Topic 4 of 6
Buffers in calculations and oceans
Follow consumption of a small acid or base addition before using equilibrium.
A-Level 9476 (2026-2027)
A buffer contains a reservoir for both added acid and added base
A weak acid/conjugate-base pair or weak base/conjugate-acid pair limits small pH changes.
An acidic buffer contains appreciable amounts of a weak acid HA and its conjugate base A-, often made by mixing the acid with a soluble salt or partially neutralising the acid. Added H+ is consumed by A-; added OH- is consumed by HA. An alkaline buffer similarly contains B and BH+, such as NH3/NH4+.
| Addition | Acidic buffer reaction | Alkaline ammonia buffer reaction |
|---|---|---|
| A small amount of acid | A- + H+ → HA | NH3 + H+ → NH4+ |
| A small amount of base | HA + OH- → A- + H2O | NH4+ + OH- → NH3 + H2O |
These reactions prevent the added strong acid or base from remaining freely in solution. The component ratio changes only modestly when the addition is small relative to both reservoirs, so pH changes modestly. A buffer does not keep pH perfectly fixed, and it fails when one component is substantially exhausted.
Buffers are used to maintain suitable pH for enzyme reactions, chemical measurements and calibration. Diluting both components by the same factor leaves their ratio approximately unchanged, so ideal buffer pH changes little, but the smaller amounts per volume give lower buffer capacity. At extreme dilution, the simple approximation also breaks down.
Neutralise the added reagent first, then calculate the new ratio
The equilibrium formula uses the remaining buffer components.
Rearrange Ka = [H+][A-]/[HA]: pH = pKa + log10([A-]/[HA]). When both components share a solution volume, their mole ratio gives the same ratio. This buffer approximation uses the analytical component amounts after any strong acid/base reaction, provided equilibrium ionisation changes are small relative to those amounts.
Worked example
A buffer before and after added acid
100.0 cm3 of buffer contains 0.0200 mol HA and 0.0300 mol A-. Ka = 1.80 × 10-5. Find its initial pH and its pH after adding 0.00500 mol HCl, with volume effects common to both components.
- pKa = -log10(1.80 × 10-5) = 4.7447.
- Initial pH = 4.7447 + log10(0.0300/0.0200) = 4.92.
- Added H+ reacts with A-: new A- amount = 0.0250 mol and new HA amount = 0.0250 mol.
- The new ratio is 1, so pH = pKa.
pH changes from 4.92 to 4.74. Substituting the added HCl concentration directly into -log[H+] would ignore its reaction with the buffer.
For a weak base B and conjugate acid BH+, either use the acid constant of BH+ in the same equation or use pOH = pKb + log10([BH+]/[B]), followed by pH = pKw - pOH. Equal amounts of NH3 and NH4+ with Kb = 1.80 × 10-5 give pH about 9.26 at 298 K.
Check your understandingCan the buffer equation be used after adding 0.0400 mol HCl to the original buffer above?Think it through, then reveal the answer
Carbonate buffering limits acidification but cannot stop unlimited carbon dioxide input
More dissolved CO2 consumes carbonate and changes the balance of the ocean carbon system.
- CO2 enters the water
CO2(g) ⇌ CO2(aq). Dissolved CO2 participates in acid-base equilibria; the simplified combined step is CO2(aq) + H2O(l) ⇌ H+(aq) + HCO3-(aq).
- Carbonate consumes some added acid
CO32-(aq) + H+(aq) → HCO3-(aq). This limits the rise in free hydrogen-ion concentration.
- The buffer balance changes
The carbonate/hydrogencarbonate ratio falls, so the buffered pH falls. Overall: CO2 + CO32- + H2O → 2HCO3-.
The pair HCO3-/CO32- is an acid/conjugate-base buffer: carbonate accepts H+, while hydrogencarbonate can consume OH- to form carbonate and water. Ocean water contains other buffering species too, but this pair links increasing atmospheric CO2 directly to carbonate availability.
The rapid atmospheric CO2 increase drives additional uptake by surface waters. Buffering consumes carbonate as it resists the added acidity, so it has finite capacity and the equilibrium pH decreases. Ocean acidification means a fall in pH; it does not require seawater to become acidic below pH 7. Reduced carbonate availability also affects the equilibria involved in forming calcium carbonate shells and skeletons.