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CHEM 102: Sample Test 5
CHAPTER 17
1. When H2SO4 is dissolved in water, which species would be found in the water at equilibrium in
measurable amounts?
a. H2SO4 b. H3SO4 + c. HSO4 - d. SO4 -2 e. OH2. Which of the following is the net ionic equation for the reaction that occurs during the titration of
nitrous acid with potassium hydroxide?
a) HNO2 + K+ + OH- Æ KNO2 + H2O
b) HNO2 + H2O Æ NO2- + H3O+
c) HNO2 + KOH Æ K+ + NO2- + H2O
d) HNO2 + OH- Æ NO2- + H2O
e) H+ + OH- Æ H2O
3. What is the pH of a solution that is 0.2 M in acetic acid (Ka = 1.8 x 10-5) and 0.2 M in sodium
acetate?
a) 4.7
b) 9.3 c) 7.0 d) 5.4 e) 8.6
4. Given 0.015 moles of NaC2H3O2and 0.05 moles of HC2H3O2 dissolved in water, what is the pH? pKa
for acetic acid is 4.74.
a. 5.22 b. 4.74 c. 4.22 d. 1.30 e. Cannot determine
5. If 50 ml of a 0.01 M HCl solution is titrated with a 0.01 M NaOH solution, what will be the initial
pH and the pH at the endpoint.
a. 2 and 3.5
b. 2 and 7 c. 3.30 and 7 d. 12 and 7 e. None of these
6. What is the pH of a 1.0 M aqueous solution of NaCl?
a) 7.0 b) greater than 7.0
c) less than 7.0
d) there is not enough information given
7. A 50.00-mL sample of 0.100 M KOH is being titrated with 0.100 M HNO3. Calculate the pH of the
solution after 52.00 mL of HNO3 is added.
a) 6.50 b) 3.01 c) 2.71 d) 2.41 e) none of these
8. If the following substance is dissolved in pure water, will the solution be acidic, neutral, or basic?
solid sodium carbonate: (Na2CO3).
a) acidic
b) neutral
c) basic
9. The Henderson-Hasselbach equation can be used to determine several factors. Which of the
following can be determined?
a. the pH for a solution of a strong acid and its conjugate base
b. the pH for a strong base that has been partially titrated
c. the mole ratio needed to make a buffer solution having a known pH
d. the pKa of an equimolar mixture of an unknown weak acid and its conjugate base without knowing
the pH
e. the change in pH if some acid or base is added to a buffered solution
10. A 1 Liter buffered solution was prepared by dissolving 1.0 mole of NH3 and 1.5 mole of NH4Cl in
water. The pH of this solution would be: (Kb = 1.8 x 10-5)
a. 11.0 b. 9.0
c. 7.0
d. 4.0
e. none of the above
11. The salt that forms an alkaline (basic) aqueous solution is
a. KNO3.
b. NH4Cl.
c. NaC2H3O2.
d. NaCl.
12. A weak acid that has different colors for acid and its conjugate base can be used as
a. an salt.
b. an indicator.
c. a buffer.
d. a base.
13. Give the pH value of a solution 0.10 M in acetic acid (Ka = 1.8 x 10-5) and 0.05 M in sodium
acetate:
a. 1.8
b. 2.4
c. 4.7
d. 4.4
e. 5
14. Methyl violet is an indicator used in acid-base titrations. Its acid form, HIn, is red, while its
conjugate base from, In-, is yellow. The color change occurs in the pH range 0.00-1.6. In a solution of
pH = 6.00 the color of the indicator will be
a. red
b. yellow
c. orange
d. colorless
e. blue
15. In the titration of a weak acid HA with 0.100 M NaOH, the stoichiometric point is known to occur
at a pH value of approximately 11. Which of the following indicators would be best to use to mark
the endpoint of this titration?
a) an indicator with Ka = 10-10
b) an indicator with Ka = 10-8
c) an indicator with Ka = 10-14
d) an indicator with Ka = 10-11
e) an indicator with Ka = 10-12
16. Calculate the pH of an aqueous solution of 2.0 M NH4Cl (Kb for NH3 =1.8 x 10-5). Choose your
answer from the following pH ranges:
a. pH = 0.00-2.99 b. pH = 3.00-5.99 c. pH = 6.00-8.99
d. pH = 9.00-10.99 e. pH = 11.00-14.00
17. Assume that an indicator works BEST when the equivalent point of a titration comes in the
middle of the indicator range. Which of the following indicators would be the best for a titration of
35.00 mL of 0.10 M HC2H3O2 (Ka=1.8 x 10-5) with 35.00 mL 0.10 M NaOH?
