14.11: Exercises

1.

Write equations that show NH₃ as both a conjugate acid and a conjugate base.

2.

Write equations that show acting both as an acid and as a base.

3.

Show by suitable net ionic equations that each of the following species can act as a Brønsted-Lowry acid:

1. (f)

4.

Show by suitable net ionic equations that each of the following species can act as a Brønsted-Lowry acid:

1. (f) HS⁻

5.

Show by suitable net ionic equations that each of the following species can act as a Brønsted-Lowry base:

1. (f)

6.

Show by suitable net ionic equations that each of the following species can act as a Brønsted-Lowry base:

1. (f)

7.

What is the conjugate acid of each of the following? What is the conjugate base of each?

1. OH⁻
2. H₂O
4. NH₃
6. H₂O₂
7. HS⁻

8.

What is the conjugate acid of each of the following? What is the conjugate base of each?

1. H₂S
3. PH₃
4. HS⁻
7. H₄N₂
8. CH₃OH

9.

Identify and label the Brønsted-Lowry acid, its conjugate base, the Brønsted-Lowry base, and its conjugate acid in each of the following equations:

1. (f)

10.

Identify and label the Brønsted-Lowry acid, its conjugate base, the Brønsted-Lowry base, and its conjugate acid in each of the following equations:

1. (f)

11.

What are amphiprotic species? Illustrate with suitable equations.

12.

State which of the following species are amphiprotic and write chemical equations illustrating the amphiprotic character of these species:

1. H₂O
3. S²⁻

13.

State which of the following species are amphiprotic and write chemical equations illustrating the amphiprotic character of these species.

1. NH₃
3. Br⁻

14.

Is the self-ionization of water endothermic or exothermic? The ionization constant for water (K_w) is 2.9 \times 10^{−14} at 40 °C and 9.3 \times 10^{−14} at 60 °C.

15.

Explain why a sample of pure water at 40 °C is neutral even though [H₃O⁺] = 1.7 \times 10^{−7} M. K_w is 2.9 \times 10^{−14} at 40 °C.

16.

The ionization constant for water (K_w) is 2.9 \times 10^{−14} at 40 °C. Calculate [H₃O⁺], [OH⁻], pH, and pOH for pure water at 40 °C.

17.

The ionization constant for water (K_w) is 9.311 \times 10^{−14} at 60 °C. Calculate [H₃O⁺], [OH⁻], pH, and pOH for pure water at 60 °C.

18.

Calculate the pH and the pOH of each of the following solutions at 25 °C for which the substances ionize completely:

1. 0.200 M HCl
2. 0.0143 M NaOH
3. 3.0 M HNO₃
4. 0.0031 M Ca(OH)₂

19.

Calculate the pH and the pOH of each of the following solutions at 25 °C for which the substances ionize completely:

1. 0.000259 M HClO₄
2. 0.21 M NaOH
3. 0.000071 M Ba(OH)₂
4. 2.5 M KOH

20.

What are the pH and pOH of a solution of 2.0 M HCl, which ionizes completely?

21.

What are the hydronium and hydroxide ion concentrations in a solution whose pH is 6.52?

22.

Calculate the hydrogen ion concentration and the hydroxide ion concentration in wine from its pH. See Figure 14.2 for useful information.

23.

Calculate the hydronium ion concentration and the hydroxide ion concentration in lime juice from its pH. See Figure 14.2 for useful information.

24.

The hydronium ion concentration in a sample of rainwater is found to be 1.7 \times 10^{−6} M at 25 °C. What is the concentration of hydroxide ions in the rainwater?

25.

The hydroxide ion concentration in household ammonia is 3.2 \times 10^{−3} M at 25 °C. What is the concentration of hydronium ions in the solution?

26.

Explain why the neutralization reaction of a strong acid and a weak base gives a weakly acidic solution.

27.

Explain why the neutralization reaction of a weak acid and a strong base gives a weakly basic solution.

