Chapter 15 1 What is the equilibrium concentration of NH4+ ions

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