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1. Define the following term: system. A) The part of the universe that is of interest to us. B) Heat transfer. C) The part of the universe surrounding a reaction. D) The universe. 2. Define the following term: surroundings. A) The part of the universe that is of interest to us. B) Heat transfer. C) The part of the universe surrounding the system. D) The universe. 3. Define the following term: open system. A) A system, which can exchange mass and energy with its surroundings. B) A system, which allows the transfer of energy but not mass with its surroundings. C) A system, which does not allow the transfer of either mass or energy. 4. Define the following term: closed system. A) A system, which can exchange mass and energy with its surroundings. B) A system, which allows the transfer of energy but not mass with its surroundings. C) A system, which does not allow the transfer of either mass or energy. 5. Define the following term: isolated system. A) A system, which can exchange mass and energy with its surroundings. B) A system, which allows the transfer of energy but not mass with its surroundings. C) A system, which does not allow the transfer of either mass or energy. 6. Define the following term: thermal energy. A) The capacity to do work. B) The energy associated with the random motion of atoms and molecules. C) The energy stored within the structural units of chemical substances. D) The energy, which comes from the sun. 7. Define the following term: chemical energy. A) The capacity to do work. B) The energy associated with the random motion of atoms and molecules. C) The energy stored within the structural units of chemical substances. D) The energy, which comes from the sun. Page 1 8. Define the following term: potential energy. A) The energy available by virtue of an object's position. B) The energy associated with the random motion of atoms and molecules. C) The energy produced by a moving object D) The energy, which comes from the sun. 9. Define the following term: kinetic energy. A) The energy available by virtue of an object's position. B) The energy stored within the structural units of chemical substances. C) The energy produced by a moving object D) The energy, which comes from the sun. 10. Define the following term: law of conservation of energy. A) The capacity to do work. B) Directed energy change resulting from a process. C) The transfer of thermal energy between two bodies that are at different temperatures. D) The total quantity of energy in the universe is assumed constant. 11. What unit of energy is most commonly employed in chemistry? A) MJ B) kJ C) cal D) kcal 12. A truck traveling at 60 kilometers per hour is brought to a complete stop at a traffic light. Does this change violate the law of conservation of energy? Explain. A) No, kinetic energy is transferred into heat through friction. B) No, kinetic energy is transferred into potential energy. C) Yes, kinetic energy should remain constant. 13. Which of the following would be an example of the interconversion between chemical and electrical energy? A) Coal is burned to produce electricity. B) Methane is burned to heat a home. C) Air turns a windmill to generate electricity. D) A switch is turned on and a lamp lights up. Page 2 14. Which of the following would be an example of the interconversion between chemical and thermal energy? A) Coal is burned to produce electricity. B) Methane is burned to heat a home. C) Air turns a windmill to generate electricity. D) A switch is turned on and a lamp lights up. 15. Which of the following would be an example of the interconversion between mechanical and electrical energy? A) Coal is burned to produce electricity. B) Methane is burned to heat a home. C) Air turns a windmill to generate electricity. D) A switch is turned on and a lamp lights up. 16. Which of the following would be an example of the interconversion between electrical energy and light? A) Coal is burned to produce electricity. B) Methane is burned to heat a home. C) Air turns a windmill to generate electricity. D) A switch is turned on and a lamp lights up. 17. Define the following term: thermochemistry. A) The study of heat change in chemical reactions B) The process of transferring thermal energy from a system to the surroundings. C) The process of transferring thermal energy from the surroundings to a system. D) The transfer of thermal energy between two bodies that are at different temperatures. 18. Define the following term: exothermic process. A) The study of heat change in chemical reactions B) The process of transferring thermal energy from a system to the surroundings. C) The process of transferring thermal energy from the surroundings to a system. D) The transfer of thermal energy between two bodies that are at different temperatures. Page 3 19. Define the following term: endothermic process. A) The study of heat change in chemical reactions B) The process of transferring thermal energy from a system to the surroundings. C) The process of transferring thermal energy from the surroundings to a system. D) The transfer of thermal energy between two bodies that are at different temperatures. 