28.

Use this list of important industrial compounds (and Figure 14.8) to answer the following questions regarding: Ca(OH)₂, CH₃CO₂H, HCl, H₂CO₃, HF, HNO₂, HNO₃, H₃PO₄, H₂SO₄, NH₃, NaOH, Na₂CO₃.

1. Identify the strong Brønsted-Lowry acids and strong Brønsted-Lowry bases.
2. Identify the compounds that can behave as Brønsted-Lowry acids with strengths lying between those of H₃O⁺ and H₂O.
3. Identify the compounds that can behave as Brønsted-Lowry bases with strengths lying between those of H₂O and OH⁻.

29.

The odor of vinegar is due to the presence of acetic acid, CH₃CO₂H, a weak acid. List, in order of descending concentration, all of the ionic and molecular species present in a 1-M aqueous solution of this acid.

30.

Household ammonia is a solution of the weak base NH₃ in water. List, in order of descending concentration, all of the ionic and molecular species present in a 1-M aqueous solution of this base.

31.

Explain why the ionization constant, K_a, for H₂SO₄ is larger than the ionization constant for H₂SO₃.

32.

Explain why the ionization constant, K_a, for HI is larger than the ionization constant for HF.

33.

Gastric juice, the digestive fluid produced in the stomach, contains hydrochloric acid, HCl. Milk of Magnesia, a suspension of solid Mg(OH)₂ in an aqueous medium, is sometimes used to neutralize excess stomach acid. Write a complete balanced equation for the neutralization reaction, and identify the conjugate acid-base pairs.

34.

Nitric acid reacts with insoluble copper(II) oxide to form soluble copper(II) nitrate, Cu(NO₃)₂, a compound that has been used to prevent the growth of algae in swimming pools. Write the balanced chemical equation for the reaction of an aqueous solution of HNO₃ with CuO.

35.

What is the ionization constant at 25 °C for the weak acid the conjugate acid of the weak base CH₃NH₂, K_b = 4.4 \times 10^{−4}.

36.

What is the ionization constant at 25 °C for the weak acid the conjugate acid of the weak base (CH₃)₂NH, K_b = 5.9 \times 10^{−4}?

37.

Which base, CH₃NH₂ or (CH₃)₂NH, is the stronger base? Which conjugate acid, or , is the stronger acid?

38.

Which is the stronger acid, or HBrO?

39.

Which is the stronger base, (CH₃)₃N or

40.

Predict which acid in each of the following pairs is the stronger and explain your reasoning for each.

1. H₂O or HF
2. B(OH)₃ or Al(OH)₃
3. or
4. NH₃ or H₂S
5. H₂O or H₂Te

41.

Predict which compound in each of the following pairs of compounds is more acidic and explain your reasoning for each.

1. or
2. NH₃ or H₂O
3. PH₃ or HI
4. NH₃ or PH₃
5. H₂S or HBr

42.

Rank the compounds in each of the following groups in order of increasing acidity or basicity, as indicated, and explain the order you assign.

1. acidity: HCl, HBr, HI
2. basicity: H₂O, OH⁻, H⁻, Cl⁻
3. basicity: Mg(OH)₂, Si(OH)₄, ClO₃(OH) (Hint: Formula could also be written as HClO₄.)
4. acidity: HF, H₂O, NH₃, CH₄

43.

Rank the compounds in each of the following groups in order of increasing acidity or basicity, as indicated, and explain the order you assign.

1. (f) basicity: BrO⁻,

44.

Both HF and HCN ionize in water to a limited extent. Which of the conjugate bases, F⁻ or CN⁻, is the stronger base?

45.

The active ingredient formed by aspirin in the body is salicylic acid, C₆H₄OH(CO₂H). The carboxyl group (−CO₂H) acts as a weak acid. The phenol group (an OH group bonded to an aromatic ring) also acts as an acid but a much weaker acid. List, in order of descending concentration, all of the ionic and molecular species present in a 0.001-M aqueous solution of C₆H₄OH(CO₂H).

