Exercises

Write equations that show ${\text{H}}_2{\text{PO}}_4{}^{\text{−}}$ acting both as an acid and as a base.

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

(a) HNO₃

(b) ${\text{PH}}_4{}^{\text{+}}$

(c) H₂S

(d) CH₃CH₂COOH

(e) ${\text{H}}_2{\text{PO}}_4{}^{\text{−}}$

(f) HS⁻

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

(a) HS⁻

(b) ${\text{PO}}_4{}^{\text{3−}}$

(c) ${\text{NH}}_2{}^{\text{−}}$

(d) C₂H₅OH

(e) O²⁻

(f) ${\text{H}}_2{\text{PO}}_4{}^{\text{−}}$

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

(a) H₂S

(b) ${\text{H}}_2{\text{PO}}_4{}^{\text{−}}$

(c) PH₃

(d) HS⁻

(e) ${\text{HSO}}_3{}^{\text{−}}$

(f) ${\text{H}}_3{\text{O}}_2{}^{\text{+}}$

(g) H₄N₂

(h) CH₃OH

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

(a) ${\text{NO}}_2{}^{\text{−}}+{\text{H}}_2\text{O}⟶{\text{HNO}}_2+{\text{OH}}^{\text{−}}$

(b) $\text{HBr}+{\text{H}}_2\text{O}⟶{\text{H}}_3{\text{O}}^{\text{+}}+{\text{Br}}^{\text{−}}$

(c) ${\text{HS}}^{\text{−}}+{\text{H}}_2\text{O}⟶{\text{H}}_2\text{S}+{\text{OH}}^{\text{−}}$

(d) ${\text{H}}_2{\text{PO}}_4{}^{\text{−}}+{\text{OH}}^{\text{−}}⟶{\text{HPO}}_4{}^{\text{2−}}+{\text{H}}_2\text{O}$

(e) ${\text{H}}_2{\text{PO}}_4{}^{\text{−}}+\text{HCl}⟶{\text{H}}_3{\text{PO}}_4+{\text{Cl}}^{\text{−}}$

(f) ${[\text{Fe}{({\text{H}}_2\text{O})}_5(\text{OH})]}^{\mathrm{2+}}+{[\text{Al}{({\text{H}}_2\text{O})}_6]}^{\mathrm{3+}}⟶{\text{[Fe}{({\text{H}}_2\text{O})}_6]}^{\mathrm{3+}}+{[\text{Al}{({\text{H}}_2\text{O})}_5(\text{OH})]}^{\mathrm{2+}}$

(g) ${\text{CH}}_3\text{OH}+{\text{H}}^{\text{−}}⟶{\text{CH}}_3{\text{O}}^{\text{−}}+{\text{H}}_2$

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

(a) H₂O

(b) ${\text{H}}_2{\text{PO}}_4{}^{\text{−}}$

(c) S²⁻

(d) ${\text{CO}}_3{}^{\text{2−}}$

(e) ${\text{HSO}}_4{}^{\text{−}}$

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.

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.

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

(a) 0.200 M HCl

(b) 0.0143 M NaOH

(c) 3.0 M HNO₃

(d) 0.0031 M Ca(OH)₂

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

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

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?

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

Use this list of important industrial compounds (and Figure 14.8) to answer the following questions regarding: CaO, Ca(OH)₂, CH₃CO₂H, CO_2, HCl, H₂CO₃, HF, HNO₂, HNO₃, H₃PO₄, H₂SO₄, NH₃, NaOH, Na₂CO₃.

(a) Identify the strong Brønsted-Lowry acids and strong Brønsted-Lowry bases.

(b) List those compounds in (a) that can behave as Brønsted-Lowry acids with strengths lying between those of H₃O⁺ and H₂O.

(c) List those compounds in (a) that can behave as Brønsted-Lowry bases with strengths lying between those of H₂O and OH⁻.

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.

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

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.

What is the ionization constant at 25 °C for the weak acid ${({\text{CH}}_3)}_2{\text{NH}}_2{}^{\text{+}},$ the conjugate acid of the weak base (CH₃)₂NH, K_b = 5.9 $\times$ 10⁻⁴?