20. Think about what is meant by a state function. Give two examples of quantities that are state functions and two that are not. A) State functions: P,V; Not: work, energy B) State functions: energy, P; Not: T, heat C) State functions: energy, V; Not: work, heat D) State functions: energy, work; Not: T, heat Use the following to answer questions 21-23: The internal energy of an ideal gas depends only on its temperature. Do a first-law analysis of the following process. A sample of an ideal gas is allowed to expand at constant temperature against atmospheric pressure. 21. Does the gas do work on its surroundings? A) Yes B) No C) Cannot be determined. 22. Is there heat exchange between the system and the surroundings? If so, in which direction? A) No B) Yes, heat is absorbed by the system. C) Yes, heat is given off by the system. 23. What is E for the gas in this process? A) Negative B) Zero C) Positive Page 4 24. The following reaction takes place: Hg(l) Hg(g) Describe the work involved in this reaction. (No calculations are necessary.) A) Work is done by the system on the surroundings. B) Work is done by the surroundings on the system. C) No work is done. 25. The following reaction takes place: 3O2(g) 2O3(g) Describe the work involved in this reaction. (No calculations are necessary.) A) Work is done by the system on the surroundings. B) Work is done by the surroundings on the system. C) No work is done. 26. The following reaction takes place: CuSO4 5H2O(s) CuSO4(s) + 5H2O(g) Describe the work involved in this reaction. (No calculations are necessary.) A) Work is done by the system on the surroundings. B) Work is done by the surroundings on the system. C) No work is done. 27. The following reaction takes place: H2(g) + F2(g) 2HF(g) Describe the work involved in this reaction at constant pressure. (No calculations are necessary.) A) Work is done by the system on the surroundings. B) Work is done by the surroundings on the system. C) No work is done. 28. A sample of nitrogen gas expands in volume from 1.6 L to 5.4 L at constant temperature. Calculate the work done in joules if the gas expands against a vacuum. A) 4.1 J B) 420 J C) –4.1 J D) 0 J Page 5 29. A sample of nitrogen gas expands in volume from 1.6 L to 5.4 L at constant temperature. Calculate the work done in joules if the gas expands against a constant pressure of 0.80 atm. A) –3.0J B) 310 J C) –310 J D) 3.8 J 30. A sample of nitrogen gas expands in volume from 1.6 L to 5.4 L at constant temperature. Calculate the work done in joules if the gas expands against a constant pressure of 3.7 atm. A) 14 J B) –14 J C) 1400 J D) –1400 J 31. A gas expands in volume from 26.7 mL to 89.3 mL at constant temperature. Calculate the work done (in joules) if the gas expands against a vacuum. A) 0 J B) 63 J C) 6300 J D) –6300 J 32. A gas expands in volume from 26.7 mL to 89.3 mL at constant temperature. Calculate the work done (in joules) if the gas expands against a constant pressure of 1.5 atm. A) –63 J B) –0.094 J C) 9.5 J D) –9.5 J 33. A gas expands in volume from 26.7 mL to 89.3 mL at constant temperature. Calculate the work done (in joules) if the gas expands against a constant pressure of 2.8 atm. A) –18000 J B) –18 J C) 0.18 J D) –0.18 J Page 6 34. A gas expands and does P-V work on the surroundings equal to 325 J. At the same time, it absorbs 127 J of heat from the surroundings. Calculate the change in energy of the gas. A) 452 J B) 198 J C) –198 J D) –452 J 35. The work done to compress a gas is 74 J. As a result, 26 J of heat is given off to the surroundings. Calculate the change in energy of the gas. A) 100 kJ B) 48 kJ C) –48 kJ D) –100 kJ 36. Calculate the work done when 50.0 g of tin are dissolved in excess acid at 1.00 atm and 25°C: Sn(s) + 2H+(aq) Sn2+(aq) + H2(g) Assume ideal gas behavior. A) 1.04 103 J B) 9.53 102 J C) –9.53 102 J D) –1.04 103 J 37. Calculate the work done in joules when 1.0 mole of water vaporizes at 1.0 atm and 100°C. Assume that the volume of liquid water is negligible compared with that of steam at 100°C and ideal gas behavior. A) 3.1 103 J B) –8.4 102 J C) –3.1 103 J D) 31 J 38. Define the following term: enthalpy. A) Heat transfer at constant temperature. B) E + PV C) Heat flow under all conditions. D) Heat transfer at constant volume. Page 7 39. Define the following term: enthalpy of reaction. A) The heat transfer from a system to a reaction. B) The transfer from a reaction to the surroundings. C) The difference between the enthalpies of the products and the enthalpies of the reactants. D) The endothermic process of breaking chemical bonds. 40. Under what condition is the heat of a reaction equal to the enthalpy change of the same reaction? A) Constant volume. B) Constant temperature. C) Constant pressure. D) Constant volume and temperature. Use the following to answer questions 41-43: Consider this reaction: 2CH3OH(l) + 3O2(g) 4H2O(l) + 2CO2(g) H = –1452.8 kJ/mol 41. What is the value of H if the equation is multiplied throughout by 2? A) –1452.8 kJ B) –2905.6 kJ C) -726.4 kJ 42. What is the value of H if the reaction is reversed so that the products become the reactants and vice versa? A) 0 kJ B) –1452.8 kJ C) 1452.8 kJ 43. What would the value of H be if the product in the reaction is water vapor instead of liquid water? The enthalpy of vaporization of water is 44 kJ/mol. A) –1628 kJ B) –1452.8 kJ C) –1408.8 kJ D) –1276.8 kJ Page 8 44. Consider the reaction 2H2O(g) 2H2(g) + O2(g) H = 483.6 kJ/mol If 2.0 moles of H2O(g) are decomposed against a pressure of 1.0 atm at 125 °C, what is E for this reaction? A) 480.3 kJ B) –2825 kJ C) 483.6 kJ D) –483.6 kJ 45. Consider the reaction H 0rxn H2(g) + Cl2(g) 2HCl(g) = 184.6 kJ/mol If 3 moles of H2 react with 3 moles of Cl2 to form HCl at a pressure of 1.0 atm