Ch 8 Practice Problems
1. What combination of substances will give a buffered solution that has a pH of 5.05? Assume each pair of substances is
dissolved in 5.0 L of water. (Ka for NH4+ = 5.6  10–10; Ka for C5H5NH+ = 5.9  10–6)
A) 1.0 mol NH3 and 1.5 mol NH4Cl
B) 1.5 mol NH3 and 1.0 mol NH4Cl
C) 1.0 mol C5H5N and 1.5 mol C5H5NHCl
D) 1.5 mol C5H5N and 1.0 mol C5H5NHCl
E)
none of these
2. You have solutions of 0.200 M HNO2 and 0.200 M KNO2 (Ka for HNO2 = 4.00  10–4). A buffer of pH 3.000 is needed.
What volumes of HNO2 and KNO2 are required to make 1 L of buffered solution?
A) 500 mL of each
B) 286 mL HNO2; 714 mL KNO2
C) 413 mL HNO2; 587 mL KNO2
D) 714 mL HNO2; 286 mL KNO2
E)
587 mL HNO2; 413 mL KNO2
3. Calculate [H+] in a solution that is 0.34 M in NaF and 0.58 M in HF. (Ka = 7.2  10–4)
A)
B)
C)
D)
E)
0.58 M
4.2  10–4 M
1.2  10–3 M
2.0  10–2 M
1.1  10–4 M
4. Calculate the pH of a solution made by mixing 33.0 mL of 0.340 M NaA (Ka for HA = 1.0  10-8) with 26.4 mL of 0.100
M HCl.
A) 8.63
B) 7.49
C) 8.51
D) 7.62
E)
8.00
5. Consider a solution of 2.0 M HCN and 1.0 M NaCN (Ka for HCN = 6.2  10–10). Which of the following statements is
true?
A)
B)
C)
D)
E)
The solution is not a buffer because [HCN] is not equal to [CN–].
The pH will be below 7.00 because the concentration of the acid is greater than that of the base.
[OH–] > [H+]
The buffer will be more resistant to pH changes from addition of strong acid than to pH changes from addition
of strong base.
All of these statements are false.
6. Which of the following will not produce a buffered solution?
A)
B)
C)
D)
E)
100 mL of 0.1 M Na2CO3 and 50 mL of 0.1 M HCl
100 mL of 0.1 M NaHCO3 and 25 mL of 0.2 M HCl
100 mL of 0.1 M Na2CO3 and 75 mL of 0.2 M HCl
50 mL of 0.2 M Na2CO3 and 5 mL of 1.0 M HCl
100 mL of 0.1 M Na2CO3 and 50 mL of 0.1 M NaOH
7. For ammonium ion, Ka is 5.6  10-10 . To make a buffered solution with pH 10.0, the ratio of NH3 to NH4Cl must be
A)
B)
C)
D)
E)
1.8 : 1
1 : 5.6
0.18 : 1
5.6 : 1
none of these
8. How much solid NaCN must be added to 1.0 L of a 0.5 M HCN solution to produce a solution with pH 7.0? Ka = 6.2 
10–10 for HCN.
A) 0.0034 g
B) 11 g
C) 160 g
D) 24 g
E)
0.15 g
9. For 105.0 mL of a buffer that is 0.40 M in HOCl and 0.36 M in NaOCl, what is the pH after 13.4 mL of 1.5 M NaOH is
added? Ka
A)
B)
C)
D)
E)
for HOCl = 3.5  10–8. (Assume the volumes are additive.)
7.03
6.91
7.41
7.50
7.88
10. What is the pH of a solution that results when 0.010 mol HNO3 is added to 655.0 mL of a solution that is 0.22 M in
aqueous ammonia and 0.50 M in ammonium nitrate. Assume no volume change. (Ka for NH4+ = 5.6  10-10 .)
A) 5.15
B) 9.66
C) 4.34
D) 8.85
E)
8.90
11. Which of the following solutions will be the best buffer at a pH of 4.74? (Ka for HC2H3O2 is 1.8  10-5; Ka for NH4+ is
5.6  10-10.)
A) 0.10 M HC2H3O2 and 0.10 M NaC2H3O2
B) 5.0 M HC2H3O2 and 5.0 M NH4Cl
C) 0.10 M NH3 and 0.10 M NH4Cl
D) 5.0 M HC2H3O2 and 5.0 M NaC2H3O2
E)
5.0 M HC2H3O2 and 5.0 M NH3
12. A solution of hydrochloric acid of unknown concentration was titrated with 0.16 M NaOH. If a 350-mL sample of the
HCl solution required exactly 17 mL of the NaOH solution to reach the equivalence point, what was the pH of the
HCl solution?
A) 11.89
B) 2.11
C) 2.13
D) 0.16
E)
–0.52
13. What is the molarity of a sodium hydroxide solution if 25.0 mL of this solution reacts exactly with 22.30 mL of 0.253
M sulfuric acid?
A) 0.113 M
B) 0.226 M
C) 0.284 M
D) 0.451 M
E)
0.567 M
14. Equal volumes of 0.1 M HCl and 0.1 M HC2H3O2 are titrated with 0.1 M NaOH. Which of the following would be
equal for both titrations?
A) the initial pH
B) the pH at the halfway point
C) the pH at the equivalence point
D) the volume of NaOH added to reach the equivalence point
E)
two of the above
15. A 74.60-mL sample of 0.0758 M HCN (Ka = 6.2  10–10) is titrated with 0.610 M NaOH. What volume of 0.610 M
NaOH is required to reach the stoichiometric point?