46.

Are the concentrations of hydronium ion and hydroxide ion in a solution of an acid or a base in water directly proportional or inversely proportional? Explain your answer.

47.

What two common assumptions can simplify calculation of equilibrium concentrations in a solution of a weak acid or base?

48.

Which of the following will increase the percent of NH₃ that is converted to the ammonium ion in water?

1. addition of NaOH
2. addition of HCl
3. addition of NH₄Cl

49.

Which of the following will increase the percentage of HF that is converted to the fluoride ion in water?

1. addition of NaOH
2. addition of HCl
3. addition of NaF

50.

What is the effect on the concentrations of HNO₂, and OH⁻ when the following are added to a solution of KNO₂ in water:

1. HCl
2. HNO₂
3. NaOH
4. NaCl
5. KNO

51.

What is the effect on the concentration of hydrofluoric acid, hydronium ion, and fluoride ion when the following are added to separate solutions of hydrofluoric acid?

1. HCl
2. KF
3. NaCl
4. KOH
5. HF

52.

Why is the hydronium ion concentration in a solution that is 0.10 M in HCl and 0.10 M in HCOOH determined by the concentration of HCl?

53.

From the equilibrium concentrations given, calculate K_a for each of the weak acids and K_b for each of the weak bases.

(a) CH₃CO₂H: = 1.34 \times 10^{−3} M;
= 1.34 \times 10^{−3} M;

[CH₃CO₂H] = 9.866 \times 10^{−2} M;

(b) ClO⁻: [OH⁻] = 4.0 \times 10^{−4} M;

[HClO] = 2.38 \times 10^{−4} M;

[ClO⁻] = 0.273 M;

(c) HCO₂H: [HCO₂H] = 0.524 M;
= 9.8 \times 10^{−3} M;
= 9.8 \times 10^{−3} M;

(d) = 0.233 M;

[C₆H₅NH₂] = 2.3 \times 10^{−3} M;
= 2.3 \times 10^{−3} M

54.

From the equilibrium concentrations given, calculate K_a for each of the weak acids and K_b for each of the weak bases.

(a) NH₃: [OH⁻] = 3.1 \times 10^{−3} M;
= 3.1 \times 10^{−3} M;

[NH₃] = 0.533 M;

(b) HNO₂: = 0.011 M;
= 0.0438 M;

[HNO₂] = 1.07 M;

(c) (CH₃)₃N: [(CH₃)₃N] = 0.25 M;
[(CH₃)₃NH⁺] = 4.3 \times 10^{−3} M;

[OH⁻] = 3.7 \times 10^{−3} M;

(d) = 0.100 M;

[NH₃] = 7.5 \times 10^{−6} M;
[H₃O⁺] = 7.5 \times 10^{−6} M

55.

Determine K_b for the nitrite ion, In a 0.10-M solution this base is 0.0015% ionized.

56.

Determine K_a for hydrogen sulfate ion, In a 0.10-M solution the acid is 29% ionized.

57.

Calculate the ionization constant for each of the following acids or bases from the ionization constant of its conjugate base or conjugate acid:

1. (f) (as a base)

58.

Calculate the ionization constant for each of the following acids or bases from the ionization constant of its conjugate base or conjugate acid:

1. (f) (as a base)

59.

Using the K_a value of 1.4

60.

Calculate the concentration of all solute species in each of the following solutions of acids or bases. Assume that the ionization of water can be neglected, and show that the change in the initial concentrations can be neglected.

1. 0.0092 M HClO, a weak acid
2. 0.0784 M C₆H₅NH₂, a weak base
3. 0.0810 M HCN, a weak acid
4. 0.11 M (CH₃)₃N, a weak base
5. 0.120 M a weak acid, K_a = 1.6 \times 10^{−7}

61.