and 25°C, then what is E for this reaction? Assume the reaction goes to completion. (Hint: First calculate the work done in kJ.) A) –568.7 kJ B) –558.8 kJ C) –553.8 kJ D) –184.6 kJ 46. Consider two metals A and B, each having a mass of 100 g and an initial temperature of 20°C. The specific heat of A is larger than that of B. Under the same heating conditions, which metal would take longer to reach a temperature of 21°C? A) Metal A B) Metal B C) Same amount of time 47. Consider the following data: Metal Al Cu Mass (g) 10 30 Specific Heat (J/g°C) 0.900 0.285 Temperature (°C) 40 60 these two metals are placed in contact, which of the following will take place? A) Heat will flow from Al to Cu because Al has a larger specific heat. B) Heat will flow from Cu to Al because Cu has a larger mass. C) Heat will flow fro Cu to Al because Cu has a larger heat capacity. D) Heat will flow from Cu to Al because Cu is at a higher temperature. E) No heat will flow in either direction. Page 9 48. A piece of silver of mass 362 g has a heat capacity of 85.7 J/°C. What is the specific heat of silver? A) 31,000 J/g ºC B) 4.22 J/g ºC C) 0.237 J/g ºC 49. A 6.22-kg piece of copper metal is heated from 20.5°C to 324.3°C. Calculate the heat absorbed (in kJ) by the metal. The specific heat of Cu is 0.385 J/ gºC. A) 7.28 10–4 kJ B) 0.728 kJ C) 728 kJ D) 7.28 105 kJ 50. Calculate the amount of heat liberated (in kJ) from 366 g of mercury when it cools from 77.0°C to 12.0°C. The specific heat of Hg is 0.139 J/ gºC. A) –3.31 kJ B) 3.31 kJ C) –10.6 kJ D) 10.6 kJ 51. A sheet of gold weighing 10.0 g and at a temperature of 18.0°C is placed flat on a sheet of iron weighing 20.0 g and at a temperature of 55.6°C. What is the final temperature of the combined metals? Assume that no heat is lost to the surroundings. (Hint: The heat gained by the gold must be equal to the heat lost by the iron.) The specific heats of Au and Fe are 0.129 J/ gºC and 0.444 J/ gºC, respectively. A) 46.3°C. B) 50.7°C. C) 36.8°C. D) 62.0°C. Use the following to answer questions 52-53: A 0.1375-g sample of solid magnesium is burned in a constant-volume bomb calorimeter that has a heat capacity of 3024 J/°C. The calorimeter contains exactly 300 g of water, and the temperature increases by 1.126°C. Page 10 52. Calculate the heat given off by the burning Mg, in kJ/g. A) 48.27 kJ/g Mg B) 3.393 kJ/g Mg C) 3.502 104 kJ/g Mg D) 24.76 kJ/g Mg 53. Calculate the heat given off by the burning Mg in kJ/mol. A) 1.14 103 kJ/mol Mg B) 82.4 kJ/mol Mg C) 8.52 103 kJ/mol Mg D) 602 kJ/mol Mg 54. A quantity of 2.00 102 mL of 0.862 M HCl is mixed with 2.00102 mL of 0.431 M Ba(OH)2 in a constant-pressure calorimeter that has a heat capacity (empty) of 453 J/°C. The initial temperature of the HCl and Ba(OH)2 solutions is the same at 20.48°C. For the process H+(aq) + OH(aq) H2O(l), the heat of neutralization is 56.2 kJ/mol. What is the final temperature of the mixed solution? A) 15.9 ºC B) 25.0 ºC C) 29.6 ºC D) 43.2 ºC 55. Which of the following standard enthalpy of formation values is not zero at 25°C? Na(s), Ne(g), CH4(g), S8(s), Hg(l), H(g). A) H(g) B) Ne(g), CH4(g), H(g) C) Na(s), S8(s) D) Hg(l) 56. H 0f The values of the two allotropes of oxygen, O2 and O3, are 0 and 142.2 kJ/mol, respectively, at 25°C. Which is the more stable form at this temperature? A) Both O2 and O3 are equally stable. B) O3 C) O2 D) Neither is stable. Page 11 57. Which is the more negative quantity at 25°C: A) They are equal. B) H2O(l) C) H2O(g) H 0f for H2O(l) or H 0f for H2O(g)? 58. H 0f Predict the value of (greater than, less than, or equal to zero) for Br2(g) at 25°C. A) Greater than zero. B) Less than zero. C) Equal to zero. 59. H 0f Predict the value of (greater than, less than, or equal to zero) for : Br2(l) at 25°C. A) Greater than zero. B) Less than zero. C) Equal to zero. 60. H 0f Predict the value of (greater than, less than, or equal to zero) for I2(g) at 25°C. A) Greater than zero. B) Less than zero. C) Equal to zero. 61. H 0f Predict the value of (greater than, less than or equal to zero) for I2(s) at 25°C. A) Greater than zero. B) Less than zero. C) Equal to zero. 62. H 0f In general, compounds with negative values are more stable than those with 0 H f H 0f positive values. H2O2(l) has a negative . Why, then, does H2O2(l) have a tendency to decompose to H2O(l) and O2(g)? A) Forming O2(g) drives the reaction. B) Forming two products drives the reaction. C) H2O2(l) has a more negative standard enthalpy of formation than does H2O(l). D) H2O(l) has a more negative standard enthalpy of formation than does H2O2(l). Page 12 63. H 0f Suggest a way that would allow you to measure the values of Ag2O(s) and CaCl2(s) from their elements. No calculations are necessary. A) Form the compound from its stable elements. B) Combust the compound and measure values from the backwards reaction. C) Measure the standard enthalpy of reaction. D) Both a and c. E) Both b and c. 64. Calculate the heat of decomposition using enthalpy of formation information for this process at constant pressure and 25°C: CaCO3(s) CaO(s) + CO2(g) A) 964.8 kJ B) 177.8 kJ C) –177.8 kJ D) –964.8 kJ 65. The standard enthalpies of formation of ions in aqueous solutions are obtained by H 0f + arbitrarily assigning a value of zero to H+ ions; that is, [H (aq)] = 0. For the following reaction HCl(g) H+(aq) + Cl(aq) H0 = 74.9kJ/mol H 0f calculate for the Cl ions A) 167.2 kJ/mol B) 17.4 kJ/mol C) 17.4 kJ/mol D) 167.2 kJ/mol 66. The standard enthalpies of