Salt Hydrolysis

When a salt dissolves in water, it dissociates into its constituent ions. These ions can interact with water molecules in a process called salt hydrolysis — the reaction of an ion with water to produce either an acidic or basic solution. The resulting pH depends entirely on the strength of the acid and base from which the salt was originally formed. This is why not all salt solutions are neutral: sodium acetate turns litmus blue, ammonium chloride turns it red, and sodium chloride leaves it unchanged. Salt hydrolysis is governed by the principles of ionic equilibrium. The hydrolysis constant, K_h, quantifies the extent to which an ion reacts with water. It is directly related to the water ionisation constant K_w and to the dissociation constants K_a or K_b of the parent acid or base. Understanding these relationships allows chemists to calculate the exact pH of any salt solution, a skill critical in biochemistry, medicine, and industrial processes. There are four categories of salts based on the strengths of the parent acid and base: (1) strong acid–strong base (SA–SB) salts such as NaCl, which are neutral; (2) strong acid–weak base (SA–WB) salts such as NH₄Cl, which are acidic; (3) weak acid–strong base (WA–SB) salts such as CH₃COONa, which are basic; and (4) weak acid–weak base (WA–WB) salts such as CH₃COONH₄, whose pH depends on the relative magnitudes of K_a and K_b.

Key takeaways
  • SA–SB Salts: No Hydrolysis (Neutral Solution)
  • SA–WB Salts: Cation Hydrolysis (Acidic Solution)
  • WA–SB Salts: Anion Hydrolysis (Basic Solution)
  • WA–WB Salts: Both Ions Hydrolyse
  • The Hydrolysis Constant K_h and Its Relationship to K_w, K_a, K_b

SA–SB Salts: No Hydrolysis (Neutral Solution)

Salts formed from a strong acid and a strong base, such as NaCl (from HCl and NaOH), do not hydrolyse. The Na⁺ ion has no tendency to react with water because NaOH is fully dissociated and does not recapture OH⁻. Similarly, Cl⁻ has no tendency to react with water because HCl is fully dissociated. The solution remains at pH = 7 at 25 °C. General rule: neither cation nor anion undergoes hydrolysis.

SA–WB Salts: Cation Hydrolysis (Acidic Solution)

Salts such as NH₄Cl are formed from the strong acid HCl and the weak base NH₃. In solution, the Cl⁻ ion does not hydrolyse, but the NH₄⁺ cation does: \ce{NH_4^+ + H_2O <=> NH_3 + H_3O^+}. This releases H₃O⁺, making the solution acidic (pH < 7). The hydrolysis constant is K_h = K_w / K_b. The pH can be calculated as pH = ½(pK_w − pK_b − log C), where C is the molar concentration of the salt.

WA–SB Salts: Anion Hydrolysis (Basic Solution)

Salts such as sodium acetate (CH₃COONa) come from the weak acid acetic acid and the strong base NaOH. The Na⁺ ion does not hydrolyse, but the acetate anion does: \ce{CH_3COO^- + H_2O <=> CH_3COOH + OH^-}. This produces OH⁻, making the solution basic (pH > 7). The hydrolysis constant is K_h = K_w / K_a. The pH is given by pH = ½(pK_w + pK_a + log C).

WA–WB Salts: Both Ions Hydrolyse

For salts like ammonium acetate (CH₃COONH₄), both ions undergo hydrolysis simultaneously: \ce{CH_3COO^- + NH_4^+ + H_2O <=> CH_3COOH + NH_3 + H_2O}. The pH is approximately pH = ½(pK_w + pK_a − pK_b), which is independent of concentration. If K_a = K_b, the solution is neutral. If K_a > K_b, the solution is acidic; if K_b > K_a, the solution is basic. The overall hydrolysis constant is K_h = K_w / (K_a · K_b).

The Hydrolysis Constant K_h and Its Relationship to K_w, K_a, K_b

The hydrolysis constant K_h is derived by combining the equilibrium expression for hydrolysis with the water autoionisation equilibrium. For anion hydrolysis: K_h = [HA][OH⁻] / [A⁻] = K_w / K_a. For cation hydrolysis: K_h = [B][H₃O⁺] / [BH⁺] = K_w / K_b. A larger K_h means greater extent of hydrolysis and a larger deviation from neutral pH. The degree of hydrolysis h = √(K_h / C) when hydrolysis is small, where C is the salt concentration.

