Buffer Action and Calculations

Specification Reference Physical Chemistry, Acids and bases 3.1.12.6

Quick Notes

A buffer solution minimises change to pH when small amounts of acid or base are added.

Acidic buffers contain a weak acid and the salt of the weak acid (its conjugate base). (e.g., CH₃COOH/CH₃COONa).
Basic buffers contain a weak base and the salt of the weak base (its conjugate acid) (e.g., NH₃/NH₄Cl).

For acidic buffers, an equilibrium is established between HA and A⁻.

AQA A-Level Chemistry equilibrium HA(aq) ⇌ H+(aq) + A−(aq) that underpins buffer action

The HA and A⁻ have a much greater concentration than H⁺, meaning that the position of the equilibrium will not be sensitive to changes in their concentrations, but it will be very sensitive to changes in the concentration of H⁺(aq) ions.
When H⁺(aq) ion concentration is increased or decreased, the position of equilibrium moves to oppose the change.

Full Notes

Buffers and example calculations have been outlined in more detail here.
This page is just what you need to know for AQA A-level Chemistry :)

What is a Buffer Solution?

A buffer solution maintains a relatively constant pH despite the addition of small amounts of acid or base. They ‘minimise’ change to pH.

Buffers are essential in biological systems and many industrial processes where a near constant pH is important.

Types of Buffers:

Acidic Buffers:
Made from a weak acid and its salt (that contains the acids conjugate base).

For example The weak acid ethanoic acid (CH₃COOH) and its salt sodium ethanoate (CH₃COONa).

AQA A-Level Chemistry diagram showing acidic buffer prepared from CH3COOH(aq) and CH3COONa(aq) producing a mixture containing CH3COOH and CH3COO−

When added to a solution of the ethanoic acid, the CH₃COONa would dissociate and release CH₃COO⁻ ions, which is the conjugate base (A⁻) of the ethanoic acid.

Basic Buffers:
Made from a weak base and its salt (that contains the conjugate acid of the base).

For example The weak base ammonia (NH₃) and its salt ammonium chloride (NH₄Cl).

AQA A-Level Chemistry diagram showing basic buffer prepared from NH3(aq) and NH4Cl(aq) producing a mixture containing NH3 and NH4+

When added to a solution of ammonia, the NH₄Cl would dissociate and release NH₄⁺ ions, which are the conjugate acid ions of the ammonia.

How Acidic Buffers Work

An equilibrium is established in the buffer system between HA, A⁻ and H⁺.

AQA A-Level Chemistry equilibrium HA(aq) ⇌ H+(aq) + A−(aq) for acidic buffer

The concentration of HA and A⁻ in the mixture must be much greater than the concentration of H⁺. This ensures the position of equilibrium is sensitive to changes in H⁺ concentration change more than changes to HA and A⁻ concentration. Equilibrium position can shift to keep H⁺ ion concentration nearly constant.

Example: Ethanoic Acid/Sodium Ethanoate Buffer
CH₃COOH ⇌ H⁺ + CH₃COO⁻

When an acid (H⁺) is added:

AQA A-Level Chemistry diagram showing A− reacting with added H+ to form HA in an acidic buffer

CH₃COO⁻ combines with added H⁺ to form CH₃COOH.
Equilibrium shifts left, reducing the increase in H⁺. [HA] increases and [A⁻] decreases.

When a base (OH⁻) is added:

AQA A-Level Chemistry diagram showing added OH− reacting with H+; HA dissociates to replace H+ in an acidic buffer

Added OH⁻ reacts with H⁺ to form H₂O.
CH₃COOH dissociates more to replace lost H⁺.
Equilibrium shifts right, replacing H⁺ ions that reacted with the added OH⁻, resisting pH increase. [HA] decreases and [A⁻] increases.

Matt’s exam tip

Remember that the concentration of HA and A^- will change when H⁺ or OH^- ions are added. When H⁺ ions are added to the mixture - the moles of HA will increase by the same as the moles of H⁺ added and moles of A^- decrease by the same amount. When OH^- ions are added - the moles of HA will decrease by the same as moles of OH^- added and the moles of A^- increase by the same amount.

Worked Example

Q: What happens to the pH when a small amount of NaOH is added to a buffer made of 0.10 mol dm^-3 CH₃COOH and 0.10 mol dm^-3 CH₃COO⁻?

Set up & assumptions
Take 1.00 dm³ of buffer so initial moles are CH₃COOH (HA) = 0.10 mol and CH₃COO⁻ (A⁻) = 0.10 mol. Add 0.010 mol NaOH (small amount). Assume volume change is negligible.
Neutralisation reaction (stoichiometry)
OH⁻ + HA → A⁻ + H₂O
New moles: HA = 0.10 − 0.010 = 0.090 mol; A⁻ = 0.10 + 0.010 = 0.110 mol; OH⁻ is fully consumed.
pH before addition (use the K_a expression)
For ethanoic acid, K_a = 1.75 × 10⁻⁵ mol dm⁻³.
K_a = \[\,[H⁺][A⁻] / [HA]\,\] ⇒ [H⁺] = K_a × [HA]/[A⁻].
With [HA] = [A⁻] = 0.10 mol dm^-3, [H⁺] = 1.75 × 10⁻⁵ mol dm^-3.
pH = −log(1.75 × 10⁻⁵) = 4.76.
pH after addition (again from K_a)
[HA] ≈ 0.090 mol dm^-3; [A⁻] ≈ 0.110 mol dm^-3 (total volume 1.00 dm³).
[H⁺] = K_a × ([HA]/[A⁻]) = 1.75 × 10⁻⁵ × (0.090/0.110) = 1.43 × 10⁻⁵ mol dm^-3.
pH = −log(1.43 × 10⁻⁵) ≈ 4.84.

A: The added OH⁻ reacts with CH₃COOH:
CH₃COOH + OH⁻ → CH₃COO⁻ + H₂O.
CH₃COOH decreases to 0.090 mol dm^-3 and CH₃COO⁻ increases to 0.110 mol dm^-3. Using the equilibrium expression for K_a, the pH changes only slightly: 4.76 → 4.84 (ΔpH ≈ +0.08), demonstrating buffer action.

How Basic Buffers Work

Basic buffers follow the same principle as for acidic buffers, with a position of equilibrium shifting to oppose a change in pH. This time however it's the OH⁻ ion concentration that the equilibrium is sensitive to.

Example Ammonia/Ammonium Chloride Buffer (NH₃/NH₄⁺)
Equilibrium reaction: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻

When an acid (H⁺) is added:

H⁺ reacts with OH⁻ to form H₂O.
Equilibrium shifts right, with more NH₃ reacting with H₂O to replace OH⁻ ions, meaning pH decrease is reduced.

When a base (OH⁻) is added:

NH₄⁺ reacts with added OH⁻ to form NH₃ and H₂O.
Equilibrium shifts left, reducing pH increase.

Summary

Idea or Concept	Explanation
Buffer Solution	Resists pH changes when small amounts of acid/base are added
Acidic Buffer	Weak acid + its salt (e.g., CH₃COOH/CH₃COO⁻)
Basic Buffer	Weak base + its salt (e.g., NH₃/NH₄⁺)
Buffer Action	Equilibrium shifts to absorb excess H⁺ or OH⁻