formation of ions in aqueous solutions are obtained by H 0f + arbitrarily assigning a value of zero to H+ ions; that is, [H (aq)] = 0. Given that 0 H f for OH¯ ions is 229.6 kJ/mol, calculate the enthalpy of neutralization when 1 mole of a strong monoprotic acid (such as HCl) is titrated by 1 mole of a strong base (such as KOH) at 25°C. A) –56.2 kJ B) –515.4 kJ C) 56.2 kJ D) 515.4 kJ Page 13 67. Calculate the heat of combustion for the following reaction from the standard enthalpies of formation: 2H2(g) + O2(g) 2H2O(l) A) –285.8 kJ B) –571.6 kJ C) 285.8 kJ D) 571.6 kJ 68. Calculate the heat of combustion for the following reaction from the standard enthalpies of formation listed in Appendix 3: 2C2H2(g) + 5O2(g) 4CO2(g) + 2H2O(l) A) 598.8 kJ B) 905.9 kJ C) –2598.8 kJ D) –905.9 kJ 69. Calculate the heat of combustion for the following reaction from the standard enthalpies of formation listed in Appendix 3: C2H4(g) + 3O2(g) 2CO2(g) + 2H2O(l) A) –731.6 kJ B) –1125.1 kJ C) –1410.9 kJ D) –1463.2 kJ 70. Calculate the heat of combustion for the following reaction from the standard enthalpies of formation listed in Appendix 3: 2H2S(g) + 3O2(g) 2H2O(l) + 2SO2(g) A) –561.8 kJ B) –1123.5 kJ C) –1143.7 kJ D) –1204.1 kJ 71. Methanol, ethanol, and n-propanol are three common alcohols. When 1.00 g of methanol (CH3OH) is burned in air, –22.6 kJ of heat is liberated. Calculate the heat of combustion of methanol in kJ/mol. A) 724 kJ/mol B) 702 kJ/mol C) –702 kJ/mol D) –724 kJ/mol Page 14 72. Methanol, ethanol, and n-propanol are three common alcohols. When 1.00 g of ethanol (C2H5OH) is burned in air, –29.7 kJ of heat is liberated. Calculate the heat of combustion of ethanol in kJ/mol. A) –1.37 103 kJ/mol B) –1.34 103 kJ/mol C) 1.34 103 kJ/mol D) 1.37 103 kJ/mol 73. Methanol, ethanol, and n-propanol are three common alcohols. When 1.00 g of npropanol (C3H7OH)is burned in air, –33.4 kJ of heat is liberated. Calculate the heat of combustion of n-propanol in kJ/mol. A) –1.97 103 kJ/mol B) –2.01 103 kJ/mol C) 1.97 103 kJ/mol D) 2.01 103 kJ/mol 74. The standard enthalpy change for the following reaction is 436.4 kJ: H2(g) H(g) + H(g) Calculate the standard enthalpy of formation of atomic hydrogen (H). A) 436.4 kJ/mol B) 218.2 kJ/mol C) –218.2 kJ/mol D) –436.4 kJ/mol 75. From the standard enthalpies of formation, calculate C6H12(l) + 9O2(g) 6CO2(g) + 6H2O(l) H 0f For C6H12(l), = 151.9 kJ/mol. A) 151.9 kJ B) –527.4 kJ C) –3923.9 kJ D) –4227.7 kJ Page 15 H 0rxn for the reaction 76. The first step in the industrial recovery of zinc from the zinc sulfide ore is roasting, that is, the conversion of ZnS to ZnO by heating: H 0rxn 2ZnS(s) + 3O2(g) 2ZnO(s) + 2SO2(g) = 879 KJ/mol Calculate the heat evolved (in kJ) per gram of ZnS roasted. A) 9.02 kJ/g ZnS B) 4.51 kJ/g ZnS. C) –4.51 kJ/g ZnS D) –9.02 kJ/g ZnS 77. Determine the amount of heat (in kJ) given off when 1.26 104 g of ammonia are produced according to the equation H 0rxn N2(g) + 3H2(g) 2NH3(g) = 92.6 kJ/mol Assume that the reaction takes place under standard-state conditions at 25°C. A) –3.43 104 kJ B) –6.86 104 kJ C) 3.43 104 kJ D) 6.86 104 kJ 78. At 850°C, CaCO3 undergoes substantial decomposition to yield CaO and CO2. H 0f Assuming that the values of the reactant and products are the same at 850°C as they are at 25°C, calculate the enthalpy change (in kJ) if 66.8 g of CO2 are produced in one reaction. A) –3.40 103 kJ B) 1.46 103 kJ C) 1.17 102 kJ D) 2.70 102 kJ 79. From these data, S(rhombic) + O2(g) SO2(g) H 0rxn = 296.06 kJ/mol H S(monoclinic) + O2(g) SO2(g) = 296.36 kJ/mol calculate the enthalpy change for the transformation S(rhombic) S(monoclinic) (Monoclinic and rhombic are different allotropic forms of elemental sulfur.) A) 592.42 kJ B) 0.30 kJ C) –0.30 kJ D) –592.42 kJ 0 rxn Page 16 80. From the following data, H 0rxn C(graphite) + O2(g) CO2(g) = 393.5 kJ/mol 0 1 H rxn = 285.8 kJ/mol H2(g) + 2 O2(g) H2O(l) H 0rxn 2C2H6(g) + 7O2(g) 4CO2(g) + 6H2O(l) = 3119.6 kJ/mol calculate the enthalpy change for the reaction 2C(graphite) + 3H2(g) C2H6(g) A) –2417.2 kJ B) –84.6 kJ C) 2417.2 kJ D) –3204.2 kJ 81. Note the following heats of combustion: H 0rxn O2(g) CO2(g) + 2H2O(l) = 726.4 kJ/mol 0 H rxn C(graphite) + O2(g) CO2(g) = 393.5 kJ/mol 0 1 H rxn H2(g) + 2 O2(g) H2O(l) = 285.8 kJ/mol Calculate the enthalpy of formation of methanol (CH3OH) from its elements: CH3OH(l) + 3 2 C(graphite) + 2H2(g) + A) –238.7 kJ/mol B) –47.1 J/mol C) 47.1 kJ/mol D) –2811.4 kJ/mol 1 2 O2(g) CH3OH(l) 82. Calculate the standard enthalpy change for the reaction 2Al(s) + Fe2O3(s) 2Fe(s) + Al2O3(s) given that 3 H 0rxn 2Al(s) + 2 O2(g) Al2O3(s) = 1601 kJ/mol 0 3 H rxn 2Fe(s) + 2 O2(g) Fe2O3(s) = 821 kJ/mol A) 2420 kJ B) 780 kJ C) –780 kJ D) –2420 kJ Page 17 83. Consider two ionic compounds A and B. A has a larger lattice energy than B. Which of the two compounds is more stable? A) A is more stable. B) B is more stable. C) They are equally stable. D) Cannot be determined. 