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Chapter 15: Acid-Base Equilibria
1. Calculate the [H+] in a solution that is 0.16 M in NaF and 0.25 M in HF. (Ka = 7.2 10–4)
A)
7.2 10–4 M
B)
1.6 M
C)
1.1 10–3 M
D)
0.20 M
E)
4.6 10–4 M
2. For a solution equimolar in HCN and NaCN, which statement is false?
A)
This is an example of the common ion effect.
B)
The [H+] is larger than it would be if only the HCN was in solution.
C)
The [H+] is equal to the Ka.
D)
Addition of more NaCN will shift the acid dissociation equilibrium of HCN to the
left.
E)
Addition of NaOH will increase [CN–] and decrease [HCN].
3. What will happen if a small amount of hydrochloric acid is added to a 0.1 M solution of HF?
A)
The percent ionization of HF will increase.
B)
The percent ionization of HF will decrease.
C)
The percent ionization of HF will remain unchanged.
D)
Ka for HF will increase.
E)
Ka for HF will decrease.
4. What will happen if a small amount of sodium hydroxide is added to a 0.1 M solution of
ammonia?
A)
Kb for ammonia will increase.
B)
Kb for ammonia will decrease.
C)
The percent ionization of ammonia will increase.
D)
The percent ionization of ammonia will decrease.
E)
The percent ionization of ammonia will remain unchanged.
5. 15.0 mL of 0.50 M HCl is added to a 100.-mL sample of 0.456 M HNO2 (Ka for
HNO2 = 4.0 10–4). What is the equilibrium concentration of NO2– ions?
A)
2.4 10–3 M
B)
1.6 10–4 M
C)
4.0 10–1 M
D)
4.9 10–2 M
E)
none of these
6. 15.0 mL of 0.50 M NaOH is added to a 100.-mL sample of 0.457 M NH3 (Kb for
NH3 = 1.8 10–5). What is the equilibrium concentration of NH4+ ions?
A)
1.0 10–2 M
B)
7.2 10–6 M
C)
4.0 10–1 M
D)
1.1 10–4 M
E)
none of these
7. What is the percent dissociation of HNO2 when 0.064 g of sodium nitrite is added to
120.0 mL of a 0.057 M HNO2 solution? Ka for HNO2 is 4.0 10–4.
A)
13%
B)
0.30%
C)
5.2%
D)
0.075%
E)
8.4%
8. Which of the following mixtures would result in a buffered solution?
A)
Mixing 100.0 mL of 0.100 M HCl with 100.0 mL of 0.100 M NaOH.
B)
Mixing 100.0 mL of 0.100 M NH3 (Kb = 1.8 10–5) with 100.0 mL of 0.100 M
NaOH.
C)
Mixing 100.0 mL of 0.100 M HCl with 100.0 mL of 0.100 M NH3 (Kb = 1.8 10–
5).
D)
Mixing 50.0 mL of 0.100 M HCl with 100.0 mL of 0.100 M NH3 (Kb = 1.8 10–5).
E)
At least two of the above mixtures would result in a buffered solution.
9. Which of the following will not produce a buffered solution?
A)
100 mL of 0.1 M Na2CO3 and 50 mL of 0.1 M HCl
B)
100 mL of 0.1 M NaHCO3 and 25 mL of 0.2 M HCl
C)
100 mL of 0.1 M Na2CO3 and 75 mL of 0.2 M HCl
D)
50 mL of 0.2 M Na2CO3 and 5 mL of 1.0 M HCl
E)
100 mL of 0.1 M Na2CO3 and 50 mL of 0.1 M NaOH
10. 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.) (Kb for NH3 = 1.8 10–5; Kb
for C5H5N = 1.7 10–9)
A)
1.0 mole NH3 and 1.5 mole NH4Cl
B)
1.5 mole NH3 and 1.0 mole NH4Cl
C)
1.0 mole C5H5N and 1.5 mole C5H5NHCl
D)
1.5 mole C5H5N and 1.0 mole C5H5NHCl
E)
none of these
11. Suppose a buffer solution is made from formic and (HCHO2) and sodium formate
(NaCHO2). What is the net ionic equation for the reaction that occurs when a small amount
of hydrochloric acid is added to the buffer?
