Download 1. Define the following term: system. A) The part of the universe that

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.
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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.
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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.
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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
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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
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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
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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.
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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
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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.
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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/ gºC.
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/ gºC.
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/ gºC and 0.444 J/ gºC, 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 2.00102 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.
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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
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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
AC
= H2
Determine the enthalpy change for the process
2B  C
A) H2  H1
B) H2 + H1
C) H1  H2
D) 2H1  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 7.01014 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
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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
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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
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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
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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
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137.
138.
139.
140.
141.
142.
143.
144.
145.
B
D
B
B
C
A
B
B
D
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