84. Why is it dangerous to add water to a concentrated acid such as sulfuric acid in a dilution process? A) The solution will rapidly get very hot. B) The solution will rapidly get very cold. C) The solution could explode and splatter you with acid. D) Both a and c. E) Both b and c. 85. The convention of arbitrarily assigning a zero enthalpy value for the most stable form of each element in the standard state of 25°C is a convenient way of dealing with enthalpies of reactions. Explain why this convention cannot be applied to nuclear reactions. A) The identity of the atoms changes during the reaction. B) The number of atoms changes during the reaction. C) The intense heat created cannot be measured. D) Both a and b. E) Both b and c. 86. Consider the following two reactions: H 0rxn A 2B = H1 0 H rxn AC = H2 Determine the enthalpy change for the process 2B C A) H2 H1 B) H2 + H1 C) H1 H2 D) 2H1 H2 Page 18 87. The standard enthalpy change H0 for the thermal decomposition of silver nitrate according to the following equation is +78.67 kJ: AgNO3(s) AgNO2(s) + ½ O2(g) The standard enthalpy of formation of AgNO3(s) is 123.02 kJ/mol. Calculate the standard enthalpy of formation of AgNO2(s). A) –201.69 kJ/mol B) –44.35 kj/mol C) 44.35 kJ/mol D) 201.69 kJ/mol 88. Hydrazine, N2H4, decomposes according to the following reaction: 3N2H4(l) 4NH3(g) + N2(g) The standard enthalpy of formation of hydrazine is 50.42 kJ/mol, Calculate H0 for its decomposition. A) 336.5 kJ B) 96.72 kJ C) –96.72 kJ D) –336.5 kJ 89. Hydrazine burns in oxygen to produce H2O(l) and N2(g). Calculate H0 for the burning of one mol of hydrazine. A) 622.0 kJ B) 336.2 kJ C) –336.2 kJ D) –622.0 kJ 90. Ammonia burns in oxygen to produce H2O(l) and N2(g). Calculate H0 for this process when the equation is balanced with the lowest possible whole numbers. A) 240 kJ B) –3240 kJ C) –1530 kJ D) –240 kJ Use the following to answer questions 91-92: Consider the reaction N2(g) + 3H2(g) 2NH3(g) H 0rxn = 92.6 kJ/mol Page 19 91. If 2.0 moles of N2 react with 6.0 moles of H2 to form NH3, calculate the work done (in kJ) against a pressure of 1.0 atm at 25°C. Assume the reaction goes to completion. A) 9.9 kJ B) 2.5 kJ C) 2.5 103 kJ D) 9.9 103 kJ 92. If 2.0 moles of N2 react with 6.0 moles of H2 to form NH3, at 1.0 atm at 25°C, then what is E for this reaction? Assume the reaction goes to completion. A) –92.6 kJ B) –175.3 kJ C) –185.3 kJ D) –195.1 kJ 93. Calculate the heat released when 2.00 L of Cl2(g) with a density of 1.88 g/L reacts with an excess of sodium metal at 25°C and 1 atm to form sodium chloride. A) 422 kJ B) 1550 kJ C) 822 kJ D) 46.4 kJ Use the following to answer questions 94-95: Photosynthesis produces glucose, C6H12O6, and oxygen from carbon dioxide and water: 6CO2 + 6H2O C6H12O6 + 6O2 94. H 0rxn How would you determine experimentally the value for this reaction? A) Form glucose at constant pressure. B) Form glucose at constant volume. C) Combust glucose in a bomb calorimeter. D) Put CO2 and H2 in a constant pressure calorimeter and allow to react. 95. Solar radiation produces about 7.01014 kg glucose a year on Earth. What is the corresponding H0 change? A) 1.1 1019 kJ B) 1.1 1013 kJ C) –1.3 1013 kJ D) –1.3 1019 kJ Page 20 96. A 2.10-mole sample of crystalline acetic acid, initially at 17.0°C, is allowed to melt at 17.0°C and is then heated to 118.1°C (its normal boiling point) at 1.00 atm. The sample is allowed to vaporize at 118.1°C and is then rapidly quenched to 17.0°C, so that it recrystallizes. Is the H0 for the total process as described: positive, negative, or zero? A) Positive B) Negative C) Zero D) Cannot be determined 97. Calculate the work done in joules by the reaction 2Na(s) + 2H2O(l) 2NaOH(aq) + H2(g) when 0.34 g of Na reacts with water to form hydrogen gas at 0°C and 1.0 atm. A) –17 J B) –34 J C) 17 J D) 34 J 98. You are given the following data: H2(g) 2H(g) H0 = 436.4 kJ/mol Br2(g) 2Br(g) H0 = 192.5 kJ/mol H2(g) + Br2(g) 2HBr(g) H0 = -72.4 kJ/mol Calculate H0 for the reaction H(g) + Br(g) HBr(g) A) –158.2 kJ B) –350.7 kJ C) –386.9 kJ D) –568.9 kJ 99. Methanol (CH3OH) is an organic solvent and is also used as a fuel in some automobile engines. From the following data, calculate the standard enthalpy of formation of methanol: H 0rxn 2CH3OH(l) + 3O2(g) 2CO2(g) + 4H2O(l) = 1452.8 kJ/mol A) –773.5 kJ/mol B) –477.4 kJ/mol C) –386.8 kJ/mol D) –238.7 kJ/mol Page 21 100. A 44.0-g sample of an unknown metal at 99.0°C was placed in a constant-pressure calorimeter containing 80.0 g of water at 24.0°C. The final temperature of the system was found to be 28.4°C. Calculate the specific heat of the metal. (The heat capacity of the calorimeter is 12.4 J/°C.) A) 2.03 J/g °C B) 1.01 J/g °C C) 0.492 J/g °C D) 0.246 J/g °C 101. Producer gas (carbon monoxide) is prepared by passing air over red-hot coke: 1 C(s) + 2 O2(g) CO(g) Water gas (mixture of carbon monoxide and hydrogen) is prepared by passing steam over red-hot coke: C(s) + H2O(g) CO(g) + H2(g) For many years, both producer gas and water gas were used as fuels in industry and for domestic cooking. The large-scale preparation of these gases was carried out by alternating between forming water gas and forming producer gas. Compare the standard enthalpies for these reactions. Which process would be used first when starting up the production of these gases? A) Forming water gas would be first. B) Forming producer gas would be first. C) It wouldn't matter which was first. 