Chapter 14: Acid-Base EquilibriaContent check — Chemistry: Atoms First, Chapter 14: Acid-Base Equilibria

(a) as a base with HBr

(b) as an acid with OH⁻

Air-saturated water has a hydronium ion concentration caused by the dissolved CO₂ of 2.0 $\times$ 10⁻⁶ M, about 20-times larger than that of pure water. Calculate the pH of the solution at 25 °C.

What is the acid ionization constant for this weak acid if an equilibrium mixture has the following composition: $[{\text{H}}_3{\text{O}}^{\text{+}}]$ = 0.027 M; $[{\text{HSO}}_4{}^{\text{−}}]=$ $0.29M;$ and $[{\text{SO}}_4{}^{\mathrm{2-}}]=$ $0.13M?$

if the composition of an equilibrium mixture is as follows: [OH⁻] = 1.3 $\times$ 10⁻⁶ M; $[{\text{H}}_2{\text{PO}}_4{}^{\text{−}}]=$ $0.042M;$ and $[{\text{HPO}}_4{}^{\mathrm{2-}}]=$ $0.341M?$

(a) K₂CO₃

(b) CaCl₂

(c) KH₂PO₄

(d) (NH₄)₂CO₃

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

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

(b) HCl

(c) NH₃

(d) CH₃CO₂H

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

(f) ${\text{HSO}}_4{}^{\text{−}}$

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

(a) H₂O

(b) OH⁻

(c) NH₃

(d) CN⁻

(e) S²⁻

(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) OH⁻

(b) H₂O

(c) ${\text{HCO}}_3{}^{\text{−}}$

(d) NH₃

(e) ${\text{HSO}}_4{}^{\text{−}}$

(f) H₂O₂

(g) HS⁻

(h) ${\text{H}}_5{\text{N}}_2{}^{\text{+}}$

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) ${\text{HNO}}_3+$ ${\text{H}}_2\text{O}⟶$ ${\text{H}}_3{\text{O}}^{\text{+}}+$ ${\text{NO}}_3{}^{\text{−}}$

(b) ${\text{CN}}^{\text{−}}+$ ${\text{H}}_2\text{O}⟶$ $\text{HCN}+$ ${\text{OH}}^{\text{−}}$

(c) ${\text{H}}_2{\text{SO}}_4+$ ${\text{Cl}}^{\text{−}}⟶$ $\text{HCl}+$ ${\text{HSO}}_4{}^{\text{−}}$

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

(e) ${\text{O}}^{\mathrm{2-}}+$ ${\text{H}}_2\text{O}⟶$ $2{\mathrm{OH}}^{\text{−}}$

(f) ${[\text{Cu}{({\text{H}}_2\text{O})}_3(\text{OH})]}^{\text{+}}+$ ${[\text{Al}{({\text{H}}_2\text{O})}_6]}^{\mathrm{3+}}⟶$ ${[\text{Cu}{({\text{H}}_2\text{O})}_4]}^{\mathrm{2+}}+$ ${[\text{Al}{({\text{H}}_2\text{O})}_5(\text{OH})]}^{\mathrm{2+}}$

(g) ${\text{H}}_2\text{S}+$ ${\text{NH}}_2{}^{\text{−}}⟶$ ${\text{HS}}^{\text{−}}+$ ${\text{NH}}_3$

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

(a) ${\text{NH}}_3$

(b) ${\text{HPO}}_4{}^{\text{2−}}$

(c) Br⁻

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

(e) ${\text{AsO}}_4{}^{\text{3−}}$

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

(a) 0.000259 M HClO₄

(b) 0.21 M NaOH

(c) 0.000071 M Ba(OH)₂

(d) 2.5 M KOH

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

(a) ${\text{HSO}}_4{}^{\text{−}}$ or ${\text{HSeO}}_4{}^{\text{−}}$

(b) NH₃ or H₂O

(c) PH₃ or HI

(d) NH₃ or PH₃

(e) H₂S or HBr

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: NaHSO₃, NaHSeO₃, NaHSO₄

(b) basicity: ${\text{BrO}}_2{}^{\text{−}},$ ${\text{ClO}}_2{}^{\text{−}},$ ${\text{IO}}_2{}^{\text{−}}$

(c) acidity: HOCl, HOBr, HOI

(d) acidity: HOCl, HOClO, HOClO₂, HOClO₃

(e) basicity: ${\text{NH}}_2{}^{\text{−}},$ HS⁻, HTe⁻, ${\text{PH}}_2{}^{\text{−}}$

(f) basicity: BrO⁻, ${\text{BrO}}_2{}^{\text{−}},$ ${\text{BrO}}_3{}^{\text{−}},$ ${\text{BrO}}_4{}^{\text{−}}$

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

(a) addition of NaOH

(b) addition of HCl

(c) addition of NH₄Cl

What is the effect on the concentrations of ${\text{NO}}_2{}^{\text{−}},$ HNO₂, and OH⁻ when the following are added to a solution of KNO₂ in water:

(a) HCl

(b) HNO₂

(c) NaOH

(d) NaCl

(e) KNO₂

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;
$[{\text{NH}}_4{}^{\text{+}}]$ = 3.1 $\times$ 10⁻³ M;

[NH₃] = 0.533 M;

(b) HNO₂: $[{\text{H}}_3{\text{O}}^{\text{+}}]$ = 0.011 M;
$[{\text{NO}}_2{}^{\text{−}}]$ = 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) ${\text{NH}}_4{}^{\text{+}}:$ $[{\text{NH}}_4{}^{\text{+}}]$ = 0.100 M;

[NH₃] = 7.5 $\times$ 10⁻⁶ M;
[H₃O⁺] = 7.5 $\times$ 10⁻⁶ M

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

(a) HTe⁻ (as a base)

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

(c) ${\text{HAsO}}_4{}^{\mathrm{2–}}$ (as a base)

(d) ${\text{HO}}_2{}^{\text{−}}$ (as a base)

(e) ${\text{C}}_6{\text{H}}_5{\text{NH}}_3{}^{\text{+}}$

(f) ${\text{HSO}}_3{}^{\text{−}}$ (as a base)

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.

(a) 0.0092 M HClO, a weak acid

(b) 0.0784 M C₆H₅NH₂, a weak base

(c) 0.0810 M HCN, a weak acid

(d) 0.11 M (CH₃)₃N, a weak base

(e) 0.120 M $\text{Fe}{({\text{H}}_2\text{O})}_6{}^{\mathrm{2+}}$ a weak acid, K_a = 1.6 $\times$ 10⁻⁷

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

(a) FeCl₃

(b) K₂CO₃

(c) NH₄Br

(d) KClO₄

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

(a) What is K_a for the acid reaction of HTe⁻ with H₂O?

(b) What is K_b for the reaction in which HTe⁻ functions as a base in water?

(c) 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.

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:

(a) HCl

(b) KCH₃CO₂

(c) NaCl

(d) KOH

(e) CH₃CO₂H

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.

(a) What is the pH of the 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 0.200 L of the original buffer?

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.

(a) no KOH added

(b) 20 mL of KOH solution added

(c) 39 mL of KOH solution added

(d) 40 mL of KOH solution added

(e) 41 mL of KOH solution added