Hydrolysis of Salts — Predicting Acidic, Basic or Neutral
Not every salt gives a neutral solution when dissolved in water — this is one of the first genuine surprises in the Equilibrium chapter. Dissolve table salt (NaCl) and the water stays at pH 7. Dissolve ammonium chloride (NH₄Cl) and the solution turns acidic. Dissolve sodium acetate (CH₃COONa) and it turns basic. The reason is salt hydrolysis — the reaction of a salt's ions with water — and once you know the four-category rule, you can predict the nature of almost any salt solution without a single calculation.
What is salt hydrolysis?
A salt is formed from the reaction of an acid and a base. When that salt dissolves in water, its cation, its anion, or both can react with water molecules — this reaction is called hydrolysis. Whether hydrolysis occurs, and which ion is responsible, depends entirely on whether the parent acid and parent base were strong or weak.
Strong acid + weak base salt → acidic (cation hydrolyses)
Weak acid + strong base salt → basic (anion hydrolyses)
Weak acid + weak base salt → depends on relative Ka and Kb
The logic behind the rule: only the ion that came from the weak parent acid or base is able to react with water and shift the pH. An ion that came from a strong acid or strong base is such a poor proton donor/acceptor (conjugate of a strong electrolyte) that it simply does not react with water at all.
Case 1 — strong acid + strong base: neutral
NaCl comes from HCl (strong acid) and NaOH (strong base). Neither Na⁺ nor Cl⁻ can react with water, so the solution stays at pH 7. The same applies to KNO₃, Na₂SO₄ and similar salts.
Case 2 — strong acid + weak base: acidic (cation hydrolysis)
NH₄Cl comes from HCl (strong acid) and NH₃/NH₄OH (weak base). The NH₄⁺ ion is the conjugate acid of the weak base NH₃, so it reacts with water and releases H₃O⁺:
This lowers the pH below 7. For a salt of concentration C, using the hydrolysis constant Kh = Kw/Kb, the pH works out to:
Case 3 — weak acid + strong base: basic (anion hydrolysis)
CH₃COONa comes from CH₃COOH (weak acid) and NaOH (strong base). The acetate ion CH₃COO⁻ is the conjugate base of the weak acid, so it reacts with water and releases OH⁻:
Following the mirror-image derivation with Kh = Kw/Ka:
Case 4 — weak acid + weak base: it depends
Both ions hydrolyse here, so the result is a tug-of-war between the acid strength of the cation and the base strength of the anion. A useful, concentration-independent formula is:
If pKa = pKb (the parent acid and parent base are equally weak), the salt is neutral. If pKa > pKb (parent acid is the weaker of the two, meaning Ka is small), the salt is basic — and vice versa.
Worked example 1 — 0.1 M NH₄Cl
NH₃ has Kb = 1.8 × 10⁻⁵, so pKb = 4.74.
pH = 7 − ½(4.74) − ½ log(0.1) = 7 − 2.37 − ½(−1) = 7 − 2.37 + 0.5 = 5.13
The solution is acidic, as predicted by Case 2.
Worked example 2 — 0.1 M CH₃COONa
CH₃COOH has Ka = 1.8 × 10⁻⁵, so pKa = 4.74.
pH = 7 + ½(4.74) + ½ log(0.1) = 7 + 2.37 − 0.5 = 8.87
The solution is basic, as predicted by Case 3.
Worked example 3 — CH₃COONH₄ (weak acid + weak base)
CH₃COOH: pKa = 4.74. NH₃: pKb = 4.74. The two values happen to be
almost equal, which is why this particular salt is the textbook example of a "weak + weak"
salt that stays close to neutral:
pH = 7 + ½(4.74 − 4.74) = 7 + 0 = ≈ 7.0
Note this result does not depend on concentration at all — a direct consequence of both ions
hydrolysing together.
Worked example 4 — NH₄CN (weak acid + weak base, but clearly basic)
This example shows Case 4 does not always give a neutral result. HCN has
Ka = 6.2 × 10⁻¹⁰ (pKa = 9.21); NH₃ has Kb = 1.8 × 10⁻⁵
(pKb = 4.74).
pH = 7 + ½(9.21 − 4.74) = 7 + ½(4.47) = 7 + 2.24 = ≈ 9.24
HCN is a far weaker acid than NH₃ is a weak base (Kb ≫ Ka), so the
cyanide ion's hydrolysis dominates and the salt is clearly basic — not neutral.
Common mistakes that cost marks
- Assuming "weak + weak" always means neutral: CH₃COONH₄ is close to neutral only because Ka(CH₃COOH) and Kb(NH₃) happen to be almost equal. NH₄CN (worked example 4) proves the general rule wrong — always compare the actual Ka and Kb values.
- Forgetting which ion hydrolyses: in NH₄Cl it is the cation (NH₄⁺) that reacts with water, not the chloride ion. In CH₃COONa it is the anion (CH₃COO⁻), not the sodium ion. Only the ion descended from the weak parent electrolyte reacts.
- Ignoring the concentration term for Cases 2 and 3: the pH of an NH₄Cl or CH₃COONa solution genuinely changes with concentration (the ½ log C term), unlike the weak-acid/weak-base case, which is independent of concentration.
- Sign errors in the pH formula: log C is negative for any concentration below 1 M, so ½ log C actually subtracts from the answer in Case 3 and adds in Case 2 — track the sign carefully rather than assuming it always increases pH.
Where salt hydrolysis appears in exams
| Exam | Typical use |
|---|---|
| CBSE/ICSE Class 11 | Equilibrium chapter — classifying salts, writing hydrolysis equations |
| JEE/NEET | MCQs on predicting pH nature of a given salt, comparing salt pairs |
| IIT-JAM / CUET-PG | Numerical pH calculation using the hydrolysis-constant formulas |
| GATE / CSIR-NET | Relating Kh, degree of hydrolysis and buffer-adjacent equilibrium problems |
Check your acid-base numericals instantly. The free pH/pOH calculator and the Henderson–Hasselbalch calculator in the ABC Chemistry Suite let you verify hydrolysis pH values the moment you finish working them out by hand.
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