102. Compare the heat produced by the complete combustion of a mole of methane (CH4) with that of a mole of water gas (0.50 mole H2 and 0.50 mole CO) under the same conditions. On the basis of your answer, which would you prefer as a fuel: methane or water gas? A) Methane B) Water gas C) No preference Page 22 103. The so-called hydrogen economy is based on hydrogen produced from water using solar energy. The gas is then burned as a fuel: 2H2(g) + O2(g) 2H2O(l) A primary advantage of hydrogen as a fuel is that it is nonpolluting. A major disadvantage is that it is a gas and therefore is harder to store than liquids or solids. Calculate the volume of hydrogen gas at 25°C and 1.00 atm required to produce an amount of energy equivalent to that produced by the combustion of a gallon of octane (C8H18). The density of octane is 2.66 kg/gal, and its standard enthalpy of formation is 249.9 kJ/mol. A) 24.5 L B) 9.15 102 L C) 5.45 103 L D) 1.09 104 L 104. Ethanol (C2H5OH) and gasoline (assumed to be all octane, C8H18) are both used as automobile fuel. If gasoline is selling for $1.20/gal, what would the price of ethanol H 0f have to be in order to provide the same amount of heat per dollar? The density and of octane are 0.7025 g/mL and 249.9 kJ/mol and of ethanol are 0.7894 g/mL and 277.0 kJ/mol, respectively. 1 gal = 3.785 L. A) $ 0.25/gal B) $ 0.72/gal C) $ 0.84/gal D) $ 1.04/gal 105. The combustion of what volume of ethane (C2H6), measured at 23.0°C and 752 mmHg, would be required to heat 855 g of water from 25.0°C to 98.0°C? A) 0.319 L B) 4.02 L C) 4.10 L D) 147 L 106. The heat of vaporization of a liquid (Hvap) is the energy required to vaporize 1.00 g of the liquid at its boiling point. In one experiment, 60.0 g of liquid nitrogen (boiling point –196°C) are poured into a Styrofoam cup containing 2.00 102 g water at 55.3°C. Calculate the molar heat of vaporization of liquid nitrogen if the final temperature of the water is 41.0°C. A) 12.0 kJ/mol B) 5.60 kJ/mol C) –5.60 kJ/mol D) –12.0 kJ/mol Page 23 107. Explain the cooling effect experienced when ethanol is rubbed on your skin, given that C2H5OH(l) C2H5OH(g) H0 = 42.2 kJ/mol A) The reaction is exothermic, absorbing heat from your skin. B) The reaction is endothermic, absorbing heat from your skin. C) The ethanol abosorbs into your skin, cooling it as it enters. 108. For which of the following reactions does (a) H2(g) + S(rhombic) H2S(g) (b) C(diamond) + O2(g) CO2(g) (c) H2(g) + CuO(s) H2O(l) + Cu(s) (d) O(g) + O2(g) O3(g) A) Reaction (a) B) Reaction (b) C) Reaction (d) D) Reactions (a) and (b) E) Reactions (a), (b), and (d) H 0rxn H 0f ? 109. Calculate the work done (in joules) when 1.0 mole of water is frozen at 0°C and 1.0 atm. The volumes of one mole of water and ice at 0°C are 0.0180 L and 0.0196 L, respectively. A) –0.16 J B) –1.6 10–3 J C) 0.16 J D) 1.6 10–3 J 110. A quantity of 0.020 mole of a gas initially at 0.050 L and 20°C undergoes a constanttemperature expansion until its volume is 0.50 L. If the gas expands against a vacuum, is work done by the gas on the surroundings, or by the surroundings on the gas? A) Work is done by the gas on the surroundings. B) Work is done by the surroundings on the gas. C) No work is done. 111. A quantity of 0.020 mole of a gas initially at 0.050 L and 20°C undergoes a constanttemperature expansion until its volume is 0.50 L. Calculate the work done (in joules) by the gas if it expands against a constant pressure of 0.20 atm. A) 0.09 J B) –0.09 J C) 9.1 J D) –9.1 J Page 24 112. A quantity of 0.020 mole of a gas initially at 0.050 L and 20°C undergoes a constanttemperature expansion. If it expands against a constant pressure of 0.20 atm, and if it is allowed to expand unchecked until its pressure is equal to the external pressure; what would its final volume be before it stopped expanding? A) 2.4 L B) 0.48 L C) 0.16 L D) 0.096 L 113. A quantity of 0.020 mole of a gas initially at 0.050 L and 20°C undergoes a constanttemperature expansion. If it expands against a constant pressure of 0.20 atm, and if it is allowed to expand unchecked until its pressure is equal to the external pressure; what would be the work done by the gas? A) –87 J B) –48 J C) –22 J D) –9.3 J 114. Calculate the standard enthalpy of formation for diamond, given that C(graphite) + O2(g) CO2(g) H0 = 393.5 kJ/mol C(diamond) + O2(g) CO2(g) H0 = 395.4 kJ/mol A) –788.9 kJ B) –1.9 kJ C) 1.9 kJ D) 788.9 kJ 115. Fermentation is a complex chemical process of wine making in which glucose is converted into ethanol and carbon dioxide: C6H12O6 2C2H5OH + 2CO2 Calculate the standard enthalpy change for the fermentation process. A) –604.0 kJ B) –66.5 kJ C) 66.5 kJ D) 604.0 kJ Page 25 116. Portable hot packs are available for skiers and people engaged in other outdoor activities in a cold climate. The air-permeable paper packet contains a mixture of powdered iron, sodium chloride, and other components, all moistened by a little water. The exothermic reaction that produces the heat is a very common one — the rusting of iron: 4Fe(s) + 3O2(g) 2Fe2O3(s) When the outside plastic envelope is removed, O2 molecules penetrate the paper, causing the reaction to begin. A typical packet contains 250 g of iron to warm your hands or feet for up to 4 hours. How much heat (in kJ) is produced by this reaction? 