Salt Hydrolysis: Why Some Salts Are Acidic or Basic

The word “salt” suggests something neutral. Dissolve table salt in water and the pH stays at 7. So it’s a surprise to find that a solution of sodium carbonate has a pH of about 11, ammonium chloride about 5, and iron(III) chloride as low as 2. These are all salts, formed by neutralisation, yet their solutions can be as acidic as vinegar or as alkaline as household ammonia.

The explanation is salt hydrolysis: some of the ions in a salt react with water, releasing H⁺ or OH⁻. Understanding when that happens lets you predict the pH of any salt solution.

What “hydrolysis” means here

Hydrolysis literally means “splitting with water”. In salt hydrolysis, an ion from the salt reacts with water molecules, taking or donating a proton:

• An anion that’s a weak base takes a proton from water, leaving OH⁻ behind → the solution becomes basic.
• A cation that’s a weak acid gives a proton to water, forming H₃O⁺ → the solution becomes acidic.

Ions that are too weak to do either are spectators and have no effect on pH.

Which ions hydrolyse?

The key is the Brønsted–Lowry link between conjugates: the weaker an acid, the stronger its conjugate base, and vice versa. See Ka × Kb = Kw.

Anions

Anion comes from…	Example	Behaviour in water
a strong acid	Cl⁻, Br⁻, I⁻, NO₃⁻, ClO₄⁻, (SO₄²⁻, only very slightly)	no hydrolysis: neutral
a weak acid	CH₃COO⁻, F⁻, CO₃²⁻, PO₄³⁻, CN⁻, S²⁻	accepts H⁺ from water: basic

Example, the ethanoate ion:

CH₃COO⁻ + H₂O ⇌ CH₃COOH + OH⁻

Cations

Cation	Example	Behaviour in water
Group 1 and heavier Group 2 metals	Na⁺, K⁺, Ca²⁺, Ba²⁺	no hydrolysis: neutral
Conjugate acid of a weak base	NH₄⁺, CH₃NH₃⁺	donates H⁺ to water: acidic
Small, highly charged metal ions	Al³⁺, Fe³⁺, Cr³⁺, Zn²⁺, Cu²⁺	hydrated ion releases H⁺: acidic

Example, the ammonium ion:

NH₄⁺ + H₂O ⇌ NH₃ + H₃O⁺

The four types of salt

Every salt can be thought of as made from an acid and a base. Their strengths predict the result:

Parent acid	Parent base	Example salt	Solution
strong	strong	NaCl, KNO₃, CaCl₂	neutral (pH 7)
weak	strong	CH₃COONa, Na₂CO₃, KF	basic (pH > 7)
strong	weak	NH₄Cl, NH₄NO₃, (CH₃NH₃)Cl	acidic (pH < 7)
weak	weak	CH₃COONH₄, (NH₄)₂CO₃	depends: compare Ka of cation with Kb of anion

A quick way to remember it: the stronger parent wins. Salt of a strong base and a weak acid behaves like a (weak) base; salt of a strong acid and a weak base behaves like a (weak) acid.

Hydrated metal ions: acidity without hydrogen

Why is iron(III) chloride acidic? It contains no ammonium ion and no acidic hydrogen.

In water, small, highly charged metal ions are surrounded by water molecules bonded through their oxygen lone pairs, forming ions such as [Fe(H₂O)₆]³⁺. The metal’s high positive charge density pulls electron density away from the O–H bonds of the attached water molecules. Those O–H bonds are weakened, and one of them can release a proton:

[Fe(H₂O)₆]³⁺ + H₂O ⇌ [Fe(H₂O)₅(OH)]²⁺ + H₃O⁺

The hydrated ion is acting as a Brønsted–Lowry acid.

Hydrated ion	pKa (approx.)	Effect
[Fe(H₂O)₆]³⁺	2.2	about as acidic as phosphoric acid’s first step
[Cr(H₂O)₆]³⁺	4.0	acidic
[Al(H₂O)₆]³⁺	5.0	similar to ethanoic acid
[Cu(H₂O)₆]²⁺	about 7.5–8	slightly acidic in practice
[Na(H₂O)ₙ]⁺	about 14 or more	negligible

The pattern follows charge density: higher charge and smaller radius mean more acidic hydrated ions. That’s why 3+ ions are much more acidic than 2+ ions, and Na⁺ and K⁺ have no effect. This is also why aluminium sulfate is added to soil to make hydrangeas blue and why iron(III) salt solutions often look orange-brown and cloudy: they’re partly hydrolysed to iron hydroxide species.

Calculating the pH of a salt solution

Treat the hydrolysing ion as a weak acid or weak base, using Ka or Kb. If you only have the constant for the parent, convert with Ka × Kb = Kw.