Which is the stronger acid, ${\text{NH}}_4{}^{\text{+}}$ or HBrO?

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

(a) H₂O or HF

(b) B(OH)₃ or Al(OH)₃

(c) ${\text{HSO}}_3{}^{\text{−}}$ or ${\text{HSO}}_4{}^{\text{−}}$

(d) NH₃ or H₂S

(e) H₂O or H₂Te

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

(a) acidity: HCl, HBr, HI

(b) basicity: H₂O, OH⁻, H⁻, Cl⁻

(c) basicity: Mg(OH)₂, Si(OH)₄, ClO₃(OH) (Hint: Formula could also be written as HClO₄).

(d) acidity: HF, H₂O, NH₃, CH₄

Both HF and HCN ionize in water to a limited extent. Which of the conjugate bases, F⁻ or CN⁻, is the stronger base? See Table 14.2.

What do we represent when we write:
${\text{CH}}_3{\text{CO}}_2\text{H}(aq)+{\text{H}}_2\text{O}(l)\rightleftharpoons {\text{H}}_3{\text{O}}^{\text{+}}(aq)+{\text{CH}}_3{\text{CO}}_2{}^{\text{−}}(aq)?$

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.

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

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

(a) addition of NaOH

(b) addition of HCl

(c) addition of NaF

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?

(a) HCl

(b) KF

(c) NaCl

(d) KOH

(e) HF

The equation for the equilibrium is:
$\text{HF}(aq)+{\text{H}}_2\text{O}(l)\rightleftharpoons {\text{H}}_3{\text{O}}^{\text{+}}(aq)+{\text{F}}^{\text{−}}(aq)$

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: $[{\text{H}}_3{\text{O}}^{\text{+}}]$ = 1.34 $\times$ 10⁻³ M;
$[{\text{CH}}_3{\text{CO}}_2{}^{\text{−}}]$ = 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 $\times$ 10⁻⁵ M;

[ClO⁻] = 0.273 M;

(c) HCO₂H: [HCO₂H] = 0.524 M;
$[{\text{H}}_3{\text{O}}^{\text{+}}]$ = 9.8 $\times$ 10⁻³ M;
$[{\text{HCO}}_2{}^{\text{−}}]$ = 9.8 $\times$ 10⁻³ M;

(d) ${\text{C}}_6{\text{H}}_5{\text{NH}}_3{}^{\text{+}}:$ $[{\text{C}}_6{\text{H}}_5{\text{NH}}_3{}^{\text{+}}]$ = 0.233 M;

[C₆H₅NH₂] = 2.3 $\times$ 10⁻³ M;
$[{\text{H}}_3{\text{O}}^{\text{+}}]$ = 2.3 $\times$ 10⁻³ M

Determine K_b for the nitrite ion, ${\text{NO}}_2{}^{\text{−}}.$ In a 0.10-M solution this base is 0.0015% ionized.

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

(a) F⁻

(b) ${\text{NH}}_4{}^{\text{+}}$

(c) ${\text{AsO}}_4{}^{\mathrm{3-}}$

(d) ${({\text{CH}}_3)}_2{\text{NH}}_2{}^{\text{+}}$

(e) ${\text{NO}}_2{}^{\text{−}}$

(f) ${\text{HC}}_2{\text{O}}_4{}^{\text{−}}$ (as a base)

For which of the following solutions must we consider the ionization of water when calculating the pH or pOH?

(a) 3 $\times$ 10⁻⁸ M HNO₃

(b) 0.10 g HCl in 1.0 L of solution

(c) 0.00080 g NaOH in 0.50 L of solution

(d) 1 $\times$ 10⁻⁷ M Ca(OH)₂

(e) 0.0245 M KNO₃

Calculate the equilibrium concentration of the nonionized acids and all ions in a solution that is 0.25 M in HCO₂H and 0.10 M in HClO.

Calculate the equilibrium concentration of the nonionized bases and all ions in a solution that is 0.25 M in CH₃NH₂ and 0.10 M in C₅H₅N (K_b = 1.7 $\times$ 10⁻⁹).

Using the K_a value of 1.4 $\times$ 10⁻⁵, place $\text{Al}{({\text{H}}_2\text{O})}_6{}^{\mathrm{3+}}$ in the correct location in Figure 14.8.

