Doc Brown's Chemistry  Theoretical Chemistry - Equilibria - Chemical Equilibrium 6.3

6.3 Buffer solutions - definition, formulation and action

Revision notes for GCE Advanced Subsidiary Level AS Advanced Level A2 IB Revise AQA GCE Chemistry OCR GCE Chemistry Edexcel GCE Chemistry Salters Chemistry CIE Chemistry revising courses for pre-university students (equal to US grade 11 and grade 12 and Honours/honors level courses) KS4 Science GCSE/IGCSE Chemistry reversible reactions-equilibrium * KS4 Science GCSE/IGCSE notes acids and bases * KS4 Science GCSE/IGCSE notes acid-base theory

Equilibria Part 6 sub-index: 6.1 Salt hydrolysis * 6.2 Acid-base indicator theory, pH curves and titrations * 6.3 Buffers - definition, formulation and action * 6.4 Buffer calculations * 6.5 Case studies of buffer function * EMAIL query?comment

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, solubility product and ion-exchange * 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

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)

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