A) 600 mL
B) 9.27 mL
C) 74.6 mL
D) 5.65 mL
E)
0.00167 mL
16. 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–  KNO2 + H2O
B) HNO2 + H2O  NO2– + H3O+
C) HNO2 + KOH  K+ + NO2– + H2O
D) HNO2 + OH–  NO2– + H2O
E)
H+ + OH–  H2O
17. The pH at the equivalence point of a titration of a weak acid with a strong base is
A)
B)
C)
D)
less than 7.00.
equal to 7.00.
greater than 7.00.
More data are needed to answer this question.
18. If 16 mL of 0.74 M HCl is added to 114 mL of 0.26 M NaOH, what is the final pH?
A)
B)
C)
D)
E)
13.14
0.86
13.36
0.64
7.00
19. A 50.00-mL sample of 0.100 M KOH is titrated with 0.100 M HNO3. Calculate the pH of the solution after the 52.00
mL of HNO3 is added.
A) 6.50
B) 3.01
C) 2.71
D) 2.41
E)
none of these
20. A 74.0-mL sample of 0.13 M HNO2 (Ka = 4.0  10–4) is titrated with 0.14 M NaOH. What is the pH after 25.8 mL of
NaOH has been added?
A) 7.00
B) 10.38
C) 3.62
D) 3.18
E)
2.97
21. A 75.0-mL sample of 0.0500 M HCN (Ka = 6.2  10–10) is titrated with 0.500 M NaOH. What is [H+] in the solution
after 3.0 mL of 0.500 M NaOH has been added?
A) 1.0  10–7 M
B) 4.1  10–10 M
C) 5.2  10–13 M
D) 9.3  10–10 M
E)
none of these
22. A 100.-mL sample of 0.10 M HCl is mixed with 50. mL of 0.10 M NH3. What is the resulting pH?
(Ka for NH4+ = 5.6  10–10)
A) 12.52
B) 3.87
C) 1.30
D) 7.85
E)
1.48
23. After adding 25.0 mL of 0.100 M NaOH to 100.0 mL of 0.100 M weak acid (HA), the pH is found to be 5.90.
Determine the value of Ka for the acid HA.
A) 1.6  10-11
B) 4.2  10-7
C) 2.1  10-5
D) 3.5  10-9
E)
none of these
24. The following plot show the pH curves for the titrations of various acids, HA, by 0.10 M NaOH. At the start of the
titration, all of the acids were 50.0 mL of 0.10 M HA.
Which pH curve corresponds to an acid with Ka = 2 × 10– 6?
A) a
B) b
C) c
D) d
E) e
25. A 50.0-mL sample of 0.10 M HNO2 is titrated with 0.10 M NaOH. What is the pH after 50.0 mL of NaOH
have been added? (Ka of HNO2 is 4.00 x 10-4)
A) 7.00
B) 1.00
C) 5.95
D) 10.25
E) 8.05
26. Silver chromate, Ag2CrO4, has a Ksp of 9.0  10–12. Calculate the solubility, in moles per liter, of silver chromate.
A)
B)
C)
D)
E)
1.3  10–4 M
7.8  10–5 M
9.5  10–7 M
1.9  10–12 M
9.8  10–5 M
27. The concentration of OH– in a saturated solution of Mg(OH)2 is 3.6  10–4 M. What is Ksp for Mg(OH)2?
A)
B)
C)
D)
E)
1.3  10–7
4.7  10–11
1.2  10–11
3.6  10–4
none of these
28. The solubility of Cd(OH)2 in water is 1.7  10–5 mol/L at 25°C. What is Ksp for Cd(OH)2?
A)
B)
C)
D)
E)
2.0  10–14
4.9  10–15
5.8 10–10
2.9  10–10
none of these
29. The solubility of La(IO3)3 in a 0.10 M KIO3 solution is 1.0  10–7 mol/L. Calculate Ksp for La(IO3)3.
1.0  10–8
2.7  10–9
1.0  10–10
2.7  10–27
none of these
A)
B)
C)
D)
E)
30. Calculate the concentration of chromate ion, CrO42–, in a saturated solution of CaCrO4 (Ksp = 7.1  10–4).
A)
B)
C)
D)
E)
0.027 M
5.0  10–7 M
7.1  10–4 M
3.5  10–4 M
3.5  10–2 M
31. The value of Ksp for AgI is 1.5  10–16. Calculate the solubility, in moles per liter, of AgI in a 0.50 M NaI solution.
1.7  10–8 mol/L
1.2  10–8 mol/L
7.5  10–17 mol/L
2.8  10-9 mol/L
3.0  10–16 mol/L
A)
B)
C)
D)
E)
32. Calculate the solubility of Ca3(PO4)2(s) (Ksp = 1.3  10–32) in a 1.0  10–2 M Ca(NO3)2 solution.
5.7  10–14 mol/L
6.2  10–7 mol/L
1.6  10–14 mol/L
3.16  10–12 mol/L
none of these
A)
B)
C)
D)
E)
Answers:
1. C
13. D
25. E
2. D
14. D
26. A
3. C
15. B
27. E
4. C
16. D
28. A
5. C
17. C
29. C
6. E
18. A
30. A
7. D
19. C
31. E
8. E
20. D
32. A
9. E
21. D
10. D
22. E
11. D
23. B
12. B
24. D