Advanced A level theoretical chemistry - acid-base equilibria

Part 6.1 Salt Hydrolysis, acidity and alkalinity of salt solutions

[Author ©  Dr Phil Brown PhD: Doc Brown's exam revision notes suitable for A level chemistry students of advanced pre–university/college advanced level theoretical–physical chemistry courses:  acid–base equilibrium revision notes on salt hydrolysis and pH [updated April 29th 2026 *]

6.1 Salt Hydrolysis, acidity and alkalinity of salt solutions

What is salt hydrolysis? 

Why can salt solutions can be either neutral, alkaline or acidic?

• Despite being taught at lower academic levels that salts e.g. sodium chloride, dissolve in water to form neutral solutions of pH 7.

• In reality, and looking at a wider variety of 'salts', the picture is much more complicated and a 'salt' solution may be acid, neutral or alkaline depending on the nature of the interaction of the salt ions with water – do any of the salt ions have an acidic or basic nature – can they act as a Bronsted–Lowry acid or base.

• The reasons are quite clear when you consider the possible Bronsted–Lowry interactions that can take place between the ions of the salt and water.

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• 6.1.1 Examples of acidic salt solutions: pH <7

  • 6.1.1a: Hexa–aqa metal cations often show acidic behaviour particularly with their salts with strong acids.

    • e.g. the hydrated aluminium ion from aqueous aluminium chloride/sulphate/nitrate.

    • [Al(H₂O)₆]³⁺_(aq) + H₂O_(l) [Al(H₂O)₅(OH)]²⁺_(aq) + H₃O⁺_(aq) 

    • or more simply: [Al(H₂O)₆]³⁺_(aq) [Al(H₂O)₅(OH)]²⁺_(aq) + H⁺_(aq) 

    • The hydrated metal ion acts as an acid (proton donor) and water acts as the base (proton acceptor) and so aqueous hydrogen/oxonium ions are formed.

    • The greater the charge on the central metal ion, the stronger the hexa–aqua ion acid. e.g.

    • [Al(H₂O)₆]³⁺_(aq) > [Mg(H₂O)₆]²⁺_(aq) > [Na(H₂O)₆]⁺_(aq)

    • for the cations of Groups 1–3 on Period 3.

    • or [Fe(H₂O)₆]³⁺_(aq) > [Fe(H₂O)₆]²⁺_(aq) (in the 3d–block transition metal example)

    • From left to right, the trend is due to a decreasing charge density effect of the central metal ion on the O–H bond of a co–ordinated water molecule. The charge density decreases as the positive charge of the central metal ion decreases and its ionic radius increases.

    • The sodium ion shows virtually no acidic behaviour.

    • Further discussion of this situation will be on the Transition Metals Appendix page section 1 (currently under production).

    • However, the anion of the salt must not be neglected for a full explanation. The anions derived from the very strong  hydrochloric/sulphuric/nitric acids are all very weak bases and so have little tendency to interact with water in an acid–base reaction. Its a general, and logical rule, that the conjugate base of a very strong acid is very weak.

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  • 6.1.1b: Salts of weak bases and strong acids give acidic solutions.

    • e.g. ammonium chloride. The chloride ion is such a weak base that there is no acid–base reaction with water, but the ammonium ion is an effective proton donor. As a general rule, the conjugate acid of a weak base is quite strong. The result here is that ammonium salt solutions have a pH of 3–4.

    • NH₄⁺_(aq) + H₂O_(l) NH_3(aq) + H₃O⁺_(aq)

    • In zinc–carbon batteries an acidic ammonium chloride paste dissolves the zinc in the cell reaction, though an oxidising agent must be added (MnO₂) to oxidise the hydrogen formed into water, or batteries might regularly explode!

    • If you place a piece of magnesium ribbon or a zinc granule in ammonium chloride or ammonium sulphate solution you will see fizzing as hydrogen gas is formed.

      • 2H₃O⁺_(aq) + M_(s) ==> M²⁺_(aq) + H₂O_(l) + H_2(g)

      • M = zinc or magnesium

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• 6.1.2 Examples of nearly neutral salt solutions: pH approx. 7

  • 6.1.2a: Salts of strong acids and strong bases e.g. sodium chloride

    • Here the hexa–aqua sodium ion shows no acidic behaviour and the chloride ion no base behaviour, so little or no interaction with water to produce either H⁺_(aq) or OH^–_(aq) to change the pH.

  • 6.1.2b: Salts of weak acids and weak bases: e.g. ammonium ethanoate

    • Here the ammonium ion can act as an acid to form H⁺_(aq) with water, but the ethanoate ion acts as a base to give OH^–_(aq) with water, so they effectively neutralise each other.

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• 6.1.3 Examples of alkaline salt solutions: pH>7

  • 6.1.3a: Salts of a weak acid and a strong base e.g. sodium ethanoate

    • The hydrated sodium ion shows no acidic character but the ethanoate ion is a strong conjugate base of a the weak ethanoic acid (pK_a = 4.76,  K_a = 1.74 x 10^–5 mol dm^–3), so an acid–base hydrolysis reaction occurs to generate hydroxide ions to raise the pH to about pH 9.

      • CH₃COO^–_(aq) + H₂O_(l) CH₃COOH_(aq) + OH^–_(aq)

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  • 6.1.3b: Potassium cyanide: is the salt of the very strong base potassium hydroxide and the very weak hydrocyanic acid (pK_a = 9.31, K_a = 4.9 x 10^–10 mol dm^–3). The hydrated potassium ion shows no acidic behaviour, but the cyanide ion is a strong conjugate base of the very weak hydrocyanic acid (HCN) which interacts with water to generate hydroxide ions. Hydrocyanic acid (pK_a = 9.4) is weaker than ethanoic acid (pK_a = 4.76) , so the equilibrium is more on the right, more OH^–, and so the pH is more alkaline, i.e. over 9.

    • CN^–_(aq) + H₂O_(l) HCN_(aq) + OH^–_(aq)

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  • 6.1.3c: Sodium carbonate is the 'salt' of the strong base sodium hydroxide and the very weak 'carbonic acid'

    • Again the hydrated sodium ion shows no acidic character but the carbonate ion is a strong conjugate base of a the weak 'carbonic' acid, so an acid–base hydrolysis reaction occurs to generate hydroxide ions to raise the pH.

    • CO₃^2–_(aq) + H₂O_(l)  HCO₃^–_(aq) + OH^–_(aq)

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WHAT NEXT?

Advanced Equilibrium Chemistry Notes Part 1. Equilibrium, Le Chatelier's Principle–rules * Part 2. K_c and K_p equilibrium expressions and calculations * Part 3. Equilibria and industrial processes * Part 4 Partition between two phases, solubility product K_sp, common ion effect, ion–exchange systems * Part 5. pH, weak–strong acid–base theory and calculations * Part 6. Salt hydrolysis, acid–base titrations–indicators, pH curves and buffers * Part 7. Redox equilibria, half–cell electrode potentials, electrolysis and electrochemical series * Part 8. Phase equilibria–vapour pressure, boiling point and intermolecular forces watch out for sub–indexes to multiple sections or pages