1. Copyright 2000, PRENTICE HALL Chapter 51 These are the remaining slides from class, but beginning with slide 22, there is another presentation from another instructor.

2. Copyright 2000, PRENTICE HALL Chapter 52 For each of the following, define a system and its surroundings and give the direction of energy transfer. a)Methane is burning in a Bunsen burner in a laboratory. b)Water drops, sitting on your skin after swimming, evaporate. c)Two chemicals mixing in a beaker give off heat. React 4

3. Copyright 2000, PRENTICE HALL Chapter 53 Hydrogen gas and oxygen gas react violently to form water. Which is lower in energy: a mixture of hydrogen and oxygen gases, or water? Sketch an energy level diagram for this reaction. React 5

4. Copyright 2000, PRENTICE HALL Chapter 54 Gas A 2 reacts with gas B 2 react to form gas AB. The bond energy of AB is much greater than the bond energy of A 2 or B 2. Is the formation of AB endothermic or exothermic? Which is lower in energy, the reactants or the products? React 6

5. Copyright 2000, PRENTICE HALL Chapter 55 Which of the following performs more work? A gas expanding against a pressure of 2 atm from 1.0 L to 4.0 L A gas expanding against a pressure of 3 atm from 1.0 L to 3.0 L React 7

6. Copyright 2000, PRENTICE HALL Chapter 56 Determine the sign of  E for each of the following with the listed conditions: An endothermic process that performs work. – *work* > *heat* – *work* < *heat* Work is done on a gas and the process is exothermic. – *work* > *heat* – *work* < *heat* React 8

7. Copyright 2000, PRENTICE HALL Chapter 57 Enthalpy and Calorimetry

8. Copyright 2000, PRENTICE HALL Chapter 58 A Coffee Cup Calorimeter Made of Two Styrofoam Cups

9. Copyright 2000, PRENTICE HALL Chapter 59 A 100.0 g sample of water at 90°C is added to a 100.0 g sample of water at 10°C. The final temperature of the water is: a) Between 50°C and 90°C. b) 50°C c) Between 10°C and 50°C. Calculate the final temperature of the water. React 9

10. Copyright 2000, PRENTICE HALL Chapter 510 A 100.0 g sample of water at 90°C is added to a 500.0 g sample of water at 10°C. The final temperature of the water is: a) Between 50°C and 90°C. b) 50°C c) Between 10°C and 50°C. Calculate the final temperature of the water. React 10

11. Copyright 2000, PRENTICE HALL Chapter 511 You have a Styrofoam cup with 50.0 g of water at 10  C. You add a 50.0 g iron ball at 90  C to the water. The final temperature of the water is: a) Between 50°C and 90°C. b) 50°C c) Between 10°C and 50°C. Calculate the final temperature of the water. React 11

12. Copyright 2000, PRENTICE HALL Chapter 512 Change in Enthalpy Can be calculated from enthalpies of formation of reactants and products.  H rxn ° =  n p  H f  (products) -  n r  H f  (reactants)

13. Copyright 2000, PRENTICE HALL Chapter 513 Hess’s Law

14. Copyright 2000, PRENTICE HALL Chapter 514 The Principle of Hess’s Law

15. Copyright 2000, PRENTICE HALL Chapter 515 Standard Enthalpies of Formation

16. Copyright 2000, PRENTICE HALL Chapter 516 Standard States Compound –For a gas, pressure is exactly 1 atmosphere. –For a solution, concentration is exactly 1 molar. –Pure substance (liquid or solid) Element –The form [N 2 (g), K(s)] in which it exists at 1 atm and 25°C.

17. Copyright 2000, PRENTICE HALL Chapter 517 A Schematic Diagram of the Energy Changes for the Reaction CH 4 (g) + 2O 2 (g)  CO 2 (g) + 2H 2 O(l)

18. Copyright 2000, PRENTICE HALL Chapter 518 Using the following data, calculate the heat of formation of the compound ICl (g) at 25°C, and show your work:  H° (kJ/mol) Cl 2 (g)  2Cl (g)242.3 I 2 (g)  I (g)151.0 ICl (g)  I (g) + Cl (g)211.3 I 2 (s)  I 2 (g)62.8 React 12

19. Copyright 2000, PRENTICE HALL Chapter 519 Present Sources of Energy

20. Copyright 2000, PRENTICE HALL Chapter 520 Energy Sources Used in the United States

21. Copyright 2000, PRENTICE HALL Chapter 521 The Earth’s Atmosphere

22. Copyright 2000, PRENTICE HALL Chapter 522 Thermochemistry Chapter 6

23. Copyright 2000, PRENTICE HALL Chapter 523 The Nature of Energy Thermodynamics is the study of energy and its transformations. Thermochemistry is the study of the relationships between chemical reactions and energy changes. Kinetic Energy and Potential Energy –Kinetic energy is the energy of motion: –Potential energy is the energy an object possesses by virtue of its position.

