PRACTICE MULTIPLE CHOICE- Chapters 17&18
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1. What is the solubility product expression for Fe2(CO3)3?
A) Ksp = [2Fe3+][3CO32–]
B) Ksp = [2Fe3+]2[3CO32–]3
C) Ksp = [Fe2+]2[CO32–]3
D) Ksp = [Fe3+]2[CO32–]3
E) Ksp = [2Fe3+]2[CO32–]3
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2. What is the solubility (in g/L) of calcium sulfate at 25°C? The solubility product constant for
calcium sulfate is 2.4  10–5 at 25°C.
A) 0.67 g/L
B) 1.6  10–3 g/L
C) 2.5 g/L
D) 3.3  10–3 g/L
E) 4.8 g/L
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3. The solubility of La(IO3)3 in a 0.89 M KIO3 solution is 1.0 10–7 mol/L. Calculate Ksp for La(IO3)3.
A) 7.0  10–1
B) 7.0  10–22
C) 8.9  10–8
D) 7.0  10–8
E) none of these
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4. A 6.0  10–4 M solution of MnSO4 is gradually made more basic by adding NaOH. At what pH
will manganese(II) hydroxide begin to precipitate? For Mn(OH)2, Ksp = 2.0  10–13.
A) 4.74
B) 9.48
C) 9.26
D) 4.52
E) 9.57
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5. If 200 mL of 1  10–7 M AgNO3 is mixed with 200 mL of 1  10–8 M NaI, what will occur? For
AgI, Ksp = 8.3  10–17.
A) No precipitate will form.
B) Silver(I) nitrate will precipitate.
C) Silver(I) iodide will precipitate.
D) Sodium nitrate will precipitate.
E) Sodium iodide will precipitate.
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6. The best explanation for the dissolution of ZnS in dilute HCl is that
A) the zinc ion is amphoteric.
B) the sulfide ion concentration is decreased by the formation of H2S.
C) the solubility product of ZnCl2 is less than that of ZnS.
D) the zinc ion concentration is decreased by the formation of a chloro complex.
E) the sulfide ion concentration is decreased by oxidation to sulfur.
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7. What is the best way to ensure complete precipitation of SnS from a saturated H2S solution?
A)
B)
C)
D)
E)
Add a strong acid.
Add a weak acid.
Add a strong base.
Add a weak base.
Add more H2S.
____
8. Cyanide ion forms very stable complex ions with a variety of metal ions. What is the molar
equilibrium concentration of uncomplexed Ag+(aq) in a solution that initially contains 1.3 mol of
Ag(CN)2 per liter of solution . Kf for Ag(CN)2 is 4.5  10 10 .
A) 5.2  10 3 M
B) 7.3  10 12 M
C) 1.9  10 4 M
D) 2.7  10 6 M
E) 3.1  10 4 M
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9. Given the following equilibrium constants,
Zn(IO3)2 Ksp = 3.9  10 6
Zn(NH3)42+ Kf = 2.9  10 9
determine Kc for the dissolution of the sparingly soluble salt Zn(IO3)2 in aqueous ammonia
(shown below).
Zn(IO3)2(s) + 4NH3(aq)
Zn(NH3)42+(aq) + 2IO3–(aq)
A) 8.8  10 5
B) 1.1  10 4
C) 7.4  10 14
D) 2.9  10 9
E) 1.3  10 15
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10. Assuming H and S are constant with respect to temperature, under what conditions will a
chemical reaction be spontaneous at all temperatures?
A) H = 0, and S is negative.
B) S = 0, and H is positive.
C) H is positive, and S is negative.
D) H is negative, and S is positive.
E) none of these
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11. For which of the following reactions is S° > 0 at 25°C?
A) 2H2(g) + O2(g)  2H2O(g)
B) 2ClBr(g)  Cl2(g) + Br2(g)
C) I2(g)  I2(s)
D) 2NO(g) + O2(g)  2NO2(g)
E) NH4HS(s)  NH3(g) + H2S(g)
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12. What is S° at 298 K for the following reaction?
