Buffer action (A-level only) — AQA A-Level Chemistry
Test yourself on Buffer action (A-level only) with AQA A-Level practice questions.
A week of Premium from your first lesson · Then the free plan continues · No card, no charge
Buffer action (A-level only) explained
A buffer holds the pH nearly steady because it contains two things at once in substantial amounts: a weak acid HA and a large reservoir of its conjugate base A⁻.
Read the full explanation
Add acid and the A⁻ removes the H⁺: A⁻ + H⁺ → HA. Add alkali and the HA removes the OH⁻: HA + OH⁻ → A⁻ + H₂O. In each case one member of the pair is converted into the other rather than being left free in solution. Since [H⁺] = Ka × [HA]/[A⁻], only the ratio matters, and when both amounts are large a small addition barely changes that ratio. Dilution divides both concentrations by the same factor, so the ratio, and therefore the pH, is unchanged. The protection is approximate and finite: add enough acid to consume the A⁻ and the pH falls sharply.
Acidic buffer solutions contain a weak acid and the salt of that weak acid.
An acidic buffer is a weak acid together with the salt of that acid, for example ethanoic acid with sodium or potassium ethanoate. The salt is fully ionic, so it supplies a large concentration of ethanoate ions, far more than the acid's own slight dissociation could ever provide. There are two ways to make one. Mix the weak acid with its solid salt. Or partially neutralise the weak acid: add the strong base to an excess of the acid, so that CH₃COOH + KOH → CH₃COOK + H₂O converts some of the acid into the salt and leaves the rest unreacted. Both routes finish with the mixture of CH₃COOH and CH₃COO⁻ that buffering requires. Added H⁺ is removed by the ethanoate ion and added OH⁻ by the ethanoic acid molecule, and naming those two species is what earns the marks.
Basic buffer solutions contain a weak base and the salt of that weak base.
A basic buffer is a weak base with the salt of that base, the classic pair being aqueous ammonia with ammonium chloride. The salt supplies the ammonium ions and the base supplies the ammonia molecules, so both halves of the conjugate pair NH₄⁺ and NH₃ are present in quantity and the solution settles above pH 7. Added acid is removed by the base: NH₃ + H⁺ → NH₄⁺. Added alkali is removed by the salt: NH₄⁺ + OH⁻ → NH₃ + H₂O. The pairs that do not work are worth recognising. Two weak bases such as ammonia and methylamine are not a conjugate pair; two weak acids are not either; and a strong acid mixed with a strong base simply gives a neutral salt solution with nothing left over to absorb further additions.
Applications of buffer solutions.
Blood plasma is the example to know. It is held between pH 7.35 and 7.45 by the carbonic acid and hydrogencarbonate pair, H₂CO₃ ⇌ H⁺ + HCO₃⁻, fed by dissolved carbon dioxide. Lactic acid produced during exercise is absorbed by HCO₃⁻ + H⁺ → H₂CO₃, and the extra carbon dioxide is then breathed out, which is part of why the system copes so well. A shift of only a few tenths of a pH unit disrupts the shapes of enzymes and of haemoglobin, so the margin is narrow. Elsewhere, shampoos and skin creams are buffered near pH 5.5 to match skin; foods are buffered with citric acid and a citrate salt to control flavour and inhibit bacteria; and buffer solutions of known pH are what a pH meter is calibrated against before use.
Students should be able to: explain qualitatively the action of acidic and basic buffers calculate the pH of acidic buffer solutions.
Qualitatively, name the species: in an acidic buffer the conjugate base removes added H⁺ and the weak acid removes added OH⁻. Quantitatively, rearrange Ka to give [H⁺] = Ka × [HA]/[A⁻]. Because both species share the same solution, the volumes cancel and moles can be used directly in the ratio. Take 20.0 cm³ of 0.100 mol dm⁻³ NaOH added to 25.0 cm³ of 0.100 mol dm⁻³ butanoic acid, for which Ka = 1.51 × 10⁻⁵ mol dm⁻³. There are 2.50 × 10⁻³ mol of acid and 2.00 × 10⁻³ mol of alkali, so 5.0 × 10⁻⁴ mol of acid survives and 2.00 × 10⁻³ mol of salt is formed. Then [H⁺] = 1.51 × 10⁻⁵ × (5.0 × 10⁻⁴)/(2.00 × 10⁻³) = 3.78 × 10⁻⁶ mol dm⁻³ and pH = 5.42.
Your focus
- Write the two equations showing how an ethanoic acid and ethanoate buffer removes added H⁺ and added OH⁻.
- Explain why diluting a buffer with water leaves its pH almost unchanged.
- Predict what happens to the pH when acid is added beyond the capacity of the buffer.
Show all 15 objectives
- Describe two ways of preparing an ethanoic acid and ethanoate buffer from reagents available in a school laboratory.
- Write the equation for the partial neutralisation of ethanoic acid by potassium hydroxide.
- Identify which species in an acidic buffer removes added H⁺ and write the equation for that reaction.
- Select, from a list of pairs of solutions, the one that forms a basic buffer and justify the choice.
- Write the two equations showing how an ammonia and ammonium chloride buffer resists added acid and added alkali.
- Explain why mixing equal volumes of hydrochloric acid and sodium hydroxide of equal concentration does not give a buffer.
- Write the equilibrium that buffers blood plasma and state the pH range it maintains.