24. Copyright 2000, PRENTICE HALL Chapter 524 Units of Energy SI unit is the joule, J. Traditionally, we use the calorie as a unit of energy. 1 cal = 4.184 J (exactly) The nutritional Calorie, Cal = 1000 cal.

25. Copyright 2000, PRENTICE HALL Chapter 525 Exercise Suppose you consume a six-pack of diet coke (1 nutritional Calorie per can). Convert the energy content of the six-pack to kJ.

26. Copyright 2000, PRENTICE HALL Chapter 526 System and Surroundings A system is the part of the universe we are interested in studying. Surroundings are the rest of the universe (i.e., the surroundings are the portions of the universe not involved in the system). Example: If we are interested in the interaction between hydrogen and oxygen in a cylinder, then the H 2 and O 2 in the cylinder form a system.

27. Copyright 2000, PRENTICE HALL Chapter 527 Transferring Energy: Work and Heat Force is a push or pull on an object. Work is the amount of force applied to an object over a distance: w = F d Heat is the energy transferred from a hotter object to a colder one. Energy is the capacity to do work or to transfer heat.

28. Copyright 2000, PRENTICE HALL Chapter 528 The first law of thermodynamics states that energy cannot be created or destroyed. The first law of thermodynamics is the law of conservation of energy. –That is, the energy of (system + surroundings) is constant. –Thus any energy transferred from a system must be transferred to the surroundings (and vice versa). The First Law of Thermodynamics

29. Copyright 2000, PRENTICE HALL Chapter 529 Internal Energy The total energy of a system is called the internal energy. –It is the sum of all the kinetic and potential energies of all components of the system. Absolute internal energy cannot be measured, only changes in internal energy. Change in internal energy;  E = E final - E initial.

30. Copyright 2000, PRENTICE HALL Chapter 530 Example A mixture of H 2 (g) and O 2 (g) has a higher internal energy than H 2 O(g). Going from H 2 (g) and O 2 (g) to H 2 O(g) results in a negative change in internal energy: H 2 (g) + O 2 (g)  2H 2 O(g)  E < 0 Going from H 2 O(g) to H 2 (g) and O 2 (g) results in a positive change in internal energy: 2H 2 O  H 2 (g) + O 2 (g)  E > 0

31. Copyright 2000, PRENTICE HALL Chapter 531 Heat and Work

32. Copyright 2000, PRENTICE HALL Chapter 532 Relating  E to Heat and Work From the first law of thermodynamics: –When a system undergoes a physical or chemical change, the change in internal energy is given by the heat added to or absorbed by the system plus the work done on or by the system:  E = q + w Heat flowing from the surroundings to the system is positive (i.e., the system feels cold to the touch because it is absorbing heat from your hand), q > 0. Work done by the surroundings on the system is positive, w > 0.

33. Copyright 2000, PRENTICE HALL Chapter 533 Endothermic and Exothermic Processes An endothermic process is one that absorbs heat from the surroundings. An endothermic reaction feels cold. An exothermic process is on that transfers heat to the surroundings. An exothermic reaction feels hot.exothermic reaction

34. Copyright 2000, PRENTICE HALL Chapter 534 Exercise 5.16 For the following process, calculate the change in the internal energy of the system and determine if the process is exothermic or endothermic. (a). A balloon is heated by adding 670 J of heat. It expands, doing 332 J of work on the atmosphere.

35. Copyright 2000, PRENTICE HALL Chapter 535 State Functions A state function depends only on the initial and final states of a system. Example: The altitude difference between Denver and Chicago does not depend on whether you fly or drive, only on the elevation of the two cities above sea level. Similarly, the internal energy of 50 g of H 2 O(l) at 25  C does not depend on whether we cool 50 g of H 2 O(l) at 100  C or heat 50 g of H 2 O(l) at 0  C.

36. Copyright 2000, PRENTICE HALL Chapter 536 State Functions A state function does not depend on how the internal energy is used. Example: A battery in a flashlight can be discharged by producing heat and light. The same battery in a toy car is used to produce heat and work. The change in internal energy of the battery is the same in both cases.

