Chapter 16 – Acid-Base Equilibria and Solubility Equilibria
39. Calculate the pH at the equivalence point for the titration of 0.20 M HCl with 0.20 M NH3
(Kb = 1.8 10–5).
40. What is the pH at the equivalence point in the titration of 100 mL of 0.10 M HCl with
0.10 M NaOH?
Chapter 16 – Acid-Base Equilibria and Solubility Equilibria
41. What is the pH at the equivalence point in the titration of 100 mL of 0.10 M HCN (Ka =
4.9 10–10) with 0.10 M NaOH?
42. Calculate the pH of the solution resulting from the addition of 10.0 mL of 0.10 M NaOH
to 50.0 mL of 0.10 M HCN (Ka = 4.9 10–10) solution.
Chapter 16 – Acid-Base Equilibria and Solubility Equilibria
43. Calculate the pH of the solution resulting from the addition of 75.0 mL of 0.15 M KOH to
35.0 mL of 0.20 M HCN (Ka (HCN) = 4.9 10–10).
44. Calculate the pH of the solution resulting from the addition of 25.0 mL of 0.20 M HCl to
50.0 mL of 0.10 M aniline (C6H5NH2). Kb (C6H5NH2) = 3.8 x 10–10
Chapter 16 – Acid-Base Equilibria and Solubility Equilibria
45. Calculate the pH of the solution resulting from the addition of 85.0 mL of 0.35 M HCl to
30.0 mL of 0.40 M aniline (C6H5NH2). Kb (C6H5NH2) = 3.8 x 10–10
46. Methyl red is a common acid-base indicator. It has a Ka equal to 6.3 10–6. Its un-ionized
form is red and its anionic form is yellow. What color would a methyl red solution have at pH
= 7.8?
Chapter 16 – Acid-Base Equilibria and Solubility Equilibria
47. What mass of sodium fluoride must be added to 250. mL of a 0.100 M HF solution to give
a buffer solution having a pH of 3.50? [Ka(HF) = 7.1 10–4]
48. What mass of ammonium nitrate must be added to 350. mL of a 0.150 M solution of
ammonia to give a buffer having a pH of 9.00? (Kb(NH3) = 1.8 10–5)
Chapter 16 – Acid-Base Equilibria and Solubility Equilibria
49. 40.0 ml of an acetic acid of unknown concentration is titrated with 0.100 M NaOH. After
20.0 mL of the base solution has been added, the pH in the titration flask is 5.10. What was
the concentration of the original acetic acid solution? [Ka(CH3COOH) = 1.8 10–5]
50. 25.0 mL of a hydrofluoric acid solution of unknown concentration is titrated with 0.200 M
NaOH. After 20.0 mL of the base solution has been added, the pH in the titration flask is 3.00.
What was the concentration of the original hydrofluoric acid solution. [Ka(HF) = 7.1 10–4]
Chapter 16 – Acid-Base Equilibria and Solubility Equilibria
51. For PbCl2 (Ksp = 2.4 10–4), will a precipitate of PbCl2 form when 0.10 L of 3.0 10–2 M
Pb(NO3)2 is added to 400 mL of 9.0 10–2 M NaCl?
52. The solubility of lead(II) iodide is 0.064 g/100 mL at 20ºC. What is the solubility product
for lead(II) iodide?
Chapter 16 – Acid-Base Equilibria and Solubility Equilibria
53. The molar solubility of magnesium carbonate is 1.8 10–4 mol/L. What is Ksp for this
compound?
54. The molar solubility of manganese(II) carbonate is 4.2 10–6 M. What is Ksp for this
compound?
Chapter 16 – Acid-Base Equilibria and Solubility Equilibria
55. The molar solubility of tin(II) iodide is 1.28 10–2 mol/L. What is Ksp for this
compound?
56. The solubility of strontium carbonate is 0.0011 g/100 mL at 20ºC. Calculate the Ksp value
for this compound.
Chapter 16 – Acid-Base Equilibria and Solubility Equilibria
57. The molar solubility of lead(II) iodate in water is 4.0 10–5 mol/L. Calculate Ksp for
lead(II) iodate.
