Doc Brown's Advanced A Level Chemistry Revision Notes

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 CH₃COO^–Na⁺ 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) 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) 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.

      • NH_3(aq) + H⁺_(aq) 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) NH_3(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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Chemical Equilibrium Notes Parts 5 & 6 Index