37. Copyright 2000, PRENTICE HALL Chapter 537 Enthalpy The heat transferred between the system and the surroundings during a chemical reaction carried out under constant pressure is called enthalpy, H. Again, we can only measure the change in enthalpy,  H. Mathematically,  H = H final - H initial = q p

38. Copyright 2000, PRENTICE HALL Chapter 538 Heat transferred from the surroundings to the system has a positive enthalpy (i.e.,  H > 0 for an endothermic reaction). Heat transferred from the system to the surroundings has a negative enthalpy (i.e.,  H < 0 for an exothermic reaction). Enthalpy is a state function. Enthalpy

39. Copyright 2000, PRENTICE HALL Chapter 539 Energy, Enthalpy, and P-V Work ConsiderConsider: A cylinder of cross- sectional area A, A piston exerting a pressure, P = F/A, on a gas inside the cylinder,

40. Copyright 2000, PRENTICE HALL Chapter 540 The volume of gas expanding through  V while the piston moves a height  h = h f – h i, The magnitude of work done = F x  h = P x A x  h = P x  V. Since work is being done by the system on the surroundings, w = - P  V From the first law of thermodynamics,  E = q - P  V If the reaction is carried out under constant volume,  V = 0 and  E = q v. If the reaction is carried out under constant pressure,  E = q p - P  V, or q p =  H =  E + P  V and  E =  H - P  V Energy, Enthalpy, and P-V Work

41. Copyright 2000, PRENTICE HALL Chapter 541 Enthalpies of Reaction The enthalpy change that accompanies a reaction is called the enthalpy of reaction. For a reaction,  H rxn = H(products) - H(reactants). Consider the thermochemical equation for the production of water: 2H 2 (g) + O 2 (g)  2H 2 O(g)  H = -483.6 kJ The equation tells us that 483.6 kJ of energy is released to the surroundings when water is formed. Enthalpy is an extensive property. Therefore, the magnitude of enthalpy is directly proportional to the amount of reactant consumed.

42. Copyright 2000, PRENTICE HALL Chapter 542 Example: If 1 mol of CH 4 is burned in oxygen to produce CO 2 and water, 890 kJ of heat is released to the surroundings. If 2 mol of CH 4 is burned, then 1780 kJ of heat is released. The sign of  H depends on the direction of the reaction. The enthalpy changes for a reaction and its reverse reaction are equal in magnitude but opposite in sign. Enthalpies of Reaction

43. Copyright 2000, PRENTICE HALL Chapter 543 CH 4 (g) + 2O 2 (g)  CO 2 (g) + 2H 2 O(l)  H = -890 kJ, but CO 2 (g) + 2H 2 O(l)  CH 4 (g) + 2O 2 (g)  H = +890 kJ. Enthalpy change depends on state. 2H 2 O(g)  2H 2 O(l)  H = -88 kJ Enthalpies of Reaction

44. Copyright 2000, PRENTICE HALL Chapter 544 Exercise 5.28 Consider the following reaction: CH 3 OH(g)  CO(g) + 2H 2 (g)  H = + 90.7 kJ (a). Is the reaction exothermic or endothermic? (b). Calculate the amount of heat transferred when 45.0 g of methanol decomposes by this reaction at constant pressure. (c). If the enthalpy change is 16.5 kJ, how many grams of hydrogen gas are produced?

45. Copyright 2000, PRENTICE HALL Chapter 545 Heat Capacity and Specific Heat Heat capacity is the amount of energy required to raise the temperature of an object by 1°C. Molar heat capacity is the heat capacity of 1 mol of a substance. Specific heat, or specific heat capacity is the amount of energy required to raise the temperature of 1 g of a substance by 1°C. Heat, q = (specific heat) x (grams of substance) x  T. Be careful of the sign of q.

46. Copyright 2000, PRENTICE HALL Chapter 546 Calorimetry Calorimetry is a measurement of heat flow. A calorimeter is an apparatus that measures heat flow. Constant-Pressure Calorimetry Most common technique: use atmospheric pressure as the constant pressure. Recall that  H = q p. Easiest method: use a coffee-cup calorimeter. q soln = (specific heat of solution) x (grams of solution) x  T = -q rxn

47. Copyright 2000, PRENTICE HALL Chapter 547 Bomb Calorimetry (Constant-Volume Calorimetry) Reactions can be carried out under conditions of constant volume instead of constant pressure. Constant-volume calorimetry is carried out in a bomb calorimeter. Calorimetry

48. Copyright 2000, PRENTICE HALL Chapter 548 The most common type of reaction studied under these conditions is combustion. If we know the heat capacity of the calorimeter, C cal, then the heat of reaction, q rxn = -C cal x  T. Since the reaction is carried out under constant volume, q corresponds to  E rather than  H. For most reactions the difference between  E and  H is small. Calorimetry

49. Copyright 2000, PRENTICE HALL Chapter 549 Exercise 5.40 The specific heat of toluene, C 7 H 8, is 1.13J/g-K. How many joules of heat are needed to raise the temperature of 40.0 g of toluene from 10.4 o C to 28.0 o C?

