8.11.A.1   pH-Dependent Solubility (Qualitative Only – No Calculations Required):

1. Le Châtelier’s Principle:

Le Châtelier’s Principle is that any dynamic equilibrium which has been disturbed by the alteration of conditions (such as concentration, pressure, or temperature) will change in an effort to counteract the disturbance and re-establish a new equilibrium.

Applied to solubility equilibria, especially when dealing with pH change, this principle can be used to predict solubility changes under acidic or basic conditions.

i. Effect of pH on Solubility – Elaborated by Le Châtelier’s Principle:

a. When a salt forms a basic anion (weak base conjugate):

Example:

$\text{Ag}_2\text{CO}_3(s) \rightleftharpoons 2\text{Ag}^+(aq) + \text{CO}_3^{2-}(aq)$

Carbonate ion $\text{CO}_3^{2-}$  is a weak base and reacts with H⁺:$$\text{CO}_3^{2-} + 2H^+ \rightarrow H_2CO_3 \rightarrow CO_2 + H_2O$$

b. In acidic solution (low pH):

* Added H⁺ combines with CO₃²⁻, lowering its concentration.
* Equilibrium moves to the right to dissolve more Ag2​CO3​.

* Solubility increases.

ii. If a salt forms an acidic cation (conjugate acid of a weak base):

Example:

$\text{Fe(OH)}_3(s) \rightleftharpoons \text{Fe}^{3+}(aq) + 3\text{OH}^-(aq)$

In acidic solution:

* Added H⁺ reacts with OH⁻ to form water.
* OH⁻ concentration drops.
* Equilibrium shifts right; more Fe(OH)₃ dissolves.
* Solubility increases.

In basic solution (high pH):

* OH⁻ increases.
* Equilibrium shifts left; precipitation occurs.
* Solubility decreases.

iii. Summary Table:

2. Weak Acids/Bases and Their Ions:

i. Salts Having the Conjugate Base of a Weak Acid:

* Example: Calcium carbonate, CaCO₃

Dissociation:$$\text{CaCO}_3(s) \rightleftharpoons \text{Ca}^{2+}(aq) + \text{CO}_3^{2-}(aq)$$

Carbonate ion (CO₃²⁻) is the conjugate base of the weak acid H₂CO₃.

ii. Low pH (acidic conditions):

* H⁺ reacts with CO₃²⁻:$$\text{CO}_3^{2-} + 2H^+ \rightarrow H_2CO_3 \rightarrow CO_2 + H_2O$$

* This removes CO₃²⁻ from the solution.
* Equilibrium shifts right, more CaCO₃ dissolves.
* Solubility increases.

iii. Salts of the Conjugate Acid of a Weak Base:

* Example: Ammonium chloride, NH₄Cl

Dissociation:$$\text{NH}_4\text{Cl} \rightarrow \text{NH}_4^+ + \text{Cl}^-$$

Ammonium ion (NH₄⁺) is the conjugate acid of the weak base NH₃.

* Effect of high pH (basic conditions):

* OH⁻ reacts with NH₄⁺:$$\text{NH}_4^+ + OH^- \rightarrow NH_3 + H_2O$$

This takes NH₄⁺ out of the solution.
* Equilibrium shifts to the right, more NH₄Cl dissolves.
* Solubility increases in basic solution.

iv. Quick Reference Table:

3. Common Ion Effect:

The common ion effect is the addition of an ion already in the equilibrium to the solution, which alters the equilibrium and the solubility.

Key Concept:

* Adding H⁺ or OH⁻ brings in a common ion that can shift the dissolution equilibrium.

i. Examples:

a. Salt with OH⁻ (e.g., Mg(OH)₂):$${\text{Mg(OH)}}_2\rightleftharpoons {\text{Mg}}^{2+}+2{\text{OH}}^{-}$$

* Add OH⁻ (↑pH) → displaces equilibrium left
* Solubility decreases

b. Salt with conjugate base (e.g., NaF):$$\text{NaF}\rightleftharpoons {\text{Na}}^{+}+{\text{F}}^{-}$$

* Add H⁺ (↓pH) → H⁺ reacts with F⁻
* Equilibrium shifts right
* Solubility increase

Summary Table:

4. Hydroxide Ion and Metal Hydroxides:

Metal hydroxide solubility in acid increases because added H⁺ reacts with OH⁻, which shifts the equilibrium.

Example:

$${\text{Fe(OH)}}_3(s)\rightleftharpoons {\text{Fe}}^{3+}(aq)+3{\text{OH}}^{-}(aq)$$

* Add H⁺ (↓pH) →

H⁺ + OH⁻ → H₂O
* Removes OH⁻ from solution
* Equilibrium shifts right → more dissolves

Key Point:

> Lower pH (more acidic) → less OH⁻ → greater metal hydroxide solubility

Applies to:

* Fe(OH)₃, Al(OH)₃, Mg(OH)₂, etc.