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Revision notes on chemical equilibrium - explaining the theory of how buffer solutions work Advanced Level Theoretical-Physical Chemistry

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Theoretical–Physical Advanced Level Chemistry – Equilibria – Chemical Equilibrium Revision Notes PART 6.3

6.3 Buffer solutions – definition, formulation and action

What is a buffer? How do buffers work?

Buffers are defined and their actions explained with appropriate examples such as ethanoic acid/sodium ethanoate and ammonia/ammonium chloride mixtures.

Chemical Equilibrium Notes Parts 5 & 6 Index

6.3 Buffer solutions – definition, formulation and action

  • 6.3.1 A buffer is a solution that minimises pH change on the addition of small amounts of acid or alkali.

    • Never say "it prevents change in pH", BUT it DO SAY it MINIMISES the CHANGE in pH.

  • Buffers and their chemical reactions must obey Le Chatelier's Equilibrium Concentration Principle, and act in a way to remove H+ and OH ions. BUT, they cannot theoretically prevent the pH being lowered/raised by the addition of acid/alkali, however small the change.

  • Note that any buffer will eventually be 'used up' if large quantities of acid or alkali are added to the solution.

  • 6.3.2 Typical buffers and their action.

    • The buffering chemistry is quite simple in principle and the ideas behind the examples described below can be applied to the design and action of most buffers.

  • Buffering action example 6.3.2a

    • A mixture of a weak acid and the salt of the weak acid with a strong base.

    • Organic acids like methanoic, ethanoic, propanoic, citric, benzenedicarboxylic etc. are frequently used in buffer mixtures i.e. those with the carboxylic acid functional group –COOH

    • The salts are usually those of the strong base–alkalis sodium and potassium hydroxide.

    • e.g. ethanoic acid CH3COOH and sodium ethanoate CH3COONa+ gives buffers in the range pH 3.7–5.6

    • CH3COOH and CH3COO constitute a conjugate acid–base pair.

    • In solution most of the weak acid is NOT ionised and the relatively high concentration of the CH3COO ion actually inhibits ionisation.

      • It is the relatively high concentration of the weak ethanoic acid that 'removes' any added hydroxide ions:

        • CH3COOH(aq) + OH(aq) (c) doc b CH3COO(aq) + H2O(l)

        • Note the use of the 'styled' reversible sign to show a bias to RHS of the equilibrium

    • The salt is fully ionised in solution to give a relatively high concentration of the ethanoate ions.

      • It is the ethanoate ion which removes most of any added hydrogen ions.

        • CH3COO(aq) + H+(aq) (c) doc b CH3COOH(aq)

        • The conjugate base of a weak acid is relatively strong!

    • How to choose the best weak acid and its corresponding salt is explained in section 6.4.1

  • Buffering action example 6.3.2b

    • A mixture of a weak base and the salt of the weak base with a strong acid.

    • e.g. ammonia NH3 and ammonium chloride NH4+Cl

    • NH4+ and NH3 constitute a conjugate acid–base pair.

    • In solution most of the ammonia is NOT ionised (and even suppressed by the ammonium ions from the salt).

      • It is the weak base that 'removes' most of any added hydrogen ions.

        • NH3(aq) + H+(aq) (c) doc b NH4+(aq)

    • The salt is fully ionised in solution giving relatively high concentrations of the ammonium ion.

      • It is the ammonium ion that removes most of any added hydroxide ions.

        • NH4+(aq) + OH(aq) (c) doc b NH3(aq) + H2O(l)

  • 6.3.3 Preparing buffer solutions.

    • Quite often several solutions of salts, weak acids/bases are prepared and then mixed in different ratios to provide buffers of a wide pH range.

    • Sometimes a single salt will do to give a single accurate pH value for calibrating a pH meter. (see Case study 6.5.1)


Chemical Equilibrium Notes Parts 5 & 6 Index

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