50. Copyright 2000, PRENTICE HALL Chapter 550 Exercise 5.43 A 2.200 g sample of quinone (C 6 H 4 O 2 ) is burned in a bomb calorimeter with a total heat capacity of 7.854kJ/ o C. The temperature of the calorimeter increases from 23.44 o C to 30.57 o C. What is the heat of combustion per gram of quinone? What is the heat of combustion per mole of quinone?

51. Copyright 2000, PRENTICE HALL Chapter 551 Hess’s Law If a reaction is carried out in a series of steps,  H for the reaction is the sum of the  H’s for each of the steps. The total change in enthalpy is independent of the number of steps. Total  H is also independent of the nature of the path.

52. Copyright 2000, PRENTICE HALL Chapter 552 Hess’s Law Example: CH 4 (g) + 2O 2 (g)  CO 2 (g) + 2H 2 O(g)  H = -802 kJ 2H 2 O(g)  2H 2 O(l)  H = -88 kJ ______________________________________________________________ CH 4 (g) + 2O 2 (g)  CO 2 (g) + 2H 2 O(l)  H = -890 kJ Therefore, for the reaction CH 4 (g) + 2O 2 (g)  CO 2 (g) + 2H 2 O(l),  H= -890 kJ. Note that  H is sensitive to the states of the reactants and products. Hess’s law allows us to calculate enthalpy data for reactions that are difficult to carry out directly: C(s) + O 2 (g) produces a mixture of CO(g) and CO 2 (g).

53. Copyright 2000, PRENTICE HALL Chapter 553 Exercise 5.52 From the following heats of reaction 2H 2 (g) + O 2 (g)  2H 2 O(g)  H= -483.6 kJ 3O 2 (g)  2O 3 (g)  H= +284.6 kJ calculate the enthalpy for the reaction 3H 2 (g) + O 3 (g)  3H 2 O(g)

54. Copyright 2000, PRENTICE HALL Chapter 554 Enthalpies of Formation Consider the formation of CO 2 (g) and 2H 2 O(l) from CH 4 (g) and 2O 2 (g). If the reaction proceeds in one step: CH 4 (g) + 2O 2 (g)  CO 2 (g) + 2H 2 O(l), then  H 1 = -890 kJ. However, if the reaction proceeds through a CO intermediate: CH 4 (g) + 2O 2 (g)  CO(g) + 2H 2 O(l) + ½O 2 (g)  H 2 = -607 kJ CO(g) + 2H 2 O(l) + ½ O 2 (g)  CO 2 (g) + 2H 2 O(l)  H 3 = -283 kJ, Then  H for the overall reaction is:  H 2 +  H 3 = -607 kJ - 283 kJ = -890 kJ =  H 1

55. Copyright 2000, PRENTICE HALL Chapter 555 If a compound is formed from its constituent elements, then the enthalpy change for the reaction is called the enthalpy of formation,  H f. Standard conditions (standard state) refer to the substance at in its pure form at 1 atm at the temperature of interest. We usually work with 25°C (298 K). Standard enthalpy,  H°, is the enthalpy measured when everything is in its standard state. Enthalpies of Formation

56. Copyright 2000, PRENTICE HALL Chapter 556

57. Copyright 2000, PRENTICE HALL Chapter 557 Standard enthalpy of formation of a compound,  H° f, is the enthalpy change for the formation of 1 mol of compound with all substances in their standard states. If there is more than one state for a substance under standard conditions, the most stable state is used. Example: When dealing with carbon we use graphite because graphite is more stable than diamond or C 60. The standard enthalpy of formation of the most stable form of an element is zero. Enthalpies of Formation

58. Copyright 2000, PRENTICE HALL Chapter 558 Using Enthalpies of Formation to Calculate Enthalpies of Reaction Example: Calculate  H for C 3 H 8 (g) + 5O 2 (g)  3CO 2 (g) + 4H 2 O(l)