- Explain, with an equation, how the hydrogencarbonate ion removes acid produced by exercising muscle.
- Suggest why a buffer is added to a shampoo and state the approximate pH chosen.
- Calculate the pH of the buffer formed when 20.0 cm³ of 0.100 mol dm⁻³ sodium hydroxide is added to 25.0 cm³ of 0.100 mol dm⁻³ butanoic acid.
- Explain why the volumes cancel in the buffer expression, so that moles may be used in place of concentrations.
- State which species removes added acid and which removes added alkali in a butanoic acid and butanoate buffer, giving an equation for each.
Buffer action (A-level only) exam tips
Marking Points
- one mark for stating that a buffer contains a weak acid together with a reservoir of its conjugate base
- one mark for an equation showing added H⁺ being removed by the conjugate base
- one mark for an equation showing added OH⁻ being removed by the weak acid
- one mark for stating that [H⁺] depends on the ratio [HA]/[A⁻], which dilution does not change
- one mark for saying the pH changes only slightly rather than staying exactly constant
- one mark for adding an excess of the weak acid to the strong base, so the acid is only partially neutralised
- one mark for stating that the mixture then contains both the weak acid and its conjugate base
- one mark for the neutralisation equation, for example CH₃COOH + KOH → CH₃COOK + H₂O
- one mark for stating that CH₃COO⁻ from the salt reacts with the added H⁺
- one mark for stating that CH₃COOH reacts with the added OH⁻
- one mark for identifying a mixture of a weak base with its salt, such as ammonia with ammonium chloride
- one mark for the equation NH₃ + H⁺ → NH₄⁺ for the removal of added acid
- one mark for the equation NH₄⁺ + OH⁻ → NH₃ + H₂O for the removal of added alkali
- one mark for rejecting a pair of two weak acids, or a strong acid with a strong base, as not forming a buffer
- one mark for naming a real application, such as blood plasma held near pH 7.4
- one mark for identifying the conjugate pair responsible, for example H₂CO₃ and HCO₃⁻
- one mark for an equation showing the added H⁺ being removed by HCO₃⁻
- one mark for stating why the pH must be controlled, such as enzymes losing activity outside a narrow range
- one mark for the initial moles of the weak acid and of the added alkali
- one mark for subtracting to find the moles of weak acid remaining and the moles of salt formed
- one mark for substituting into [H⁺] = Ka × [HA]/[A⁻], in which the volumes cancel
- one mark for pH = −log₁₀[H⁺] quoted to two decimal places
- one mark for naming the species that removes added H⁺ and the species that removes added OH⁻
Examiner Tips
- 💡Every buffer explanation needs two equations, one for added acid and one for added alkali; one equation caps the marks.
- 💡Use the word 'approximately' or 'only slightly' about the pH; 'constant' on its own is often refused.
- 💡'Add excess ethanoic acid to the potassium hydroxide' is the phrasing that scores; 'mix them' is not enough.
- 💡Name CH₃COO⁻ and CH₃COOH explicitly in the explanation rather than talking about 'the equilibrium'.
- 💡In a multiple-choice buffer question, look for two species that differ by one H⁺; that is the conjugate pair.
- 💡Write ionic equations for basic buffers, since NH₃ and NH₄⁺ are the species that actually react.
- 💡Give a number for the pH being maintained; a named application with no value rarely gets full credit.
- 💡Follow every application with the consequence of losing control, such as loss of enzyme activity.
- 💡Set the working out as a table of initial moles, change and final moles; the marks follow that structure.
- 💡Check the answer against pKa: a buffer with more salt than acid must have a pH above the pKa.
Common Mistakes
- writing that a buffer keeps the pH exactly constant
- saying the buffer 'neutralises' the added acid without naming the species that reacts
- claiming dilution changes the pH because both concentrations fall
- using OH⁻ as the species that removes added acid in an acidic buffer
- assuming a buffer works for any amount of added acid, ignoring its finite capacity
- adding an excess of potassium hydroxide, which destroys all the weak acid and leaves no buffer
- describing the mixture as a weak acid and a strong base, when the base has been fully consumed
- writing the salt as CH₃COOHK instead of CH₃COOK or CH₃COO⁻K⁺
- explaining the action only by saying 'the equilibrium shifts', which examiners ignore unless a species is named
- choosing a strong acid and its salt, which cannot buffer because the acid is fully dissociated
- choosing two weak bases such as ammonia and methylamine, which are not a conjugate pair
- choosing hydrochloric acid with sodium hydroxide, which gives only a neutral salt solution
- writing NH₄OH as the species present in aqueous ammonia
- using OH⁻ rather than the ammonia molecule as the species that removes added acid
- stating that a basic buffer holds the pH at exactly 7
- naming the blood buffer as carbon dioxide and water rather than the H₂CO₃ and HCO₃⁻ pair
- stating that blood is buffered at pH 7.0 rather than at about 7.4
- claiming a buffer can restore the pH after any amount of acid is added
- describing a shampoo as 'pH balanced' with no named acid and salt
- converting one species to a concentration using the total volume but leaving the other as moles
- inverting the ratio, using [A⁻]/[HA], so the pH comes out too low
- forgetting that the moles of salt formed equal the moles of alkali added
- treating all of the original weak acid as still present after the alkali has reacted
- subtracting rather than adding the log term when using pH = pKa + log([A⁻]/[HA])