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

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?

The pH of a 0.20-M solution of HF is 1.92. Determine K_a for HF from these data.

The pH of a 0.10-M solution of caffeine is 11.16. Determine K_b for caffeine from these data:
${\text{C}}_8{\text{H}}_{10}{\text{N}}_4{\text{O}}_2(aq)+{\text{H}}_2\text{O}(l)\rightleftharpoons {\text{C}}_8{\text{H}}_{10}{\text{N}}_4{\text{O}}_2{\text{H}}^{\text{+}}(aq)+{\text{OH}}^{\text{−}}(aq)$

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

(a) Al(NO₃)₃

(b) RbI

(c) KHCO₂

(d) CH₃NH₃Br

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.

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

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.

(a) 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).

(b) 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).

(c) Under some conditions, aspirin reacts with water and forms a solution of salicylic acid and acetic acid:
${\text{CH}}_3{\text{CO}}_2{\text{C}}_6{\text{H}}_4{\text{CO}}_2\text{H(}aq)+{\text{H}}_2\text{O}(l)⟶{\text{HOC}}_6{\text{H}}_4{\text{CO}}_2\text{H}(aq)+{\text{CH}}_3{\text{CO}}_2\text{H}(aq)$

i. Which of the acids, salicylic acid or acetic acid, produces more hydronium ions in such a solution?

ii. What are the concentrations of molecules and ions in a solution produced by the hydrolysis of 0.50 g of aspirin dissolved in enough water to give 75 mL of solution?

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

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.

What is [H₃O⁺] in a solution of 0.075 M HNO₂ and 0.030 M NaNO₂?
${\text{HNO}}_2(aq)+{\text{H}}_2\text{O}(l)\rightleftharpoons {\text{H}}_3{\text{O}}^{\text{+}}(aq)+{\text{NO}}_2{}^{\text{−}}(aq)K_{\text{a}}=4.5\times {10}^{\mathrm{-5}}$

What is [OH⁻] in a solution of 1.25 M NH₃ and 0.78 M NH₄NO₃?
${\text{NH}}_3(aq)+{\text{H}}_2\text{O}(l)\rightleftharpoons {\text{NH}}_4{}^{\text{+}}(aq)+{\text{OH}}^{\text{−}}(aq)K_{\text{b}}=1.8\times {10}^{\mathrm{-5}}$

What concentration of NaF is required to make [H₃O⁺] = 2.3 $\times$ 10⁻⁴ in a 0.300-M solution of HF?

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:

(a) KI

(b) NH₃

(c) HI

(d) NaOH

(e) NH₄Cl

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.

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.)

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.

(a) What is the pH of this buffer solution?

(b) Is the solution acidic or basic?

(c) What is the pH of a solution that results when 3.00 mL of 0.034 M HCl is added to the solution?

Which acid in Table 14.2 is most appropriate for preparation of a buffer solution with a pH of 3.7? Explain your choice.

Which base in Table 14.3 is most appropriate for preparation of a buffer solution with a pH of 9.20? Explain your choice.

What is the pH of 1.000 L of a solution of 100.0 g of glutamic acid (C₅H₉NO₄, a diprotic acid; K₁ = 8.5 $\times$ 10⁻⁵, K₂ = 3.39 $\times$ 10⁻¹⁰) to which has been added 20.0 g of NaOH during the preparation of monosodium glutamate, the flavoring agent? What is the pH when exactly 1 mol of NaOH per mole of acid has been added?

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

Why can we ignore the contribution of water to the concentration of OH⁻ in a solution of the following bases:

0.0784 M C₆H₅NH₂, a weak base

0.11 M (CH₃)₃N, a weak base

but not the contribution of water to the concentration of H₃O⁺?

Draw a curve similar to that shown in Figure 14.23 for a series of solutions of NH₃. Plot [OH⁻] on the vertical axis and the total concentration of NH₃ (both ionized and nonionized NH₃ molecules) on the horizontal axis. Let the total concentration of NH₃ vary from 1 $\times$ 10⁻¹⁰ M to 1 $\times$ 10⁻² M.

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).