The Key Equation

Kh=KwKa(for WA–SB salt),pH=12(pKw+pKa+log⁡C)K_h = \frac{K_w}{K_a} \quad \text{(for WA–SB salt)}, \qquad \mathrm{pH} = \frac{1}{2}\left(\mathrm{p}K_w + \mathrm{p}K_a + \log C\right)
  • K_h — Hydrolysis constant (dimensionless): Equilibrium constant for the reaction of an ion with water; measures the extent of hydrolysis.
  • K_w — Water ionisation constant (mol² dm⁻⁶): Product of [H⁺][OH⁻] in pure water; equals 1.0 × 10⁻¹⁴ at 25 °C.
  • K_a — Acid dissociation constant (mol dm⁻³): Equilibrium constant for the dissociation of the parent weak acid; a smaller K_a means a weaker acid and greater hydrolysis.
  • C — Molar concentration of salt (mol dm⁻³): The initial concentration of the dissolved salt; affects the degree of hydrolysis and final pH.
  • \mathrm{p}K_a — Negative logarithm of Ka (dimensionless): pKa = −log(Ka); a higher pKa indicates a weaker acid and a more basic salt solution.

Worked Example

Problem. Calculate the pH of a 0.10 mol dm⁻³ solution of sodium acetate (CH₃COONa) at 25 °C. Given: Ka(CH₃COOH) = 1.8 × 10⁻⁵, Kw = 1.0 × 10⁻¹⁴.

pH=12(14.00+4.745−1.000)=8.87\mathrm{pH} = \frac{1}{2}(14.00 + 4.745 - 1.000) = 8.87

Answer. pH = 8.87 (basic solution, as expected for a weak acid–strong base salt)

Figure
Diagram
Salt Hydrolysis: pH Summary of Four Salt TypesSA – SBe.g. NaClNo hydrolysisNa⁺ + Cl⁻ → no rxnpH = 7NeutralSA – WBe.g. NH₄ClCation hydrolysesNH₄⁺+H₂O⇌NH₃+H₃O⁺pH < 7AcidicWA – SBe.g. CH₃COONaAnion hydrolysesCH₃COO⁻+H₂O⇌HA+OH⁻pH > 7BasicWA – WBe.g. CH₃COONH₄Both hydrolysepH ≈ ½(pKw+pKa−pKb)pH ≈ 7 *Depends on Ka vs KbpH— Neutral pH = 7 —SA-WBNH₄ClpH≈5.1WA-SBCH₃COONapH≈8.9NaCl pH=7WA-WB≈70pH Scale Effect of Salt Type on Solution pHAcidic ←→ Basic* For WA–WB: if Ka = Kb, pH = 7; if Ka > Kb, pH < 7; if Kb > Ka, pH > 7
Origins · 1 minSvante Arrhenius and how this was discovered
SA
Svante Arrhenius
1859–1927 · Swedish

Arrhenius developed the ionic theory of electrolytic dissociation (1884), proposing that salts, acids, and bases break into ions in aqueous solution. This foundational concept directly underlies our understanding of salt hydrolysis and ionic equilibrium. He defined acids as substances that produce H⁺ ions in water and bases as substances that produce OH⁻ ions — the Arrhenius definitions still taught in introductory chemistry. His work earned him the Nobel Prize in Chemistry in 1903. Arrhenius's doctoral thesis on ionic dissociation — the work that won him the Nobel Prize — was initially awarded the lowest possible passing grade by his examining committee, who did not believe ions could exist independently in solution!

  • Chemical Equilibrium and Le Chatelier's Principle — Salt hydrolysis is an equilibrium process; understanding how the system responds to concentration changes requires Le Chatelier's principle and the equilibrium constant expression.
  • Weak Acids, Ka, and pH Calculations — The hydrolysis constant K_h = Kw/Ka directly involves Ka; calculating the pH of a WA–SB salt solution requires the same ICE-table method used for weak acid equilibria.
  • Buffer Solutions and the Henderson–Hasselbalch Equation — Partially hydrolysed salt solutions of weak acids or bases act as buffers; the Henderson–Hasselbalch equation is derived from the same equilibrium relationships used in salt hydrolysis.
  • Water Autoionisation and Kw — K_h is always defined relative to Kw = 1.0 × 10⁻¹⁴; understanding the temperature dependence and magnitude of Kw is essential for all ionic equilibrium calculations.