Grade 11

Hydrolysis of salts

Predict whether a salt solution is acidic, basic or nearly neutral by asking whether its ions are conjugate acids or bases that react with water. Quantify the effect with hydrolysis constants, equilibrium calculations and examples.

IntuitionThe ions can change water

Dissolving a salt separates its ions. Some simply remain solvated; others exchange a proton with water, producing extra hydronium or hydroxide. That reaction—salt hydrolysis—explains why equal concentrations of different salts can have different pH values.

Select NaCl, NH₄Cl, CH₃COONa or CH₃COONH₄ in the scheme. The flask gives an indicative universal-indicator colour; boxes identify the reacting ion and spectator. The second display plots the general weak-acid/strong-base titration curve, whose equivalence pH is above 7 because the conjugate base hydrolyses.

SchoolClassifying salts

A salt of a strong acid and strong base is approximately neutral (NaCl). A strong acid/weak base salt is acidic (NH₄Cl); a weak acid/strong base salt is basic (CH₃COONa). If both parent acid and base are weak, compare their Ka and Kb: the larger hydrolysis tendency determines the pH.

Carbonate hydrolysis occurs stepwise; the first step is appreciably more basic. A 0.10 M Na₂CO₃ solution is typically about pH 11.6 (approximate, at 25 °C). Hydrated Al³⁺ and Fe³⁺ also acidify water: coordinated water molecules lose protons, for example $\ce{[M(H2O)6]^3+ + H2O <=> [M(H2O)5(OH)]^2+ + H3O+}$.

For a dilute weak-base anion salt of concentration C, $[\ce{OH-}]\approx\sqrt{K_hC}$ and $pH\approx7+\tfrac12(pK_a+\log C)$. For a weak-acid cation salt, $[\ce{H3O+}]\approx\sqrt{K_hC}$ and $pH\approx7-\tfrac12(pK_b+\log C)=\tfrac12(pK_a-\log C)$ at 25 °C. These approximations assume small hydrolysis and ideal dilute solutions.

For a weak–weak salt, both ions hydrolyse. Ammonium acetate is approximately neutral when Ka(CH₃COOH) and Kb(NH₃) are equal: $pH=\tfrac12(pK_w+pK_a-pK_b)=7.00$ at 25 °C for pKa=pKb=4.76. If they differ, use the expression rather than assuming pH 7.

Some ion combinations drive hydrolysis forward because a weakly soluble hydroxide precipitates and a gas escapes. Aluminium sulfide decomposes in water, and mixing Al³⁺ with carbonate can yield Al(OH)₃ and CO₂ rather than a stable aluminium carbonate.

Applications include baking soda’s mildly basic solution, alkaline soap solutions and alum in water treatment: hydrolysed aluminium species form flocs that capture suspended particles. Soil pH can also be affected by salt hydrolysis and metal ions.

UndergraduateExact equilibria and controlling factors

For $\ce{A- + H2O <=> HA + OH-}$, let initial salt concentration be C and hydrolysed amount x. Then $K_h=x^2/(C-x)$, so $x=[\ce{OH-}]=\frac{-K_h+\sqrt{K_h^2+4K_hC}}{2}$. For $\ce{BH+}$ the same expression gives $[\ce{H3O+}]$. The small-x approximation is valid when x/C is small; otherwise use the quadratic.

The hydrolysis degree is $h=x/C$. In the weak-hydrolysis regime $h\approx\sqrt{K_h/C}$, so dilution increases h even though the absolute ion concentration and pH shift toward neutral water. Hydrolysis is generally endothermic; increasing temperature often increases its extent, while the neutral pH itself also changes because $K_w$ increases.

At comparable coordination, higher charge and smaller ionic radius increase charge density and polarise O–H bonds in coordinated water, making proton loss easier and pKa lower. These quoted values are approximate and depend on ionic strength, temperature and conventions.