a. methyl violet 0.0-1.6 b. methyl orange 3.2-4.4
c. bromocresol green 3.8-5.4 d. methyl red 4.8-6.0
e. phenolphthalein 8.2-10.0
18. Given a equimolar mixture of Benzoic acid and sodium benzoate what would be the pH of the
solution? Ka of HC7H5O2 = 6.28 x 10-5
a. 4.20
b. 9.80 c. 2.10 d. 6.28
e. Cannot determine
19. A weak monoprotic acid (HA) is 1.41% dissociated in a 1.0 M solution. The Ka for this acid is
a. 4.0 x 10-6. b. 5.6 x 10-5. c. 2.0 x 10-4. d. 9.9 x 10-3.
20. Buffer solution is one which
a. contains the maximum amount of solute possible for a particular temperature.
b. contains more than the expected amount of solute for a particular temperature and is therefore unstable.
c. resists changes in pH upon addition of acid or base.
d. contains an equal number of hydronium and hydroxide ions.
e. changes color upon addition of strong base.
21. Which pair of compounds could be used to make a buffer solution?
a. NaCl/HCl
b. (NH4)2SO4/H2SO4
c. NaCH3COO/CH3COOH
d. K3PO4/KH2PO4
e. KOH/KBr
22. A salt which can be used to make a buffer with formic acid, HCOOH, is
a. NH4COOH
b. KCOOH
c. NaCl
d. Na2CO3
e. (NH4)3PO4
23. Which two of the following compounds in aqueous solution would make an effective buffer solution?
KNO3
CH3COOH
NH4NO3
NH3
HNO3
a. HNO3 & NH4NO3
b. HNO3 & KNO3
c. CH3COOH & NH3
d. NH3 & NH4NO3
e. CH3COOH & HNO3
24. Which combination of solutions is the best choice for making a buffer solution?
a. equal volumes of 0.1 M formic acid and 0.1 M sodium formate
b. equal volumes of 0.05 M hydrochloric acid and 0.075 ammonium chloride
c. equal volumes of 0.1 M sulfuric acid and 0.001 M sodium sulfate
d. equal volumes of 1 M acetic acid and 0.005 M sodium acetate
e. equal volumes of 0.5 M nitric acid and 0.5 M sodium hydroxide
25. Calculate the pH of a solution which is 0.05 M in lactic acid and 0.04 M in sodium lactate. The Ka for
lactic acid is 1.8 × 10-4.
a. 10.35
b. 3.74
c. 3.65
d. 0.097
e. 2.25 × 10-4
26. Calculate the pH of a solution that contains 0.45 M benzoic acid and 0.40 M sodium benzoate. The Ka
for benzoic acid is 6.5 × 10-5.
a. 7.3 × 10-5
b. 4.53
c. 4.19
d. 4.14
e. 3.79
27.Which formula represents the Henderson-Hasselbach equation for the generic acid HA?
[A - ]
[HA]
a.
pK a = pH + log
b.
pH = pK a − log
c.
log
d.
pH = pK a +
e.
[A - ]
pH = pK a + log
[HA]
[A - ]
[HA]
[A - ]
= pK a − pH
[HA ]
[A - ]
[HA]
28. Which pair of reagents is the best choice to make a buffer of pH = 4.50?
a. HF/FKa = 7.2 × 10-4
Ka = 1.8 × 10-5
b. CH3COOH/CH3COO
2Ka = 6.2 × 10-8
c. H2PO4 /HPO4
+
Ka = 5.6 × 10-10
d. NH4 /NH3
23Ka = 3.6 × 10-13
e. HPO4 /PO4
29.Consider a buffer solution made up of H2PO4- and HPO42-, which has a Ka of 6.2 × 10-8. What ratio of
HPO42- to H2PO4- will give a pH of 7.35?
a. 1.38 to 1
b. 0.140 to 1
c. 1 to 1
d. 0.725 to 1
e. More information is needed to answer this question.