A)
H3O+(aq) + OH–(aq) → 2H2O(l)
B)
H3O+(aq) + HCHO2(aq) → H2O(l) + H2CHO2+(aq)
C)
HCl(aq)+ OH–(aq) → H2O(l) + Cl–(aq)
D)
HCl(aq) + CHO2–(aq) → HCHO2(aq) + Cl–(aq)
E)
H3O+(aq) + CHO2–(aq) → HCHO2(aq) + H2O(l)
12. Suppose a buffer solution is made from formic acid, HCHO2, and sodium formate,
NaCHO2. What is the net ionic equation for the reaction that occurs when a small amount of
sodium hydroxide is added to the buffer?
A)
NaOH(aq) + H3O+(aq) → Na+(aq) + 2H2O(l)
B)
H3O+(aq) + OH–(aq) → 2H2O(l)
C)
OH–(aq) + HCHO2(aq) → CHO2–(aq) + H2O(l)
D)
NaOH(aq) + HCHO2(aq) → NaCHO2(aq) + H2O(l)
E)
Na+(aq) + HCHO2(aq) → NaH(aq) + HCO2+(aq)
13. A weak acid, HF, is in solution with dissolved sodium fluoride, NaF. If HCl is added, which
ion will react with the extra hydrogen ions from the HCl to keep the pH from changing?
A)
OH–
B)
Na+
C)
F–
D)
Na–
E)
none of these
14. Which of the following is true for a buffered solution?
A)
The solution resists change in its [H+].
B)
The solution will not change its pH very much even if a concentrated acid is
added.
C)
The solution will not change its pH very much even if a strong base is added.
D)
Any H+ ions will react with a conjugate base of a weak acid already in solution.
E)
All of these.
15. A 100. mL sample of 0.10 M HCl is mixed with 50. mL of 0.11 M NH3. What is the
resulting pH? (Kb for NH3 = 1.8 10–5)
A)
3.09
B)
10.91
C)
12.48
D)
1.35
E)
1.52
16. The following question refers to a 2.0-liter buffered solution created from 0.72 M NH3
(Kb = 1.8 10–5) and 0.26 M NH4F. What is the pH of this solution?
A)
9.26
B)
9.70
C)
4.30
D)
5.18
E)
8.81
17. The following question refers to a 2.0-liter buffered solution created from 0.31 M NH3
(Kb = 1.8 10–5) and 0.26 M NH4F. When 0.10 mol of H+ ions is added to the solution what
is the pH?
A)
4.82
B)
4.66
C)
10.53
D)
9.18
E)
7.88
18. You have a 250.-mL sample of 1.250 M acetic acid (Ka = 1.8 10–5). Assuming no volume
change, how much NaOH must be added to make the best buffer?
A)
6.25 g
B)
12.5 g
C)
16.3 g
D)
21.3 g
E)
none of these
19. You have a 250.-mL sample of 1.28 M acetic acid (Ka = 1.8 10–5). Calculate the pH of the
best buffer.
A)
7.00
B)
4.74
C)
4.25
D)
9.26
E)
none of these
20. You have a 250.0-mL sample of 1.00 M acetic acid (Ka = 1.8 10–5). Calculate the pH after
adding 0.0050 mol of NaOH to 1.0 liter of the best buffer.
A)
7.05
B)
2.41
C)
3.54
D)
4.78
E)
none of these
21. You have a 250.0-mL sample of 1.00 M acetic acid (Ka = 1.8 10–5). Calculate the pH after
adding 0.0040 mol HCl to 1.0 liter of the best buffer.
A)
4.72
B)
2.35
C)
3.12
D)
6.98
E)
none of these
22. 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
liter 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
23. A solution contains 0.250 M HA (Ka = 1.0 10–6) and 0.45 M NaA. What is the pH after
0.17 mole of HCl is added to 1.00 L of this solution?