58. The solubility product for chromium(III) fluoride is Ksp = 6.6 10–11. What is the molar
solubility of chromium(III) fluoride?
Chapter 16 – Acid-Base Equilibria and Solubility Equilibria
59. The solubility product for barium sulfate is 1.1 10–10. Calculate the molar solubility of
barium sulfate.
60. The Ksp for silver(I) phosphate is 1.8 10–18. Calculate the molar solubility of silver(I)
phosphate.
Chapter 16 – Acid-Base Equilibria and Solubility Equilibria
61. The solubility product for calcium phosphate is Ksp = 1.3 10–26. What is the molar
solubility of calcium phosphate?
62. The Ksp value for lead(II) chloride is 2.4 10–4. What is the molar solubility of lead(II)
chloride?
Chapter 16 – Acid-Base Equilibria and Solubility Equilibria
63. Calculate the silver ion concentration in a saturated solution of silver(I) carbonate (Ksp =
8.1 10–12).
64. The Ksp for silver(I) phosphate is 1.8 10–18. Determine the silver ion concentration in a
Chapter 16 – Acid-Base Equilibria and Solubility Equilibria
65. Calculate the silver ion concentration in a saturated solution of silver(I) sulfate (Ksp = 1.4
10–5).
66. Calculate the concentration of chloride ions in a saturated lead(II) chloride (Ksp = 2.4
10–4) solution.
Chapter 16 – Acid-Base Equilibria and Solubility Equilibria
67. Calculate the concentration of fluoride ions in a saturated barium fluoride (Ksp = 1.7 10–
6) solution.
68. Which of the following would decrease the Ksp for PbI2?
Chapter 16 – Acid-Base Equilibria and Solubility Equilibria
69. Calculate the minimum concentration of Mg2+ that must be added to 0.10 M NaF in order
to initiate a precipitate of magnesium fluoride. (For MgF2 , Ksp = 6.9 10–9.)
70. Calculate the minimum concentration of Cr3+ that must be added to 0.095 M NaF in order
to initiate a precipitate of chromium(III) fluoride. (For CrF3 , Ksp = 6.6 10–11.)
Chapter 16 – Acid-Base Equilibria and Solubility Equilibria
71. Will a precipitate form (yes or no) when 50.0 mL of 1.2 10–3 M Pb(NO3)2 are added to
50.0 mL of 2.0 10–4 M Na2S? If so, identify the precipitate.
72. Will a precipitate of magnesium fluoride form when 300. mL of 1.1 10–3 M MgCl2 are
added to 500. mL of 1.2 10–3 M NaF? [Ksp (MgF2) = 6.9 10–9]
Chapter 16 – Acid-Base Equilibria and Solubility Equilibria
73. Will a precipitate of magnesium fluoride form when 200. mL of 1.9 10–3 M MgCl2 are
added to 300. mL of 1.4 10–2 M NaF? [Ksp (MgF2) = 6.9 10–9]
74. Will a precipitate (ppt) form when 300. mL of 5.0 10–5 M AgNO3 are added to 200. mL
of 2.5 10–7 M NaBr? Answer yes or no, and identify the precipitate if there is one.
Chapter 16 – Acid-Base Equilibria and Solubility Equilibria
75. Will a precipitate (ppt) form when 20.0 mL of 1.1 10–3 M Ba(NO3)2 are added to 80.0
mL of 8.4 10–4 M Na2CO3?
76. Will a precipitate (ppt) form when 300. mL of 2.0 10–5 M AgNO3 are added to 200. mL
of 2.5 10–9 M NaI? Answer yes or no, and identify the precipitate if there is one.
Chapter 16 – Acid-Base Equilibria and Solubility Equilibria
77. Which response has both answers correct? Will a precipitate form when 250 mL of 0.33
M Na2CrO4 are added to 250 mL of 0.12 M AgNO3? [Ksp(Ag2CrO4) = 1.1 10–12] What is
the concentration of the silver ion remaining in solution?
78. To 1.00 L of a 0.100 M aqueous solution of benzoic acid (C6H5COOH) is added 1.00 mL
of 12.0 M HCl. What is the percentage ionization of the benzoic acid in the resulting solution?
[Ka(C6H5COOH) = 6.5 10–5]