30. One liter of a buffer is prepared using equimolar amounts of ascorbic acid and sodium ascorbate,
producing a solution with pH = 4.10 After addition of 10 mL of 1 M NaOH, the most likely value of
the pH is
a. 2.00
b. 4.10
c. 4.12
d. 5.90
e. 8.00
31. Calculate the pH of a buffer that is 0.14 M NH3 and 0.10 M NH4NO3. The Kb for ammonia is 1.8 ×
10-5.
a. 4.74
b. 4.89
c. 9.11
d. 9.25
e. 9.40
32. Which is the smallest amount of solid NaOH shown that will exceed the buffer capacity of a 500. mL
solution that is 0.40 M in acetic acid and 0.15 M in sodium acetate. The Ka for acetic acid is 1.8 × 10-5.
a. 3.00 g
b. 4.31 g
c. 4.74 g
d. 5.16 g
e. 6.00 g
33. Which statement about the titration of 0.10 M HNO3 with 0.10 M KOH is not correct?
a. The pH at the equivalence point is 7.00.
b. The initial pH is 1.00.
c. At the equivalence point the pH decreases sharply.
d. At the equivalence point the volume of base added will be equal to the original volume of acid.
e. The net ionic equation is H3O+ + OH- → 2H2O.
34. Calculate the pH of a titration mixture when 25.00 mL of 0.106 M NaOH has been added to a 50.00
mL sample of 0.0950 M HNO3.
a. 0.97
b. 1.02
c. 1.20
d. 1.45
e. 1.55
35. Calculate the volume of 0.106 M NaOH needed to neutralize a 50.00 mL sample of 0.0950 M HNO3.
a. 5.19 mL
b. 44.81 mL
c. 50.00 mL
d. 55.19 mL
e. 55.79 mL
36. Calculate the pH of a mixture of 51.0 mL of 0.106 M NaOH and 50.00 mL of 0.0950 M HNO3.
a. 11.81
b. 11.02
c. 9.31
d. 4.74
e. 2.19
37. Calculate the pH of a titration mixture when 35.00 mL of 0.106 M NaOH has been added to a 50.00
mL sample of 0.0950 M CH3COOH. Ka = 1.8 × 10-5.
a. 4.64
b. 4.74
c. 5.30
d. 8.70
e. 9.26
38. Calculate the pH of a titration mixture at the midpoint of the titration of a 50.00 mL sample of 0.0950
M CH3COOH with 0.106 M NaOH. Ka = 1.8 × 10-5.
a. 4.85
b. 4.74
c. 4.64
d. 1.02
e. 0.97
39. Calculate the pH of a titration mixture at the equivalence point of the titration of a 50.00 mL sample
of 0.0950 M CH3COOH with 0.106 M NaOH. Ka = 1.8 × 10-5.
a. 10.56
b. 9.53
c. 9.26
d. 8.72
e. 4.47
40. Calculate the pH of a titration mixture in which a 50.00 mL sample of 0.0950 M CH3COOH has
reacted with 50.00 mL of 0.106 M NaOH. Ka = 1.8 × 10-5.
a. 12.72
b. 11.74
c. 10.58
d. 9.26
e. 8.71
41. Which combination of acid and base would give the titration curve shown?
a. NaOH and HCl
b. KOH and H3PO4
c. NH4OH and CH3COOH
d. NaOH and H2SO4
e. NH4OH and H3PO4
42. Acid "rain" is defined as any precipitation with a pH value less than _____.
a. 1.0
b. 4.2
c. 5.6
d. 7.0
e. 8.5
43. Which list contains only compounds that directly contribute to acid rain?
a. CaO, H2Se, Kr
b. CO2, SiO2, GeO2
c. Fe2S3, CaCO3, Na2CO3
d. H2O, NH3, CH4
e. CO2, NO2, SO2
44. Write the Ksp expression for silver phosphate, Ag3PO4.
a. [Ag+]3[PO43-]
b.
45.
46.
47.
48.
49.
50.
51.
52.
[Ag + ]3 [PO 34- ]
[Ag 3 PO 4 ]
c. 3x[Ag+][PO43-]
d. 3x[Ag+]3 + [PO43-]
e. [H+][OH-]
Calculate the molar solubility of silver chloride, AgCl. The Ksp = 1.8 × 10-10 for silver chloride.
a. 9.0 × 10-11
b. 2.6 × 10-5
c. 1.3 × 10-5
d. 9.5 × 10-6
e. 6.5 × 10-6
Calculate the molar solubility of silver chloride, AgCl, in a solution in which [Cl-] = 0.045 M. Ksp =
1.8 × 10-10 for silver chloride.
a. 1.3 × 10-5
b. 2.5 × 10-8
c. 4.0 × 10-9
d. 1.8 × 10-10
e. 8.1 × 10-12
In which solution would calcium phosphate have the greatest solubility?
a. 0.1 M phosphoric acid
b. 0.1 M sodium phosphate
c. saturated calcium hydroxide
d. 0.1 M calcium nitrate
e. distilled water
Which of the factors affecting solubility explains the observation that silver phosphate is more
soluble in water than in Na3PO4?
a. The solubility of most salts increases as temperature increases.
b. The solubility of many salts is affected by the pH of the solution.
c. A common ion displaces the solubility equilibrium toward the undissolved solute.
d. The formation of complex ions displaces the solubility equilibrium toward the aqueous ions.
e. Some insoluble compounds are amphoteric.