0 (Hint: See values for H f .) A) B) C) D) 3.68 103 kJ –1.84 103 kJ 1.84 103 kJ 3.68 103 kJ 117. A person ate 0.50 pound of cheese (an energy intake of 4000 kJ). Suppose that none of the energy was stored in his body. What mass (in grams) of water would he need to perspire in order to maintain his original temperature? (It takes 44.0 kJ to vaporize 1 mole of water.) A) 5.05 g B) 90.9 g C) 9.09 102 g D) 1.64 103 g 118. The total volume of the Pacific Ocean is estimated to be 7.2 108 km3. A medium-sized atomic bomb produces 1.0 1015 J of energy upon explosion. Calculate the number of atomic bombs needed to release enough energy to raise the temperature of the water in the Pacific Ocean by 1°C. A) 3.0 103 atomic bombs B) 3.0 106 atomic bombs C) 3.0 109 atomic bombs D) 3.0 1012 atomic bombs 119. A 19.2-g quantity of dry ice (solid carbon dioxide) is allowed to sublime (evaporate) in an apparatus like the one shown in Figure 6.10. Calculate the work done by the gas against a constant external pressure of 0.995 atm and at a constant temperature of 22°C. Assume that the initial volume of dry ice is negligible and that CO2 behaves like an ideal gas. A) –0.0104 kJ B) –0.079 kJ C) –1.06 kJ D) –10.5 kJ Page 26 120. The enthalpy of combustion of benzoic acid (C6H5COOH) is commonly used as the standard for calibrating constant-volume bomb calorimeters; its value has been accurately determined to be –3226.7 kJ/mol. When 1.9862 g of benzoic acid are burned, the temperature rises from 21.84°C to 25.67°C. What is the heat capacity of the calorimeter? (Assume that the quantity of water surrounding the calorimeter is exactly 2000 g.) A) 5.35 kJ/°C B) 4.56 kJ/°C C) –4.56 kJ/°C D) –5.35 kJ/°C Use the following to answer questions 121-122: Lime is a term that includes calcium oxide (CaO, also called quicklime) and calcium hydroxide [Ca(OH)2, also called slaked lime]. It is used in the steel industry to remove acidic impurities, in air-pollution control to remove acidic oxides such as SO2, and in water treatment. Quicklime is made industrially by heating limestone (CaCO3) above 2000°C: CaCO3(s) CaO(s) + CO2(g) H0 = 177.8 kJ/mol Slaked lime is produced by treating quicklime with water: CaO(s) + H2O(l) Ca(OH)2(s) H0 = 65.2 kJ/mol The exothermic reaction of quicklime with water and the rather small specific heats of both quicklime (0.946 J/g °C) and slaked lime (1.20 J/g °C) make it hazardous to store and transport lime in vessels made of wood. Wooden sailing ships carrying lime would occasionally catch fire when water leaked into the hold. 121. If a 500-g sample of water reacts with an equimolar amount of CaO (both at an initial temperature of 25°C), what is the final temperature of the product, Ca(OH)2? Assume that the product absorbs all of the heat released in the reaction. A) 462 °C B) 711 °C C) 761 °C D) 1497 °C 122. Given that the standard enthalpies of formation of CaO and H2O are 635.6 kJ/mol and 285.8 kJ/mol, respectively, calculate the standard enthalpy of formation of Ca(OH)2. A) 986.6 kJ/mol B) 921.4 kJ/mol C) –921.4 kJ/mol D) –986.6 kJ/mol Page 27 123. Calcium oxide (CaO) is used to remove sulfur dioxide generated by coal-burning power stations: 2CaO(s) + 2SO2(g) + O2(g) 2CaSO4(s) Calculate the enthalpy change for this process if 6.6 105 g of SO2 are removed by this process every day. A) –1.0 107 kJ B) –5.2 106 kJ C) –5.6 106 kJ D) –1.0 103 kJ 124. Glauber's salt, sodium sulfate decahydrate (Na2SO4 10H2O), undergoes a phase transition (that is, melting or freezing) at a convenient temperature of about 32°C: Na2SO4 10H2O(s) Na2SO4 +10H2O(l) H0 = 74.4 kJ/mol As a result, this compound is used to regulate the temperature in homes. It is placed in plastic bags in the ceiling of a room. During the day, the endothermic melting process absorbs heat from the surroundings, cooling the room. At night, it gives off heat as it freezes. Calculate the mass of Glauber's salt in kilograms needed to lower the temperature of air in a room by 8.2°C at 1.0 atm. The dimensions of the room are 2.80 m 10.6 m 17.2 m, the specific heat of air is 1.2 J/g °C, and the molar mass of air may be taken as 29.0 g/mol. A) 0.25 kg B) 25 kg C) 2.4 kg D) 240 kg 125. A balloon 16 m in diameter is inflated with helium at 18°C. Calculate the mass of He in the balloon, assuming ideal behavior and a pressure of 98.7 kPa. A) 2.1 104 g He B) 4.3 104 g He C) 3.4 105 g He D) 5.6 106 g He 126. A balloon 16 m in diameter is inflated with helium at 18°C. Assuming ideal behavior, calculate the work done (in joules) during the inflation process if the atmospheric pressure is 98.7 kPa. A) –2.0 105 J B) –2.1 108 J C) –2.1 1010 J D) 2.0 105 J Page 28 127. An excess of zinc metal is added to 50.0 mL of a 0.100 M AgNO3 solution in a constantpressure calorimeter like the one pictured in Figure 6.7. As a result of the reaction Zn(s) + 2Ag+(aq) Zn2+(aq) + 2Ag(s) the temperature rises from 19.25°C to 22.17°C. If the heat capacity of the calorimeter is 98.6 J/°C, calculate the enthalpy change for the above reaction on a molar basis. Assume that the density and specific heat of the solution are the same as those for water, and ignore the specific heats of the metals. A) 360 kJ/mol Zn (180 kJ/mol Ag+) B) 130 kJ/mol Zn (65 kJ/mol Ag+) C) –360 kJ/mol Zn (–180 kJ/mol Ag+) D) –130 kJ/mol Zn (–65 kJ/mol Ag+) 128. A person drinks four glasses of cold water (3.0°C) every day. The volume of each glass is 2.5 102 mL. How much heat (in kJ) does the body have to supply to raise the temperature of the water to 37°C, the body temperature? A) 1.4 105 kJ B) 1.4 102 kJ C) 1.4 103 kJ D) –1.4 105 kJ 129. How much heat (in kJ) would your body lose if you were to ingest 8.0 102 g of snow at 0°C to quench thirst? (Normal body temperature is 37°C. The amount of heat necessary to melt snow is 6.01 kJ/mol.) A) 3.9 102 kJ B) 4.9 102 kJ C) 4.9 103 kJ D) 1.2 105 kJ Use the following to answer questions 130-131: At 25°C the standard enthalpy of formation of HF(aq) is 320.1 kJ/mol; of OH-(aq), it is 229.6 kJ/mol; of F- (aq), it is 329.1 kJ/mol; and of H2O(l), it is 285.8 kJ/mol. 130. Calculate the standard enthalpy of neutralization of HF(aq): HF(aq) + OH(aq) F(aq) + H2O(l) A) –65.2 kJ B) –506.4 kJ C) –524.4 kJ D) –1164.6 kJ Page 29 131. Use the value of 56.2 kJ as the standard enthalpy change for the reaction H+(aq) + OH(aq) H2O(l) Calculate the standard enthalpy change for the reaction HF(aq) H+(aq) + F(aq) A) 121.4 kJ B) –9.0 kJ C) –121.4 kJ D) –468.2 kJ 132. From the standard enthalpy of formation for CO2 (-393.5 kJ/mol), and the following information, calculate the standard enthalpy of formation for carbon monoxide (CO). 