Lesson 8.6: Acid–Base Properties of Salt Solutions

Acid–Base Properties of Salt Solutions

Introduction to Salt Solutions

Salts are ionic compounds composed of cations and anions arranged in a crystalline lattice. When dissolved in water, salts dissociate into their constituent ions, which may or may not affect the pH of the solution. The acid–base properties of salt solutions depend on the nature of the ions produced upon dissolution.

• Electrolyte: A substance that dissociates into ions in water, conducting electricity.

• Acidic salt: Produces hydrogen ions (H+) in water, lowering pH. (acidic: pH < 7)

• Basic salt: Produces hydroxide ions (OH–) in water, raising pH. (basic: pH > 7)

• Neutral salt: Does not affect the concentration of H+ or OH– in water.

Salts That Produce Neutral Solutions

Salts formed from the cation of a strong base and the anion of a strong acid do not hydrolyze in water, resulting in a neutral solution (pH = 7).

• Example: Sodium chloride (NaCl) and potassium chloride (KCl)

• The cation (e.g., Na+, K+) and anion (e.g., Cl–, NO3–) do not react with water.

Equation:

$\mathrm{KCl(s) \rightarrow K^+(aq) + Cl^-(aq)}$

Neither ion reacts with water, so the solution remains neutral.

Salts That Produce Basic Solutions

Salts containing the cation of a strong base and the anion of a weak acid produce basic solutions. The anion hydrolyzes with water to generate OH– ions.

• Example: Potassium ethanoate (KC2H3O2)

• The ethanoate ion (C2H3O2–) is the conjugate base of a weak acid (acetic acid) and reacts with water.

Hydrolysis Reaction:

$\mathrm{C_2H_3O_2^-(aq) + H_2O(l) \rightleftharpoons HC_2H_3O_2(aq) + OH^-(aq)}$

This reaction increases [OH–], making the solution basic.

Calculating the pH of a Basic Salt Solution (Better procedure is in the handout package notes)

1. Identify major entities in solution (e.g., NaNO2(aq) → Na+, NO2–).

2. Determine which ions hydrolyze (NO2– is the conjugate base of a weak acid, so it hydrolyzes).

3. Write the hydrolysis equation: $\mathrm{NO_2^-(aq) + H_2O(l) \rightleftharpoons HNO_2(aq) + OH^-(aq)}$

4. Write the equilibrium constant expression: $K_b = \frac{[HNO_2][OH^-]}{[NO_2^-]}$

5. Calculate $K_b$ using $K_w$ and $K_a$: $K_w = K_a K_b \implies K_b = \frac{K_w}{K_a}$

6. Set up an ICE table and solve for [OH–], then calculate pOH and pH.

Example Calculation: For 0.30 mol/L NaNO2(aq), $K_a$ for HNO2 is $4.6 \times 10^{-4}$, $K_w = 1.0 \times 10^{-14}$.

• $K_b = \frac{1.0 \times 10^{-14}}{4.6 \times 10^{-4}} = 2.17 \times 10^{-11}$

• Solve for $x$ in $K_b = \frac{x^2}{0.30}$, $x = [OH^-] = 2.55 \times 10^{-6}$

• $pOH = -\log(2.55 \times 10^{-6}) = 5.59$

• $pH = 14.00 - 5.59 = 8.41$ (basic)

Salts That Produce Acidic Solutions

Salts containing the cation of a weak base and the anion of a strong acid produce acidic solutions. The cation hydrolyzes with water to generate H+ ions.

• Example: Ammonium chloride (NH4Cl)

• The ammonium ion (NH4+) is the conjugate acid of a weak base (NH3) and reacts with water.

Hydrolysis Reaction:

$\mathrm{NH_4^+(aq) + H_2O(l) \rightleftharpoons NH_3(aq) + H_3O^+(aq)}$

This reaction increases [H+], making the solution acidic.

Calculating the pH of an Acidic Salt Solution

1. Identify major entities in solution (e.g., NH4Cl(aq) → NH4+, Cl–).

2. Determine which ions hydrolyze (NH4+ hydrolyzes; Cl– does not).

3. Write the hydrolysis equation: $\mathrm{NH_4^+(aq) + H_2O(l) \rightleftharpoons NH_3(aq) + H_3O^+(aq)}$

4. Write the equilibrium constant expression: $K_a = \frac{[NH_3][H^+]}{[NH_4^+]}$

5. Calculate $K_a$ using $K_w$ and $K_b$: $K_a = \frac{K_w}{K_b}$

6. Set up an ICE table and solve for [H+], then calculate pH.

Example Calculation: For 0.10 mol/L NH4Cl(aq), $K_b$ for NH3 is $1.8 \times 10^{-5}$, $K_w = 1.0 \times 10^{-14}$.