Which of the factors affecting solubility explains the observation that the solubility of silver chloride
is increased by the addition of NH3 to the mixture?
a. The solubility of most salts increases as temperature increases.
b. The solubility of many salts is affected by the pH of the solution.
c. A common ion displaces the solubility equilibrium toward the undissolved solute.
d. The formation of complex ions displaces the solubility equilibrium toward the aqueous ions.
e. Some insoluble compounds are amphoteric.
Which group contains only solutes that would decrease the solubility of barium sulfate?
a. Ba(OH)2, NaOH, NH4OH
b. HNO3, H2SO4, HCH3COO
c. Na2SO4, NaOH, NaCH3COO
d. Ba(NO3)2, Na2SO4, H2SO4
e. SO2, CO2, NH3
Which mixture would dissolve the largest amount of barium sulfate?
a. 0.1 M HCl
b. 0.1 M Ba(NO3)2
c. 0.1 M Mg(OH)2
d. 0.1 M Na2SO4
e. 0.1 M Al2(SO4)3
Which mixture would dissolve the smallest amount of barium sulfate?
a. 0.1 M HCl
b. 0.1 M Ba(NO3)2
c. 0.1 M Mg(OH)2
d. 0.1 M Na2SO4
e. 0.1 M Al2(SO4)3
53. Which of the following species would increase the solubility of CuI?
a. NH3
b. SO42c. OHd. S2O32e. CN54. Will a precipitate form when 10.0 mL of 0.5 M NaCl is added to 10.0 mL of 0.05 M AgNO3? The Ksp
for AgCl is 1.8 × 10-10.
a. Yes because Q = K
b. Yes because Q > K
c. Yes because Q < K
d. No because Q < K
e. More information is needed to answer this question.
55. Calculate the volume of 0.10 M NaCl that must be added to 10.0 mL of 0.05 M AgNO3 to begin
forming a precipitate. The Ksp for AgCl is 1.8 × 10-10. Dilution effects can be neglected; 1 mL = 20
drops.
a. less than one drop
b. 1.0 mL
c. 2.5 mL
d. 5.0 mL
e. 10.0 mL
Answers to Sample Test 4 Questions
1. c. HSO42. d) HNO2 + OH- ---- NO2- + H2O
3. a) 4.7
4. c) 4.22
5. b) 2 and 7
6. a. 7.0
7. c. 2.71
8. c. basic
9. d. the pKa of an equimolar mixture of an unknown weak acid and its conjugate base without
knowing the pH
10. b 9.0
11. c. NaC2H3O2
12. b. an indicator.
13. c. 4.4
14. b. yellow
15. b. an indicator with Ka = 10-11
16. b. pH = 3.00-5.00
17. d. methyl red 4.8-6.0
18. a. 4.20
19. c. 2.0 x 10-4.
20. c. resists changes in pH upon addition of acid or base.
21. c. NaCH3COO/CH3COOH
22. b. KCOOH
23. d. NH3 & NH4NO3
24. a. equal volumes of 0.1 M formic acid and 0.1 M sodium formate
25. c. 3.65
26. d. 4.14
27. e.
pH = pK a + log
[A - ]
[HA]
28. b. CH3COOH/CH3COOKa = 1.8 × 10-5
29. a. 1.38 to 1
30. c. 4.12
31. e. 9.40
32. e. 6.00 g
33. c. At the equivalence point the pH decreases sharply.
34. e. 1.55
35. b. 44.81 mL
36. a. 11.81
37. c. 5.30
38. b. 4.74
39. d. 8.
40. b. 11.74
41. b. KOH and H3PO4
42. c. 5.6
43. e. CO2, NO2, SO2
44. a. [Ag+]3[PO43-]
45. c. 1.3 × 10-5
46. c. 4.0 × 10-9
47. e. distilled water
48. c. A common ion displaces the solubility equilibrium toward the undissolved solute.
49. d. The formation of complex ions displaces the solubility equilibrium toward the aqueous ions.
50. d. Ba(NO3)2, Na2SO4, H2SO4
51. a. 0.1 M HCl
52. e. 0.1 M Al2(SO4)3
53. a. NH3
54. b. Yes because Q > K
55. a. less than one drop