Propionic acid, C₂H₅CO₂H (K_a = 1.34 \times 10^{−5}), is used in the manufacture of calcium propionate, a food preservative. What is the pH of a 0.698-M solution of C₂H₅CO₂H?

62.

White vinegar is a 5.0% by mass solution of acetic acid in water. If the density of white vinegar is 1.007 g/cm³, what is the pH?

63.

The ionization constant of lactic acid, CH₃CH(OH)CO₂H, an acid found in the blood after strenuous exercise, is 1.36 \times 10^{−4}. If 20.0 g of lactic acid is used to make a solution with a volume of 1.00 L, what is the concentration of hydronium ion in the solution?

64.

Nicotine, C₁₀H₁₄N₂, is a base that will accept two protons (K_b1 = 7 \times 10^{−7}, K_b2 = 1.4 \times 10^{−11}). What is the concentration of each species present in a 0.050-M solution of nicotine?

65.

The pH of a 0.23-M solution of HF is 1.92. Determine K_a for HF from these data.

66.

The pH of a 0.15-M solution of is 1.43. Determine K_a for from these data.

67.

The pH of a 0.10-M solution of caffeine is 11.70. Determine K_b for caffeine from these data:

68.

The pH of a solution of household ammonia, a 0.950 M solution of NH_3, is 11.612. Determine K_b for NH₃ from these data.

69.

Determine whether aqueous solutions of the following salts are acidic, basic, or neutral:

1. Al(NO₃)₃
2. RbI
3. KHCO₂
4. CH₃NH₃Br

70.

Determine whether aqueous solutions of the following salts are acidic, basic, or neutral:

1. FeCl₃
2. K₂CO₃
3. NH₄Br
4. KClO₄

71.

Novocaine, C₁₃H₂₁O₂N₂Cl, is the salt of the base procaine and hydrochloric acid. The ionization constant for procaine is 7 \times 10^{−6}. Is a solution of novocaine acidic or basic? What are [H₃O⁺], [OH⁻], and pH of a 2.0% solution by mass of novocaine, assuming that the density of the solution is 1.0 g/mL.

72.

Which of the following concentrations would be practically equal in a calculation of the equilibrium concentrations in a 0.134-M solution of H₂CO₃, a diprotic acid: [OH⁻], [H₂CO₃], No calculations are needed to answer this question.

73.

Calculate the concentration of each species present in a 0.050-M solution of H₂S.

74.

Calculate the concentration of each species present in a 0.010-M solution of phthalic acid, C₆H₄(CO₂H)₂.

75.

Salicylic acid, HOC₆H₄CO₂H, and its derivatives have been used as pain relievers for a long time. Salicylic acid occurs in small amounts in the leaves, bark, and roots of some vegetation (most notably historically in the bark of the willow tree). Extracts of these plants have been used as medications for centuries. The acid was first isolated in the laboratory in 1838.

1. Both functional groups of salicylic acid ionize in water, with K_a = 1.0 \times 10^{−3} for the—CO₂H group and 4.2 \times 10^{−13} for the −OH group. What is the pH of a saturated solution of the acid (solubility = 1.8 g/L).
2. Aspirin was discovered as a result of efforts to produce a derivative of salicylic acid that would not be irritating to the stomach lining. Aspirin is acetylsalicylic acid, CH₃CO₂C₆H₄CO₂H. The −CO₂H functional group is still present, but its acidity is reduced, K_a = 3.0 \times 10^{−4}. What is the pH of a solution of aspirin with the same concentration as a saturated solution of salicylic acid (See Part a).

76.

The ion HTe⁻ is an amphiprotic species; it can act as either an acid or a base.

1. What is K_a for the acid reaction of HTe⁻ with H₂O?
2. What is K_b for the reaction in which HTe⁻ functions as a base in water?
3. Demonstrate whether or not the second ionization of H₂Te can be neglected in the calculation of [HTe⁻] in a 0.10 M solution of H₂Te.

