oxide is one of the components of cement.

26. Thermochemical Equations CaO(s) + H 2 O(l) → Ca(OH) 2 (s) + 65.2 kJ A chemical equation that includes the enthalpy change is called a thermochemical equation.

27. Thermochemical Equations The heat of rxn is the ΔH for the chemical equation exactly as it is written. Heats of reaction are reported as ΔH. The physical state of the reactants and products must also be given. The standard conditions are that the reaction is carried out at 101.3 kPa (1 atm) and 25°C. Heats of Reaction

28. Thermochemical Equations CaO(s) + H 2 O(l) → Ca(OH) 2 (s) Each mole of calcium oxide and water that reacts to form calcium hydroxide produces 65.2 kJ of heat. ΔH = –65.2 kJ In exothermic processes, the chemical potential energy of the reactants is higher than the chemical potential energy of the products. Heats of Reaction

29. Thermochemical Equations 2NaHCO 3 (s) + 85 kJ → Na 2 CO 3 (s) + H 2 O(l) + CO 2 (g) Baking soda decomposes when it is heated. This process is endothermic. Heats of Reaction The carbon dioxide released in the reaction causes muffins to rise while baking.

30. Thermochemical Equations 2NaHCO 3 (s) + 85 kJ → Na 2 CO 3 (s) + H 2 O(l) + CO 2 (g) Heats of Reaction Remember that ΔH is positive for endothermic reactions. Therefore, you can write the reaction as follows: 2NaHCO 3 (s) → Na 2 CO 3 (s) + H 2 O(l) + CO 2 (g)ΔH = 85 kJ

31. Thermochemical Equations The amount of heat released or absorbed during a rxn depends on the number of moles of the reactant involved. The decomposition of 2 mol of sodium bicarbonate requires 85 kJ of heat. Therefore, the decomposition of 4 mol of the same substance would require twice as much heat, or 170 kJ. Heats of Reaction

32. Thermochemical Equations Heats of Reaction To see why the physical state of the reactants and products must be stated, compare the following two equations. difference = 44.0 kJ H 2 O(l) → H 2 (g) + O 2 (g)ΔH = 285.8 kJ 1 2 H 2 O(g) → H 2 (g) + O 2 (g)ΔH = 241.8 kJ 1 2

33. Calculate the amount of heat (in kJ) required to decompose 2.24 mol NaHCO 3 (s). Using the Heat of Reaction to Calculate Enthalpy Change 2NaHCO 3 (s) + 85 kJ → Na 2 CO 3 (s) + H 2 O(l) + CO 2 (g) Sample Problem 17.4

34. Assignment: #s 14 and 15 on page 567 (show work), #s 16, 17, 20, and 21 on page 568 (show work) Day 4 2-26

35. Question of the Day: 3C 2 H 2(g)  C 6 H 6(l) + 630 kJ Is this reaction endo OR exo? How much heat is involved if you start with 2 grams of acetylene (acetylene = C 2 H 2 )? EXO 16.2 kJ

36. Thermochemical Equations Heats of Combustion The heat of combustion is the heat of reaction for the complete burning of one mole of a substance.

37. Thermochemical Equations Heats of Combustion Small amounts of natural gas within crude oil are burned off at oil refineries. This is an exothermic reaction. Burning 1 mol of methane releases 890 kJ of heat. The heat of combustion (ΔH) for this reaction is –890 kJ per mole of methane burned. CH 4 (g) + 2O 2 (g) → CO 2 (g) + 2H 2 O(l) + 890 kJ

38. Like other heats of reaction, heats of combustion are reported as the enthalpy changes when the reactions are carried out at 101.3 kPa and 25°C. Interpret Data Heats of Combustion at 25°C SubstanceFormulaΔH (kJ/mol) HydrogenH2(g)H2(g) –286 Carbon C(s, graphite) –394 MethaneCH 4 (g) –890 AcetyleneC2H2(g)C2H2(g)–1300 EthanolC 2 H 6 O(l)–1368 PropaneC3H8(g)C3H8(g)–2220 GlucoseC 6 H 12 O 6 (s)–2808 OctaneC 8 H 18 (l)–5471 SucroseC 12 H 22 O 11 (s)–5645

39. Which of the following thermochemical equations represents an endothermic reaction? A.C graphite (s) + 2 kJ C diamond (s) B.2H 2 (g) + O 2 (g) 2H 2 O + 483.6 kJ

40. Vocab. from the book: Heats of reaction Heats of combustion Molar heat of fusion (melting) Molar heat of solidification (freezing) ΔH fus = -ΔH solid

41. Vocab. from the book: Molar heat of vaporization Molar heat of condensation ΔH vap = -ΔH cond Molar heat of solution (dissolving)

42. Assignment: #s 14 and 15 on page 567 (show work), #s 16, 17, 20, and 21 on page 568 (show work) Day 4 2-26 ADD: # 47, 53, 56, 59, and 70 on pages 586-587 Day 5 2-27

43. http://wps.prenhall.co m/esm_brown_chemist ry_9/2/660/169060.cw/i ndex.html Homework # 2 – w/ discussion partner – show me successful screen

44. 43 Spontaneity Main question of chemistry: Will a reaction GO! when reactants are put together? Spontaneous reactions Reaction occurs by itself, activation E may be needed May have fast, moderate, or slow rate Non-spontaneous reactions Require constant E supply to occur Reaction stops when E supply is removed Reversing reactions If a reaction is spontaneous under specified conditions, the reverse is non-spontaneous at those same conditions

45. 44 Predicting Sponaneity Why is  H a pretty good predictor of spontaneity? -  H : heat out +  H: heat supply needed When might  H not indicate spontaneity well?

46. 45 Predicting Spontaneity H 2 O (s) → H 2 O ( l ) ∆H = + 6.0 kJ; 1 atm, 0°C

47. 46 Predicting Spontaneity H 2 O ( l ) → H 2 O (g ) ∆H = + 40.7 kJ; 1 atm, 100°C

48. 47 Predicting Spontaneity CaCO 3(s) → CaO (s) + CO 2(g) ∆H = +178.0 kJ; 298 K vs. 1100 K

49. 48 Predicting Spontaneity

50. 49 Entropy Entropy is the measure of the disorder or randomness of a system. Entropy is a state function. Entropy changes follow Hess’ Law. A system has high entropy if it _______________________________________ Is Very dsorigniaezd (Very disorganized) Has freedom of motion

51. 50 Entropy & Spontaneity Reactions will be spontaneous if ________ is increased. Examples ice cubes melting vs. putting water in freezer to make ice cubes room gets messy vs. doing work to put things in order etnopry

52. 51 Standards & Constants Conditions & Units T = 25 o C = 298K P = 1 atm Units: J/K = Joules/Kelvin What is the relationship between entropy and temperature? As T  0 K, S  0 T ↓, S ↓ and T ↑, S ↑

53. 52 Standards & Constants Why are the entropy values for elements and compounds always positive? T always > 0K Therefore everything is always moving How do entropies compare among the phases of matter? S o solid < S o liquid < S o gas

54. 53 Standard Entropy Changes The equation ∆ S o =  S o products –  S o reactants Example: 2HCI (aq) + 2Ag (s) → 2AgCl (s) + H 2(g) 1 atm, 25 o C Predict if this reaction will have + ∆S o or – ∆S o. Calculate ∆S o.

55. 54 Standard Entropy Changes The equation  S o =  nS o products –  mS o reactants Example: 2 H 2(g) + O 2(g) → 2 H 2 O (g) 1 atm, 25 o C Predict if this reaction will have + ∆S o or – ∆S o. Calculate ∆S o.