Effects of pH on Solubility

The Effect of Acid–Base Equilibria on the Solubility of Salts

• The solubility of compounds is heavily influenced by pH, especially for salts with anions that are conjugate bases of weak acids.
• When a salt $MA$ dissolves in water:
  $MA(s) \rightleftharpoons M^+(aq) + A^-(aq)$
• The anion $A^-$ can react with water:
  $A^-(aq) + H_2O(l) \rightleftharpoons HA(aq) + OH^-(aq)$
• Adding a strong acid increases the solubility of salts with basic anions by reacting with $A^-$ to form $HA$, reducing $[A^-]$ and increasing the dissolution of $MA$.
• pH has little effect on salts whose anions are conjugate bases of strong acids (e.g., chlorides, bromides, iodides, sulfates).
• Example:
  $Mg(OH)<em>2(s) \rightleftharpoons Mg^{2+}(aq) + 2OH^-(aq)$ Adding acid: $H^+(aq) + OH^-(aq) \rightarrow H</em>2O(l)$
  Overall:
  $Mg(OH)<em>2(s) + 2H^+(aq) \rightleftharpoons Mg^{2+}(aq) + 2H</em>2O(l)$
• Similarly, for $CaF<em>2$: $CaF</em>2(s) \rightleftharpoons Ca^{2+}(aq) + 2F^-(aq)$
  Adding strong acid:
  $2H^+(aq) + 2F^-(aq) \rightleftharpoons 2HF(aq)$
  Net reaction:
  $CaF_2(s) + 2H^+(aq) \rightarrow Ca^{2+}(aq) + 2HF(aq)$
• Sparingly soluble salts of weak acids are more soluble in acidic solutions.

Solubility Products and pH

• Examples of pH-dependent solubility include limestone caves (e.g., Carlsbad Caverns).
• Reactions for limestone cave formation:
  $CO<em>2(aq) + H</em>2O(l) \rightleftharpoons H^+(aq) + HCO<em>3^-(aq)$$HCO</em>3^-(aq) \rightleftharpoons H^+(aq) + CO<em>3^{2-}(aq)$$Ca^{2+}(aq) + CO</em>3^{2-}(aq) \rightleftharpoons CaCO_3(s)$
• When $CaCO<em>3$ saturated solution rises and is heated: $Ca^{2+}(aq) + 2HCO</em>3^-(aq) \rightleftharpoons CaCO<em>3(s) + CO</em>2(g) + H_2O(l)$

Acidic, Basic, and Amphoteric Oxides and Hydroxides

• Classification of oxides and hydroxides is based on their solubility in acidic vs. basic solutions.
• Basic oxides/hydroxides react with water to produce a basic solution or dissolve in aqueous acid.
• Acidic oxides/hydroxides react with water to produce an acidic solution or dissolve in aqueous base.
• Oxides of metallic elements are generally basic; oxides of nonmetallic elements are acidic.
• Example: Cesium oxide (basic):
  $Cs<em>2O(s) + H</em>2O(l) \rightarrow 2Cs^+(aq) + 2OH^-(aq)$
• Sulfur trioxide (acidic):
  $SO<em>3(g) + H</em>2O(l) \rightarrow H<em>2SO</em>4(aq)$
• Oxides of metals in high oxidation states also tend to be acidic.
• Amphoteric oxides dissolve in both acidic and basic solutions.
• Example: Chromium(III) hydroxide:
  $Cr(OH)<em>3(s) + 3H^+(aq) \rightarrow Cr^{3+}(aq) + 3H</em>2O(l)$
  $Cr(OH)<em>3(s) + OH^-(aq) \rightarrow [Cr(OH)</em>4]^-(aq)$
• Aluminum hydroxide is also amphoteric:
  $Al(OH)<em>3(s) + 3H^+(aq) \rightarrow Al^{3+}(aq) + 3H</em>2O(l)$
  $Al(OH)<em>3(s) + OH^-(aq) \rightarrow [Al(OH)</em>4]^-(aq)$

Selective Precipitation Using pH

• Dissolved metal ions can be separated by selective precipitation by controlling pH.
• Example: Separating $Zn^{2+}$ and $Cd^{2+}$ as sulfide salts:
  $ZnS(s) \rightleftharpoons Zn^{2+}(aq) + S^{2-}(aq)$
  $CdS(s) \rightleftharpoons Cd^{2+}(aq) + S^{2-}(aq)$
• Sulfide concentrations needed for precipitation:
  [S^{2-}] > 1.6 \times 10^{-21} M for $ZnS$
  [S^{2-}] > 8.0 \times 10^{-24} M for $CdS$
• To control sulfide concentration, adjust pH by adding acid:
  $H_2S(aq) \rightleftharpoons 2H^+(aq) + S^{2-}(aq)$
• Adjusting $[H^+]$ allows selective precipitation of $CdS$ without precipitating $ZnS$.