• $K_a = \frac{1.0 \times 10^{-14}}{1.8 \times 10^{-5}} = 5.56 \times 10^{-10}$

• Solve for $x$ in $K_a = \frac{x^2}{0.10}$, $x = [H^+] = 7.45 \times 10^{-6}$

• $pH = -\log(7.45 \times 10^{-6}) = 5.13$ (acidic)

Salts with Both Acidic and Basic Ions

Some salts contain both a cation that can act as an acid and an anion that can act as a base (e.g., NH4C2H3O2). The resulting solution can be acidic, basic, or neutral depending on the relative strengths (Ka and Kb) of the ions.

• If Ka > Kb: Solution is acidic.

• If Kb > Ka: Solution is basic.

• If Ka = Kb: Solution is neutral.

Hydrolysis of Amphiprotic Ions

Amphiprotic ions (e.g., HSO4–) can act as either acids or bases. The dominant reaction (acidic or basic) depends on the relative values of Ka and Kb.

• Example: Sodium hydrogen sulfate (NaHSO4)

• Acid hydrolysis: $\mathrm{HSO_4^-(aq) + H_2O(l) \rightleftharpoons H_3O^+(aq) + SO_4^{2-}(aq)}$

• Base hydrolysis: $\mathrm{HSO_4^-(aq) + H_2O(l) \rightleftharpoons OH^-(aq) + H_2SO_4(aq)}$

• Compare Ka and Kb to predict the dominant effect.

Hydrolysis of Metallic and Non-metallic Oxides

Oxides can also affect the pH of aqueous solutions:

• Metallic oxides (e.g., CaO, MgO) react with water to form basic solutions:

$\mathrm{CaO(s) + H_2O(l) \rightarrow Ca^{2+}(aq) + 2OH^-(aq)}$

• Non-metallic oxides (e.g., CO2, SO2) react with water to form acidic solutions:

$\mathrm{CO_2(g) + 2H_2O(l) \rightarrow H_3O^+(aq) + HCO_3^-(aq)}$

Metallic oxides are called basic oxides, while non-metallic oxides are called acidic oxides.

Summary Table: Acid–Base Properties of Various Types of Salts

Type of Salt	Examples	Comment	pH of Solution
Cation of Group 1 or 2 (except Be); anion from strong acid	KCl, NaCl, NaNO3	Neither ion acts as acid or base	Neutral
Cation of Group 1 or 2 (except Be); anion from weak acid	NaC2H3O2, KCN, NaF	Anion acts as base; cation has no effect	Basic
Cation is conjugate acid of weak base; anion from strong acid	NH4Cl, NH4NO3	Cation acts as acid; anion has no effect	Acidic
Cation is conjugate acid of weak base; anion is conjugate base of weak acid	NH4C2H3O2, NH4CN	Cation acts as acid; anion acts as base	Acidic if Ka > Kb Basic if Kb > Ka Neutral if Ka = Kb
Cation is highly charged metal ion; anion from strong acid	Al(NO3)3, FeCl3	Hydrated cation acts as acid; anion has no effect	Acidic

Summary Table: Acid Ionization Constants for Some Metal Cations

Metal Cation	Ka
Zr4+	2.1
Sn2+	2.0 × 10–2
Fe3+	1.5 × 10–3
Cr3+	1.0 × 10–4
Al3+	9.8 × 10–6
Be2+	3.2 × 10–7
Fe2+	1.8 × 10–7
Pb2+	1.6 × 10–8
Cu2+	1.0 × 10–8

Additional info: These values refer to hydrated metal ions, which are weak acids in water.

Key Concepts and Summary

• Salts from strong acid + strong base: neutral solution.

• Salts from strong base + weak acid: basic solution (anion hydrolyzes).

• Salts from weak base + strong acid: acidic solution (cation hydrolyzes).

• Salts with both hydrolyzing ions: compare Ka and Kb to predict pH.

• Metallic oxides form basic solutions; non-metallic oxides form acidic solutions.

Practice Problems

1. Calculate the pH of a 0.35 mol/L solution of sodium methanoate, NaCHO2(aq). [Answer: 8.64]

2. Calculate the pH of a 0.85 mol/L solution of sodium hypochlorite, NaClO(aq). [Answer: 10.69]

3. Calculate the pH of a 0.525 mol/L solution of ammonium chloride, NH4Cl(aq). (Kb for NH3 is 1.8 × 10–5). [Answer: 4.77]

4. What is the pH of a 0.0250 mol/L solution of codeine hydrochloride, DHCl(aq), if Kb for codeine is 1.62 × 10–6? [Answer: 4.91]

Glossary

• Hydrolysis: The reaction of an ion with water to produce acidic or basic solutions by generating H3O+ or OH– ions.

• Amphiprotic: A species that can act as either an acid or a base.

• Basic oxide: An oxide that reacts with water to form a basic solution (usually metallic oxides).

• Acidic oxide: An oxide that reacts with water to form an acidic solution (usually non-metallic oxides).