1 CO(g) + 2 O2(g) CO2(g) H0 = 283.0 kJ/mol A) –676.5 kJ B) –110.5 kJ C) 110.5 kJ D) 676.5 kJ 133. A 46-kg person drinks 500 g of milk, which has a “caloric” value of approximately 3.0 kJ/g. If only 17 percent of the energy in milk is converted to mechanical work, how high (in meters) can the person climb based on this energy intake? [Hint: The work done in ascending is given by mgh, where m is the mass (in kilograms), g the gravitational acceleration (9.8 m/s2), and h the height (in meters).] A) 5.7 104 m B) 3.3 104 m C) 3.3 103 m D) 5.7 102 m Use the following to answer questions 134-136: The height of Niagara Falls on the American side is 51 meters. Consider 1.0 g of water at the top of the falls. 134. Calculate the potential energy of 1.0 g of water at the top of the falls relative to the ground level. A) 0.50 J B) 4.9 J C) 1.9 102 J D) 5.0 102 J Page 30 135. What is the speed of 1.0 g of falling water if all of the potential energy is converted to kinetic energy? A) 1000 m/s B) 320 m/s C) 32 m/s D) 10 m/s 136. What would be the increase in temperature of 1.0 g of water if all the kinetic energy were converted to heat? A) 0.12 ºC B) 1.2 ºC C) 45 ºC D) 120 ºC 137. Determine the standard enthalpy of formation of ethanol (C2H5OH) from its standard enthalpy of combustion (1367.4 kJ/mol). A) 688.1 kJ/mol B) –277.0 kJ/mol C) –2046.7 kJ/mol D) –3011.4 kJ/mol 138. Acetylene (C2H2) and benzene (C6H6) have the same empirical formula. In fact, benzene can be made from acetylene as follows: 3C2H2(g) C6H6(l) The enthalpies of combustion for C2H2 and C6H6 are 1299.4 kJ/mol and 3267.4 kJ/mol, respectively. Calculate the enthalpy change for the formation of C6H6 from C2H2. A) 2588.1 kJ B) 728.8 kJ C) –177.6 kJ D) –630.8 kJ 139. Ice at 0°C is placed in a Styrofoam cup containing 361 g of a soft drink at 23°C. The specific heat of the drink is about the same as that of water. Some ice remains after the ice and soft drink reach an equilibrium temperature of 0°C. Determine the mass of ice that has melted. Ignore the heat capacity of the cup. (Hint: It takes 334 J to melt 1 g of ice at 0°C.) A) 10.4 g B) 104 g C) 11.6 g D) 116 g Page 31 140. A gas company in Massachusetts charges $1.30 for 15 ft3 of natural gas (CH4) measured at 20°C and 1.0 atm. Calculate the cost of heating 200 mL of water (enough to make a cup of coffee or tea) from 20°C to 100°C. Assume that only 50 percent of the heat generated by the combustion is used to heat the water; the rest of the heat is lost to the surroundings. A) 0.76 cents, $0.0076 B) 1.1 cents, $0.011 C) 5.1 cents, $0.051 D) 7.0 cents, $0.070 Use the following to answer questions 141-142: A Goodyear blimp is filled with helium gas at 1.2 105 Pa. The volume of the blimp is 5.5 103 m3 . 141. Calculate the internal energy of the gas in the blimp. A) 37 J B) 3.7 103 J C) 9.9 108 J D) 9.9 1010 J 142. If all of the internal energy of the gas in the blimp were used to heat 10.0 tons of copper at 21°C, calculate the final temperature of the metal. The specific heat of copper is 0.385 J/g°C. (Hint: Calculate the internal energy of a gas. 1 ton = 9.072 105 g.) A) 305 °C B) 153 °C C) 1.0 °C D) 30.5 °C 143. Acetylene (C2H2) can be made by reacting calcium carbide (CaC2) with water. What is the maximum amount of heat (in kJ) that can be obtained from the combustion of acetylene, starting with 74.6 g of CaC2? A) 3.02 103 kJ B) 1.51 103 kJ C) 5.16 103 kJ D) 3.92 103 kJ Page 32 144. The average temperature in deserts is high during the day but quite cool at night, whereas that in regions along the coastline is more moderate. Which of the following give an explanation to this fact? A) The quantity of sand in the desert is much greater than that at the beach. B) The specific heat of water is much larger than that of sand; therefore the water will hold the thermal energy longer. C) The moon is able to transfer thermal energy to the coastlines at night. D) The coastline regions are closer to the equator than the deserts. 145. When 0.8436 mole of naphthalene (C10H8), a solid, is burned in a constant-volume bomb calorimeter at 298 K, 5150 kJ of heat is evolved. Calculate E for the reaction on a molar basis. A) 5154 kJ/mol B) 5145 kJ/mol C) –5145 kJ/mol D) –5153 kJ/mol Page 33 Answer Key 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. 31. 32. 33. 34. 35. 36. 37. 38. 39. 40. 41. 42. 43. 44. A C A B C B C A A D B A A B C D A B C C A B B A B A C D C D A D B C B D C B C C B C D A Page 34 45. 46. 47. 48. 49. 50. 51. 52. 53. 54. 55. 56. 57. 58. 59. 60. 61. 62. 63. 64. 65. 66. 67. 68. 69. 70. 71. 72. 73. 74. 75. 76. 77. 78. 79. 80. 81. 82. 83. 84. 85. 86. 87. 88. 89. 90. C A D C C A B D D B A C B A C A C D D B A A B C C B D A B B C C A D B B A C A D D A B D D C Page 35 91. 92. 93. 94. 95. 96. 97. 98. 99. 100. 101. 102. 103. 104. 105. 106. 107. 108. 109. 110. 111. 112. 113. 114. 115. 116. 117. 118. 119. 120. 121. 122. 123. 124. 125. 126. 127. 128. 129. 130. 131. 132. 133. 134. 135. 136. A B D C A C A B D C B A D C C B B A A C D A B C B B D C C A C D B B C B C B A A B B D A C A Page 36 137. 138. 139. 140. 141. 142. 143. 144. 145. B D B B C A B B D Page 37