Buffers and Their Action

Specification Reference Topic 12, points 18–22, 24

Quick Notes

A Buffer solution resists pH change when small amounts of acid or base are added to it.
Buffers are typically made by:
- Mixing a weak acid and its conjugate base (e.g. ethanoic acid + sodium ethanoate)
- Mixing Excess weak acid with strong base, forming some salt in situ.
Acidic buffers maintains equilibrium:
HA ⇌ H⁺ + A⁻
- If acid (H⁺) is added: A⁻ removes it by forming more HA.
- If base (OH⁻) is added: HA donates H⁺ to neutralise OH⁻.
Buffers are important in biological, industrial, and laboratory settings.
In blood: carbonic acid–hydrogencarbonate buffer maintains pH ~7.4.

Full Notes

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

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. For example, in living organisms buffers maintain an optimum pH to prevent enzymes from being denatured.

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.

Buffers in Titration Curves

In a weak acid + strong base titration:

Edexcel A-Level Chemistry titration curve highlighting the buffer region for a weak acid with strong base.

Before the equivalence point, both HA and A⁻ are present, forming a buffer system.
The pH changes only minimally as base is added, however when the concentration of HA falls too low, the mixture is no longer able to act as a buffer and the pH changes rapidly as more base is added.

Buffers in Blood

Blood pH is controlled by a hydrogencarbonate buffer system:

H₂CO₃ ⇌ H⁺ + HCO₃⁻

If H⁺ is increased then HCO₃⁻ removes H⁺, forming H₂CO₃.
If OH⁻ is added then H₂CO₃ releases H⁺, neutralising the OH⁻.

This maintains blood pH around 7.4.

Summary

Buffers resist pH change by having significant amounts of HA and A⁻ (or base and its conjugate acid) present.
Acidic buffers: weak acid plus its salt and basic buffers: weak base plus its salt.
Added H⁺ is removed by A⁻; added OH⁻ is neutralised by HA.
Buffer pH can be calculated from K_a, [HA], and [A⁻].
Buffer regions appear before equivalence in weak acid–strong base titrations.
The hydrogencarbonate buffer maintains blood pH near 7.4.