Basic salt

pH of 0.10 mol/dm³ sodium ethanoate (Ka of ethanoic acid = 1.8 × 10⁻⁵).

• Kb(CH₃COO⁻) = 1.0 × 10⁻¹⁴ ÷ 1.8 × 10⁻⁵ = 5.6 × 10⁻¹⁰
• [OH⁻] ≈ √(5.6 × 10⁻¹⁰ × 0.10) = 7.5 × 10⁻⁶ mol/dm³
• pOH = 5.13, so pH = 8.87

Acidic salt

pH of 0.20 mol/dm³ ammonium chloride (Kb of NH₃ = 1.8 × 10⁻⁵).

• Ka(NH₄⁺) = 1.0 × 10⁻¹⁴ ÷ 1.8 × 10⁻⁵ = 5.6 × 10⁻¹⁰
• [H⁺] ≈ √(5.6 × 10⁻¹⁰ × 0.20) = 1.06 × 10⁻⁵ mol/dm³
• pH = 4.98

Metal ion salt

pH of 0.10 mol/dm³ iron(III) nitrate (pKa of [Fe(H₂O)₆]³⁺ = 2.2).

• Ka = 10⁻²·² = 6.3 × 10⁻³
• Shortcut √(Ka × c) = 0.025, which is 25% of c, so solve the quadratic: x² + (6.3 × 10⁻³)x − 6.3 × 10⁻⁴ = 0 → x = 0.022 mol/dm³
• pH ≈ 1.65

A solution of an iron(III) salt can be more acidic than many common acids. More worked examples are in pH of weak acids and pH of weak bases.

Weak acid + weak base salt

For a salt like ammonium ethanoate, both ions hydrolyse. The pH depends on which is stronger:

• Ka(NH₄⁺) = 5.6 × 10⁻¹⁰ and Kb(CH₃COO⁻) = 5.6 × 10⁻¹⁰. They’re equal, so ammonium ethanoate solution is almost exactly neutral.
• Ammonium carbonate: Kb(CO₃²⁻) = 2.1 × 10⁻⁴ is far bigger than Ka(NH₄⁺), so the solution is basic.
• Ammonium fluoride: Ka(NH₄⁺) = 5.6 × 10⁻¹⁰ vs Kb(F⁻) = 1.5 × 10⁻¹¹. The cation wins slightly, so the solution is slightly acidic.

For these salts, a useful approximation (independent of concentration) is: pH ≈ ½(pKw + pKa(cation acid) − pKb(anion)), which for ammonium ethanoate gives ½(14.00 + 9.25 − 9.26) ≈ 7.0.

Amphiprotic anions

Ions like HCO₃⁻, HSO₄⁻ and H₂PO₄⁻ can act as either an acid or a base. Which effect wins depends on comparing their Ka and Kb:

• NaHCO₃: HCO₃⁻ is a slightly better base than acid, so the solution is mildly basic (about pH 8.3).
• NaHSO₄: HSO₄⁻ is a fairly strong acid (pKa about 2), so the solution is clearly acidic (about pH 1–2 at typical concentrations).
• NaH₂PO₄: slightly acidic (about pH 4.5); Na₂HPO₄: basic (about pH 9.8).

For an amphiprotic ion, a good estimate is pH ≈ ½(pKa₁ + pKa₂), where the two pKa values are those on either side of the ion. For HCO₃⁻: ½(6.35 + 10.33) = 8.34.

Everyday examples

• Baking soda (NaHCO₃) is mildly alkaline, which is why it neutralises acid indigestion and reacts with acidic ingredients in baking.
• Washing soda (Na₂CO₃·10H₂O) is strongly alkaline enough to cut grease.
• Ammonium sulfate fertiliser slowly acidifies soil.
• Alum and aluminium sulfate are acidic and are used in water treatment and to adjust soil pH.
• Soap is the sodium salt of a weak fatty acid, so soap solutions are alkaline (pH 9–10).
• Trisodium phosphate (TSP), a heavy-duty cleaner, is strongly alkaline because PO₄³⁻ is a fairly strong base.

Common mistakes

• Assuming all salts are neutral.
• Thinking Na⁺ or Cl⁻ affect pH. Neither hydrolyses measurably.
• Forgetting metal ions. Al³⁺ and Fe³⁺ salts are acidic even though they contain no hydrogen.
• Using Ka of the parent acid directly for the anion. Convert to Kb first.

Key takeaways

• Salt hydrolysis is the reaction of a salt’s ions with water, releasing H⁺ or OH⁻.
• Anions of weak acids make solutions basic; cations of weak bases and small, highly charged metal ions make them acidic.
• Salts of strong acids and strong bases are neutral; for weak acid–weak base salts, compare Ka and Kb.
• Calculate salt pH by treating the hydrolysing ion as a weak acid or base, using Ka × Kb = Kw.
• Check any result with the pH calculator.