A)
0.77
B)
7.82
C)
5.82
D)
1.97
E)
8.18
24. The following question refers to the following system: A 1.0-liter solution contains 0.25 M
HF and 0.83 M NaF (Ka for HF is 7.2 10–4).
What is the pH of this solution?
A)
3.14
B)
3.66
C)
2.62
D)
0.52
E)
10.34
25. The following question refers to the following system: A 1.0-liter solution contains 0.25 M
HF and 0.32 M NaF (Ka for HF is 7.2 10–4).
If one adds 0.30 liters of 0.020 M KOH to the solution, what will be the change in pH?
A)
0.02
B)
3.27
C)
0.13
D)
–0.11
E)
–0.28
26. How many moles of HCl need to be added to 150.0 mL of 0.50 M NaZ to have a solution
with a pH of 6.50? (Ka of HZ is 2.3 10–5)? Assume negligible volume of the HCl.
A)
6.8 10–3
B)
7.5 10–2
C)
5.0 10–1
D)
1.0 10–3
E)
none of these
27. Calculate the pH of a solution that is 0.50 M in HF (Ka = 7.2 10–4) and 0.95 M in NaF.
A)
3.14
B)
3.42
C)
0.28
D)
10.58
E)
2.86
28. Calculate the pH of a solution that is 2.00 M HF, 1.00 M NaOH, and 0.642 M NaF.
(Ka = 7.2 10–4)
A)
3.14
B)
3.36
C)
2.65
D)
2.93
E)
none of these
29. Consider a solution consisting of the following two buffer systems:
H2CO3 HCO3– + H+ pKa = 6.4
H2PO4– HPO42– + H+ pKa = 7.2
At pH 6.4, which one of the following is true of the relative amounts of acid and conjugate
base present?
A)
[H2CO3] > [HCO3–] and [H2PO4–] > [HPO42–]
B)
[H2CO3] = [HCO3–] and [H2PO4–] > [HPO42–]
C)
[H2CO3] = [HCO3–] and [HPO42–] > [H2PO4–]
D)
[HCO3–] > [H2CO3] and [HPO42–] > [H2PO4–]
E)
[H2CO3] > [HCO3–] and [HPO42–] > [H2PO4–]
30. Given 100.0 mL of a buffer that is 0.50 M in HOCl and 0.77 M in NaOCl, what is the pH
after 10.0 mL of 1.0 M NaOH has been added? (Ka for HOCl = 3.5 10–8)
A)
7.68
B)
7.74
C)
7.46
D)
7.12
E)
7.79
31. How many moles of solid NaF would have to be added to 1.0 L of 2.39 M HF solution to
achieve a buffer of pH 3.35? Assume there is no volume change. (Ka for HF = 7.2 10–4)
A)
3.9
B)
0.50
C)
0.67
D)
1.0
E)
1.6
32. What is the pH of a solution that results when 0.010 mol HNO3 is added to 500. mL of a
solution that is 0.10 M in aqueous ammonia and 0.55 M in ammonium nitrate. Assume no
volume change. (The Kb for NH3 = 1.8 10–5.)
A)
9.26
B)
5.05
C)
10.11
D)
8.40
E)
8.71
33. How many mmoles of HCl must be added to 100 mL of a 0.100 M solution of methylamine
(pKb = 3.36) to give a buffer having a pH of 10.00?
A)
8.1
B)
18.7
C)
20.0
D)
41.5
E)
12.7
34. Calculate the pH of a solution made by mixing 100.0 mL of 0.635 M NH3 with 100.0 mL of
0.100 M HCl. (Kb for NH3 = 1.8 10–5)
A)
9.98
B)
4.02
C)
8.53
D)
9.26
E)
none of these
35. A solution contains 0.500 M HA (Ka = 1.0 10–8) and 0.432 M NaA. What is the [H+] after
0.10 mole of HCl is added to 1.00 L of this solution?
A)
1.0 10–8 M
B)
3.0 10–8 M
C)
5.5 1021 M
D)
1.8 10–8 M
E)
none of these
36. 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)
The solution is not a buffer because [HCN] is not equal to [CN–].