59. Copyright 2000, PRENTICE HALL Chapter 559 We start with the reactants, decompose them into elements, then rearrange the elements to form products. The overall enthalpy change is the sum of the enthalpy changes for each step (Hess’s Law). Decomposing into elements (note O 2 is already elemental, so we concern ourselves with C 3 H 8 ); we have: C 3 H 8 (g)  3C(s) + 4H 2 (g)  H 1 = -  H° f [C 3 H 8 (g)] Next we form CO 2 and H 2 O from their elements: 3C(s) + 3O 2 (g)  3CO 2 (g)  H 2 = 3  H° f [CO 2 (g)] 4H 2 (g) + 2O 2 (g)  4H 2 O(l)  H 3 = 4  H° f [H 2 O(l)] Using Enthalpies of Formation to Calculate Enthalpies of Reaction

60. Copyright 2000, PRENTICE HALL Chapter 560 We look up the values and add:  H° rxn = -1(-103.85 kJ) + 3(-393.5 kJ) + 4(-285.8 kJ) = -2220 kJ Notice that CO 2 and H 2 O are products and propane is the reaction ! In general:  H° rxn =  n  H° f (products) -  m  H° f (reactants) Where n and m are the stoichiometric coefficients Using Enthalpies of Formation to Calculate Enthalpies of Reaction

61. Copyright 2000, PRENTICE HALL Chapter 561 Exercise 5.60 Many cigarette lighters contain liquid butane, C 4 H 10 ( l ). Using enthalpies of formation, calculate the quantity of heat produced when 1.0 g of butane is completely combusted in air.

62. Copyright 2000, PRENTICE HALL Chapter 562 Exercise 5.67 FG05_025.JPG One of the cleanest burning octanes is sketched above. The complete combustion of 1 mol of this compound to H 2 O(g) and CO 2 (g) leads to  H o = -5069 kJ. (a). Write a balanced equation for the complete combustion of this compound. (b). By using the information in this problem and the data in Table 5.3, calculate the standard enthalpy of formation of this compound at 298 K.

63. Copyright 2000, PRENTICE HALL Chapter 563 Foods and Fuels Fuel value is the energy released when 1 g of substance is burned. Foods Fuel value is usually measured in Calories (1 nutritional Calorie, 1 Cal = 1000 cal). Most energy in our bodies comes from the oxidation of carbohydrates and fats. In the intestines carbohydrates are converted into glucose, C 6 H 12 O 6, or blood sugar.

64. Copyright 2000, PRENTICE HALL Chapter 564 In the cells glucose reacts with O 2 in a series of steps which ultimately produce CO 2, H 2 O, and energy. C 6 H 12 O 6 (s) + 6O 2 (g)  6CO 2 (g) + 6H 2 O(l)  H° = -2803 kJ Fats, for example, tristearin, react with O 2 as follows: 2C 57 H 110 O 6 (s) + 163O 2 (g)  114CO 2 (g) + 110H 2 O(l)  H° = -75,250 kJ Fats contain more energy than carbohydrates. Fats are not water soluble. Therefore, fats are good for energy storage. Foods and Fuels

65. Copyright 2000, PRENTICE HALL Chapter 565 Fuels In the United States we use about 1.0 x 10 6 kJ of energy per person per day. Most of this energy comes from petroleum and natural gas. The remainder of the energy comes from coal, nuclear, and hydroelectric sources. Coal, petroleum, and natural gas are fossil fuels. They are not renewable.

66. Copyright 2000, PRENTICE HALL Chapter 566 Natural Gas Natural gas consists largely of carbon and hydrogen. Compounds such as CH 4, C 2 H 6, C 3 H 8, and C 4 H 10 are typical constituents. Petroleum is a liquid consisting of hundreds of compounds. Impurities include S, N, and O compounds. Coal contains high molecular weight compounds of C and H. In addition compounds containing S, O, and N are present as impurities. Common fuels have typical fuel values around 30 kJ/g. Hydrogen has great potential as a fuel with a fuel value of 142 kJ/g.

67. Copyright 2000, PRENTICE HALL Chapter 567 Other Energy Sources Nuclear energy: energy released in splitting or fusion of nuclei of atoms. Fossil fuels and nuclear energy are nonrenewable sources of energy. Renewable energy sources include: –Solar energy –Wind energy –Geothermal energy –Hydroelectric energy –Biomass energy These are virtually inexhaustible and will become increasingly important as fossil fuels are depleted.