CH4(g) + N2(g)  HCN(g) + NH3(g); H° = 164.1 kJ; G° = 159.1 kJ at 298 K
A)
B)
C)
D)
E)
____
2.0 J/K
5.5  102 J/K
1.1  103 J/K
5.3  102 J/K
17 J/K
13. Given the following, determine G°f at 298 K for PbO.
Pb(s) + PbO2(s)  2PbO(s) ; G° = –158.5 kJ at 298K
Substance
PbO(s)
PbO2(s)
A)
B)
C)
D)
E)
G°f (kJ/mol) at 298 K
?
–217.3
–187.9 kJ/mol
29.4 kJ/mol
–375.8 kJ/mol
117.6 kJ/mol
58.8 kJ/mol
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14. Which of the following is correct for the condensation of gaseous oxygen at –188°C? The normal
boiling point of oxygen is –183°C.
A) H > 0, S > 0, and G > 0.
B) H < 0, S < 0, and G < 0.
C) H < 0, S > 0, and G > 0.
D) H > 0, S < 0, and G < 0.
E) H = 0, S = 0, and G < 0.
____
15. For a reaction that has an equilibrium constant of 3  10–6, which of the following statements must
be true?
A) G° is positive.
B) S° is positive.
C) H° is negative.
D) G° is negative.
E) H° is positive.
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16. The standard free energy of formation of nitric oxide, NO, at 1000. K (roughly the temperature in
an automobile engine during ignition) is 78.4 kJ/mol. Calculate the equilibrium constant for the
reaction
2NO(g)
at 1000. K. (R = 8.31 J/(Kmol))
A) 8.0  10–5
B) 6.4  10–9
C) –15
D) 1.6  105
E) 0.95
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17. Consider the following reaction:
2C(s) + 3H2(g)  C2H6(g); H° = –84.68 kJ; S° = –173.8 J/K at 298 K
What is the equilibrium constant at 298 K for this reaction?
A) 5.8  105
B) 1.7  10–6
C) 8.6  10–10
D) 7.0  1014
E) 1.0
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18. What is G° at 500.0 K for the following reaction?
Mg(s) + H2O(g)  MgO(s) + H2(g)
Substance
Mg(s)
H2O(g)
MgO(s)
H2(g)
A)
B)
C)
D)
E)
H°f (kJ/mol) at 298 K
0
–241.8
–601.6
0
S° (J/(molK)) at 298 K
32.7
188.7
27.0
130.6
340.8 kJ
–340.8 kJ
–327.9 kJ
–391.7 kJ
327.9 kJ
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19. For the reaction SrCO3(s)  SrO(s) + O2(g) at 1 atm pressure, the values of H and S are both
positive, and the process is spontaneous at high temperatures. Which of the following statements
about this reaction is true?
A) The reverse reaction is endothermic.
B) The change in entropy is the driving force for the reaction.
C) G at room temperature is negative.
D) The process is exothermic at high temperatures and endothermic at room
temperature.
E) The reverse reaction is nonspontaneous at room temperature.
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20. Condensation is a process for which
A) G is negative at high temperature but positive at low temperature.
B) H, S, and G are positive at all temperatures.
C) G is positive when condensation occurs spontaneously.
D) H and S are positive at all temperatures.
E) H and S are negative at all temperatures.
PRACTICE MULTIPLE CHOICE- Chapters 17&18
Answer Section
1. ANS: D
PTS: 1
DIF: easy
REF: 17.1
OBJ: Write solubility product expressions. (Example 17.1)
TOP: solubility | solubility equilibria
KEY: solubility product constant
MSC: general chemistry
2. ANS: A
PTS: 1
DIF: easy
REF: 17.1
OBJ: Calculate the solubility from Ksp. (Example 17.4)
TOP: solubility | solubility equilibria
KEY: solubility product constant
MSC: general chemistry
3. ANS: D
PTS: 1
DIF: moderate
REF: 17.2
OBJ: Calculate the solubility of a slightly soluble salt in a solution of a common ion. (Example
17.5) TOP:
solubility | solubility equilibria
KEY: solubility and the common-ion effect
MSC: general
chemistry
4. ANS: C
PTS: 1
DIF: easy
REF: 17.3
OBJ: Predict whether precipitation will occur (given ion concentrations). (Example 17.6)
TOP: solubility | solubility equilibria
KEY: precipitation calculations | criterion for precipitation
MSC: general chemistry
5. ANS: C
PTS: 1
DIF: easy
REF: 17.3
OBJ: Predict whether precipitation will occur (given solution volumes and concentrations).