77.

Explain why a buffer can be prepared from a mixture of NH₄Cl and NaOH but not from NH₃ and NaOH.

78.

Explain why the pH does not change significantly when a small amount of an acid or a base is added to a solution that contains equal amounts of the acid H₃PO₄ and a salt of its conjugate base NaH₂PO₄.

79.

Explain why the pH does not change significantly when a small amount of an acid or a base is added to a solution that contains equal amounts of the base NH₃ and a salt of its conjugate acid NH₄Cl.

80.

What is [H₃O⁺] in a solution of 0.25 M CH₃CO₂H and 0.030 M NaCH₃CO₂?

81.

What is [H₃O⁺] in a solution of 0.075 M HNO₂ and 0.030 M NaNO₂?

82.

What is [OH⁻] in a solution of 0.125 M CH₃NH₂ and 0.130 M CH₃NH₃Cl?

83.

What is [OH⁻] in a solution of 1.25 M NH₃ and 0.78 M NH₄NO₃?

84.

What is the effect on the concentration of acetic acid, hydronium ion, and acetate ion when the following are added to an acidic buffer solution of equal concentrations of acetic acid and sodium acetate:

1. HCl
2. KCH₃CO₂
3. NaCl
4. KOH
5. CH₃CO₂H

85.

What is the effect on the concentration of ammonia, hydroxide ion, and ammonium ion when the following are added to a basic buffer solution of equal concentrations of ammonia and ammonium nitrate:

1. KI
2. NH₃
3. HI
4. NaOH
5. NH₄Cl

86.

What will be the pH of a buffer solution prepared from 0.20 mol NH₃, 0.40 mol NH₄NO₃, and just enough water to give 1.00 L of solution?

87.

Calculate the pH of a buffer solution prepared from 0.155 mol of phosphoric acid, 0.250 mole of KH₂PO₄, and enough water to make 0.500 L of solution.

88.

How much solid NaCH₃CO₂•3H₂O must be added to 0.300 L of a 0.50-M acetic acid solution to give a buffer with a pH of 5.00? (Hint: Assume a negligible change in volume as the solid is added.)

89.

What mass of NH₄Cl must be added to 0.750 L of a 0.100-M solution of NH₃ to give a buffer solution with a pH of 9.26? (Hint: Assume a negligible change in volume as the solid is added.)

90.

A buffer solution is prepared from equal volumes of 0.200 M acetic acid and 0.600 M sodium acetate. Use 1.80 \times 10^{−5} as K_a for acetic acid.

1. What is the pH of the solution?
2. Is the solution acidic or basic?
3. What is the pH of a solution that results when 3.00 mL of 0.034 M HCl is added to 0.200 L of the original buffer?

91.

A 5.36–g sample of NH₄Cl was added to 25.0 mL of 1.00 M NaOH and the resulting solution

diluted to 0.100 L.

1. What is the pH of this buffer solution?
2. Is the solution acidic or basic?
3. What is the pH of a solution that results when 3.00 mL of 0.034 M HCl is added to the solution?

92.

Explain how to choose the appropriate acid-base indicator for the titration of a weak base with a strong acid.

93.

Explain why an acid-base indicator changes color over a range of pH values rather than at a specific pH.

94.

Calculate the pH at the following points in a titration of 40 mL (0.040 L) of 0.100 M barbituric acid (K_a = 9.8 \times 10^{−5}) with 0.100 M KOH.

1. no KOH added
2. 20 mL of KOH solution added
3. 39 mL of KOH solution added
4. 40 mL of KOH solution added
5. 41 mL of KOH solution added

95.

The indicator dinitrophenol is an acid with a K_a of 1.1 \times 10^{−4}. In a 1.0 \times 10^{−4}-M solution, it is colorless in acid and yellow in base. Calculate the pH range over which it goes from 10% ionized (colorless) to 90% ionized (yellow).