14.4 Hydrolysis of Salts

Salt hydrolysis explains why dissolving a salt in water doesn't always give you a neutral solution. Even though salts come from acid-base neutralization reactions, the resulting solution can be acidic, basic, or neutral depending on the strengths of the parent acid and base.

When ions from a dissolved salt interact with water, they can produce or consume $H^+$ or $OH^-$ ions. This is the hydrolysis reaction, and it's the key to predicting the pH of any salt solution.

Hydrolysis of Salts

pH Prediction for Salt Solutions

The pH of a salt solution depends entirely on what kind of acid and base originally combined to form that salt. There are four cases to know:

Strong acid + strong base → neutral solution (pH = 7) Neither ion hydrolyzes. The cation (like $Na^+$ or $K^+$) and the anion (like $Cl^-$ or $NO_3^-$) are spectator ions that don't react with water. Examples: $NaCl$, $KNO_3$.

Weak acid + strong base → basic solution (pH > 7) The anion is the conjugate base of a weak acid, so it hydrolyzes by accepting a proton from water, generating $OH^-$ ions. For example, in a solution of sodium acetate ($CH_3COONa$), the acetate ion reacts with water:

$CH_3COO^- + H_2O \rightleftharpoons CH_3COOH + OH^-$

The $Na^+$ ion doesn't hydrolyze, so the net effect is a basic solution.

Strong acid + weak base → acidic solution (pH < 7) The cation is the conjugate acid of a weak base, so it hydrolyzes by donating a proton to water, generating $H_3O^+$ ions. For example, in a solution of ammonium chloride ($NH_4Cl$), the ammonium ion reacts with water:

$NH_4^+ + H_2O \rightleftharpoons NH_3 + H_3O^+$

The $Cl^-$ ion doesn't hydrolyze, so the net effect is an acidic solution.

Weak acid + weak base → depends on relative strengths Both ions hydrolyze, so you need to compare the $K_a$ of the cation with the $K_b$ of the anion:

• If $K_a > K_b$, the solution is acidic (pH < 7)
• If $K_b > K_a$, the solution is basic (pH > 7)
• If $K_a = K_b$, the solution is neutral (pH ≈ 7)

The underlying idea here is conjugate acid-base pairs. Every weak acid produces a conjugate base that can hydrolyze, and every weak base produces a conjugate acid that can hydrolyze. Strong acids and bases produce conjugates that are too weak to hydrolyze meaningfully.

Ion Concentrations in Hydrolyzed Salts

To calculate the actual pH or ion concentrations, you'll use an ICE table along with the appropriate equilibrium constant.

For salts of weak acids and strong bases (e.g., $CH_3COONa$):

1. Write the hydrolysis equation: $CH_3COO^- + H_2O \rightleftharpoons CH_3COOH + OH^-$

2. Set up an ICE table using the initial concentration of the salt as the starting concentration of the anion.

3. Use $K_b$ for the conjugate base to solve for $[OH^-]$:

$K_b = \frac{[CH_3COOH][OH^-]}{[CH_3COO^-]}$

Remember that $K_b$ for the conjugate base is related to the $K_a$ of the parent weak acid by: $K_w = K_a \times K_b$, so $K_b = \frac{K_w}{K_a}$.

For salts of strong acids and weak bases (e.g., $NH_4Cl$):

1. Write the hydrolysis equation: $NH_4^+ + H_2O \rightleftharpoons NH_3 + H_3O^+$
2. Set up an ICE table using the initial concentration of the salt as the starting concentration of the cation.
3. Use $K_a$ for the conjugate acid to solve for $[H_3O^+]$:

$K_a = \frac{[NH_3][H_3O^+]}{[NH_4^+]}$

Here, $K_a = \frac{K_w}{K_b}$, where $K_b$ is for the parent weak base ($NH_3$).

One thing to watch for: the common ion effect can shift the hydrolysis equilibrium. If a solution already contains one of the product ions (say, extra $OH^-$), the hydrolysis reaction will be suppressed.

Hydrated Metal Ions as Acids

Small, highly charged metal ions don't just float around in water. They form hydrated complexes, like $Al(H_2O)_6^{3+}$, where water molecules coordinate directly to the metal ion.

The high charge density of the metal ion pulls electron density away from the $O-H$ bonds in those coordinated water molecules. This weakens the bonds and makes the hydrogen atoms much more likely to dissociate as $H^+$. The result is a hydrolysis reaction:

$Al(H_2O)_6^{3+} + H_2O \rightleftharpoons Al(H_2O)_5OH^{2+} + H_3O^+$

This produces $H_3O^+$, making the solution acidic. That's why solutions of salts like $AlCl_3$ or $FeCl_3$ are noticeably acidic even though $HCl$ is a strong acid.

Two factors control how acidic a hydrated metal ion will be:

• Higher charge → stronger polarization → more acidic ($Al^{3+}$ is more acidic than $Mg^{2+}$)
• Smaller ionic radius → higher charge density → more acidic

This behavior is well explained by Lewis acid-base theory: the metal ion acts as a Lewis acid (electron pair acceptor), and the water molecules act as Lewis bases (electron pair donors). The strong Lewis acidity of the metal weakens the $O-H$ bonds in the coordinated water.

Equilibrium and pH Control

Salt hydrolysis is an equilibrium process, so Le Chatelier's principle applies. Adding more of a product (like $OH^-$ or $H_3O^+$) will shift the hydrolysis equilibrium back toward reactants, suppressing further hydrolysis. Increasing temperature generally favors the endothermic direction of the reaction.

This connects directly to buffer solutions. A buffer is a mixture of a weak acid and its conjugate base (or a weak base and its conjugate acid). In practice, many buffers are made by combining a weak acid or base with one of its salts. For example, mixing acetic acid with sodium acetate creates an acetate buffer that resists pH changes. The hydrolysis behavior of the salt's ions is exactly what allows the buffer to neutralize added acid or base.