(Example 17.7)
TOP: solubility | solubility equilibria
KEY: precipitation calculations | criterion for precipitation
MSC: general chemistry
6. ANS: B
PTS: 1
DIF: easy
REF: 17.4
OBJ: Explain the qualitative effect of pH on solubility of a slightly soluble salt.
TOP: solubility | solubility equilibria
KEY: effect of pH on solubility | qualitative effect of pH
MSC: general chemistry
7. ANS: C
PTS: 1
DIF: moderate
REF: 17.4
OBJ: Explain the basis for the sulfide scheme to separate a mixture of metal ions.
TOP: solubility | applications of solubility equilibria
MSC: general chemistry
8. ANS: C
PTS: 1
DIF: easy
REF: 17.5
OBJ: Calculate the concentration of a metal ion in equilibrium with a complex ion. (Example
17.9) TOP:
solubility | complex ion equilibria
9. ANS: B
PTS: 1
DIF: easy
REF: 17.6
OBJ: Calculate the solubility of a slightly soluble ionic compound in a solution of the complex
ion. (Example 17.11)
TOP: solubility | complex ion equilibria
10. ANS: D
PTS: 1
DIF: easy
REF: 18.2 | 18.7
OBJ: Describe how deltaH – TdeltaS functions as a criterion of a spontaneous reaction.
TOP: thermochemistry | thermodynamics
11. ANS: E
PTS: 1
DIF: easy
REF: 18.3
OBJ: Predict the sign of the entropy change of a reaction. (Example 18.2)
TOP: thermochemistry | thermodynamics
KEY: third law of thermodynamics | entropy change for a reaction
MSC: general chemistry
12. ANS: E
PTS: 1
DIF: easy
REF: 18.4
OBJ: Calculate deltaG from deltaH and deltaS. (Example 18.4)
TOP: thermochemistry | thermodynamics
KEY: free energy | standard free energy change
MSC: general chemistry
13. ANS: A
PTS: 1
DIF: easy
REF: 18.4
OBJ: Calculate deltaG from standard free energies of formation. (Example 18.5)
TOP: thermochemistry | thermodynamics
14. ANS: B
PTS: 1
DIF: moderate
REF: 18.4
OBJ: State the rules for using deltaG as a criterion for spontaneity.
TOP: thermochemistry | thermodynamics
KEY: free energy | spontaneity
MSC: general chemistry
15. ANS: A
PTS: 1
DIF: easy
REF: 18.6
OBJ: Relate the standard free-energy change to the thermodynamic equilibrium constant.
TOP: thermochemistry | thermodynamics
KEY: thermodynamic equilibrium constant (K)
MSC: general chemistry
16. ANS: B
PTS: 1
DIF: easy
REF: 18.6
OBJ: Calculate K from the standard free-energy change (molecular equation). (Example 18.8)
TOP: thermochemistry | thermodynamics
KEY: thermodynamic equilibrium constant (K)
MSC: general chemistry
17. ANS: A
PTS: 1
DIF: moderate
REF: 18.6
OBJ: Calculate K from the standard free-energy change (molecular equation). (Example 18.8)
TOP: thermochemistry | thermodynamics
KEY: thermodynamic equilibrium constant (K)
MSC: general chemistry
18. ANS: C
PTS: 1
DIF: moderate
REF: 18.7
OBJ: Describe how deltaG at a given temperature (deltaG T) is approximately related to deltaH
and deltaS at that temperature.
TOP: thermochemistry | thermodynamics
KEY: temperature dependence of free energy | calculation of free energy change at various
temperatures
MSC: general chemistry
19. ANS: B
PTS: 1
DIF: easy
REF: 18.7
OBJ: Describe the how spontaneity or nonspontaneity of a reaction is related to each of the four
possible choices of signs of deltaH and deltaS.
TOP: thermochemistry | thermodynamics
KEY: temperature dependence of free energy | spontaneity and temperature change
MSC: general chemistry
20. ANS: E
PTS: 1
DIF: moderate
REF: 18.7
OBJ: Describe the how spontaneity or nonspontaneity of a reaction is related to each of the four
possible choices of signs of deltaH and deltaS.
TOP: thermochemistry | thermodynamics
KEY: temperature dependence of free energy | spontaneity and temperature change
MSC: general chemistry