B)
The pH will be below 7.00 because the concentration of the acid is greater than
that of the base.
C)
[OH–] > [H+]
D)
The buffer will be more resistant to pH changes from addition of strong acid than
of strong base.
E)
All of the above are false.
37. Which of the following solutions will be the best buffer at a pH of 9.26? (Ka for HC2H3O2 is
1.8 10–5, Kb for NH3 is 1.8 10–5).
A)
0.10 M HC2H3O2 and 0.10 M Na C2H3O2
B)
5.0 M HC2H3O2 and 5.0 M Na C2H3O2
C)
0.10 M NH3 and 0.10 M NH4Cl
D)
5.0 M NH3 and 5.0 M NH4Cl
E)
5.0 M HC2H3O2 and 5.0 M NH3
You have two buffered solutions. Buffered solution 1 consists of 5.0 M HOAc and 5.0 M
NaOAc; buffered solution 2 is made of 0.050 M HOAc and 0.050 M NaOAc.
38. How do the pHs of the buffered solutions compare?
A)
The pH of buffered solution 1 is greater than that of buffered solution 2.
B)
The pH of buffered solution 2 is greater than that of buffered solution 1.
C)
The pH of buffered solution 1 is equal to that of buffered solution 2.
D)
Cannot be determined without the Ka values.
E)
None of these (A-D).
39. Buffered solution 1 has a greater buffering capacity than buffered solution 2.
40. In titrating 0.20 M hydrochloric acid, HCl, with 0.20 M NaOH at 25°C, the solution at the
equivalence point is
A)
0.20 M NaCl
B)
very acidic
C)
slightly acidic
D)
0.10 M HCl and 0.20 M NaOH
E)
0.10 M NaCl
41. One milliliter (1.00 mL) of acid taken from a lead storage battery is pipetted into a flask.
Water and phenolphthalein indicator are added, and the solution is titrated with 0.55 M
NaOH until a pink color appears; 12.0 mL are required. The number of grams of H2SO4
(formula weight = 98) present in one liter of the battery acid is:
A)
647
B)
323
C)
30
D)
1294
E)
54
42. You are given 5.00 mL of an H2SO4 solution of unknown concentration. You divide the
5.00-mL sample into five 1.00-mL samples and titrate each separately with 0.1000 M
NaOH. In each titration the H2SO4 is completely neutralized. The average volume of NaOH
solution used to reach the endpoint is 15.6 mL. What was the concentration of H2SO4 in the
5.00-mL sample?
A)
1.56 M
B)
3.90 M
C)
0.780 M
D)
0.156 M
E)
7.80 M
43. What is the molarity of a sodium hydroxide solution if 28.7 mL of this solution reacts
exactly with 22.30 mL of 0.253 M sulfuric acid?
A)
0.197 M
B)
0.786 M
C)
7.26 M
D)
0.393 M
E)
0.221 M
44. If 25.0 mL of 0.451 M NaOH solution is titrated with 0.253 M H2SO4, the flask at the
endpoint will contain (besides the indicator phenolphthalein) as the principal components:
A)
sodium hydroxide, sulfuric acid, and water
B)
dissolved sodium sulfate and water
C)
sodium hydroxide, sodium sulfate, and water
D)
dissolved sodium sulfate, sulfuric acid, and water
E)
precipitated sodium sulfate and water
45. A 19.0-mL sample of tartaric acid is titrated to a phenolphthalein endpoint with 20. mL of
1.0 M NaOH. Assuming tartaric acid is diprotic, what is the molarity of the acid?
A)
1.0 M
B)
0.53 M
C)
2.0 M
D)
1.05 M
E)
impossible to determine
46. If 25 mL of 0.750 M HCl are added to 100. mL of 0.352 M NaOH, what is the final pH?
A)
13.12
B)
0.88
C)
13.45
D)
0.55
E)
7.00
47. A 50.00-mL sample of 0.100 M KOH is titrated with 0.202 M HNO3. Calculate the pH of
the solution after 52.00 mL of HNO3 is added.
