Doc Brown's Advanced A Level Chemistry Revision Notes

Theoretical–Physical Advanced Level Chemistry – Equilibria – Chemical Equilibrium Revision Notes PART 6.1

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?

Chemical Equilibrium Notes Parts 5 & 6 Index

 (I shouldn't use it, but M = old fashioned shorthand for mol dm–3 !)



6.1 Salt Hydrolysis, acidity and alkalinity of salt solutions

  • 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.

  • 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.

  • 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) (c) doc b [Al(H2O)5(OH)]2+(aq) + H3O+(aq) 

        • or more simply: [Al(H2O)6]3+(aq) (c) doc b [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) (the cations of Gps 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.

    • 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) (c) doc b NH3(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) + H2(g)

        • M = zinc or magnesium

  • 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.

  • 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 (pKa = 4.76,  Ka = 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) (c) doc b CH3COOH(aq) + OH(aq)

    • 6.1.3b: Potassium cyanide: is the salt of the very strong base potassium hydroxide and the very weak hydrocyanic acid (pKa = 9.31, Ka = 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 (pKa = 9.4) is weaker than ethanoic acid (pKa = 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) (c) doc b HCN(aq) + OH(aq)

    • 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.

      • CO32–(aq) + H2O(l) (c) doc b HCO3(aq) + OH(aq)

Chemical Equilibrium Notes Parts 5 & 6 Index


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