A)
12.73
B)
0.99
C)
1.27
D)
13.01
E)
none of these
48. A solution of hydrochloric acid of unknown concentration was titrated with 0.23 M NaOH.
If a 100.-mL sample of the HCl solution required exactly 10. mL of the NaOH solution to
reach the equivalence point, what was the pH of the HCl solution?
A)
12.4
B)
1.6
C)
–0.4
D)
3.3
E)
6.6
49. A titration of 200.0 mL of 1.32 M H2A was done with 1.00 M NaOH. For the diprotic acid
H2A, Ka1 = 2.5 10–5, Ka2 = 3.1 10–9. Calculate the pH before any 1.00 M NaOH has been
added.
A)
11.76
B)
4.48
C)
9.52
D)
8.96
E)
2.24
50. A titration of 200.0 mL of 1.00 M H2A was done with 1.38 M NaOH. For the diprotic acid
H2A, Ka1 = 2.5 10–5, Ka2 = 3.1 10–9. Calculate the pH after 100.0 mL of 1.38 M NaOH
have been added.
A)
9.05
B)
8.86
C)
5.14
D)
9.90
E)
4.95
51. A titration of 200.0 mL of 1.00 M H2A was done with 1.02 M NaOH. For the diprotic acid
H2A, Ka1 = 2.5 10–5, Ka2 = 3.1 10–9. Calculate the pH after 600.0 mL of 1.02 M NaOH
have been added.
A)
13.423
B)
0.577
C)
13.712
D)
0.288
E)
9.423
52. Consider the titration of 300.0 mL of 0.577 M NH3 (Kb = 1.8 10–5) with 0.500 M HNO3.
After 150.0 mL of 0.500 M HNO3 have been added, the pH of the solution is:
A)
4.63
B)
11.37
C)
6.37
D)
9.37
E)
none of these
53. Consider the titration of 300.0 mL of 0.450 M NH3 (Kb = 1.8 10–5) with 0.450 M HNO3.
How many milliliters of 0.450 M HNO3 are required to reach the stoichiometric point of the
reaction?
A)
3.50 102 mL
B)
4.00 102 mL
C)
4.50 102 mL
D)
3.00 102 mL
E)
none of these
54. Consider the titration of 300.0 mL of 0.450 M NH3 (Kb = 1.8 10–5) with 0.450 M HNO3.
At the stoichiometric point of this titration, the pH is:
A)
4.80
B)
2.70
C)
4.95
D)
4.74
E)
7.00
55. Consider the titration of 500.0 mL of 0.200 M NaOH with 0.800 M HCl. How many
milliliters of 0.800 M HCl must be added to reach a pH of 13.000?
A)
55.6 mL
B)
24.6 mL
C)
18.5 mL
D)
12.9 mL
E)
4.32 mL
56. What quantity of NaOH(s) must be added to 2.00 L of 0.434 M HCl to achieve a pH of
13.00? (Assume no volume change.)
A)
0.67 mol
B)
1.07 mol
C)
0.20 mol
D)
1.00 10–13 mol
E)
none of these
57. A 50.0-mL sample of 0.10 M HNO2 (Ka = 4.0 10–4) is titrated with 0.12 M NaOH. The pH
after 25.0 mL of NaOH have been added is
A)
10.43
B)
7.00
C)
6.57
D)
3.57
E)
none of these
58. The pH at the equivalence point of the titration of a strong acid with a strong base is:
A)
3.9
B)
4.5
C)
7.0
D)
8.2
E)
none of these
59. The pH at the equivalence point of a titration of a weak acid with a strong base will be
A)
less than 7.00
B)
equal to 7.00
C)
greater than 7.00
D)
equal to the pKa of the acid
E)
more data needed to answer this question
60. A 75.0-mL sample of 0.0650 M HCN (Ka = 6.2 10–10) is titrated with 0.65 M NaOH. What
volume of 0.65 M NaOH is required to reach the stoichiometric point?
A)
750. mL
B)
7.50 mL
C)
3.75 mL
D)
75.0 mL
E)
cannot determine without knowing the pH at the stoichiometric point
