AQA A-Level Chemistry: Acids, Bases and Buffers — Complete Revision Guide (7405)
AQA A-Level Chemistry: Acids, Bases and Buffers — Complete Revision Guide (7405)
Acids, bases and buffers is one of the most calculation-heavy topics on AQA A-Level Chemistry (7405), and one of the most rewarding to master once the pattern clicks. The entire section, from the proton-transfer definition through to a multi-step buffer-after-addition problem, reduces to a small set of equilibrium expressions — Ka, Kw and Kb — combined with a logarithmic conversion to pH. Once you can move between concentrations and pH fluently, the standard AQA questions become clean substitutions rather than algebraic puzzles: strong acid pH, weak acid pH, buffer composition, pH-curve interpretation, indicator choice and the half-equivalence determination of pKa.
This topic also sits at the intersection of physical and analytical chemistry, so it rewards a methodical approach to units, significant figures and assumptions. Most students who lose marks here lose them through procedure, not concept.
Synoptic Preview — Where Acids and Buffers Fit Into 7405
Acids and buffers is one of the most synoptic blocks in AQA 7405. The equilibrium constants here — Ka, Kb, Kw — are specific cases of the Kc framework developed in kinetics and equilibrium. The acid-base behaviour of carboxylic acids (pKa around 4 to 5) and amino-acid zwitterions in organic advanced is described directly by Ka. The titration arithmetic — moles, concentrations, end points — builds on the mole concept and titration calculations introduced in atomic structure and amounts.
Recognising those connections is worth real marks. AQA's longer-response items often blend topics: a buffer question can hide inside an amino-acid problem; a titration calculation can be the first step of a Required Practical write-up. Time spent on acids and buffers returns across the rest of the course.
Guide Overview — Eight Sub-Topics
This guide walks through the AQA 7405 acids, bases and buffers content in eight focused sections. Each gives you the core ideas, the common pitfalls, a worked example and a direct link to the matching LearningBro lesson.
- Brønsted-Lowry Acids and Bases
- Acids, Bases and pH
- Ka, Kw and pKa
- pH of Weak Acids and Bases
- Buffer Solutions and pH Curves
- pH Curves and Indicators
- Polyprotic Acids and Salt Hydrolysis
- Required Practical 9: Titration and pH Curves
Together these match the lesson order of the LearningBro Acids, Bases and Buffers course, so you can pair this guide directly with the course quizzes and AI tutor feedback.
What the AQA 7405 Specification Covers
AQA A-Level Chemistry (7405) is examined through three papers: Paper 1 (Inorganic and Physical, 2h, 105 marks), Paper 2 (Organic and Physical, 2h, 105 marks) and Paper 3 (synoptic plus practical, 2h, 90 marks). The acids, bases and buffers content sits in specification section §3.1.12 and is examined heavily on Paper 1 and on Paper 3 through practical and synoptic items.
| Sub-topic | Spec area | Typical paper weight |
|---|---|---|
| Brønsted-Lowry acids and bases | §3.1.12.1 | 2-4 marks |
| Definition and measurement of pH | §3.1.12.2 | 3-5 marks |
| Ka, Kw and pKa | §3.1.12.3-4 | 3-6 marks |
| pH of weak acids and bases | §3.1.12.4 | 4-6 marks |
| Buffer action and pH curves | §3.1.12.5-6 | 4-8 marks |
| Indicator choice | §3.1.12.6 | 2-4 marks |
| Polyprotic acids and salt hydrolysis | §3.1.12 (synoptic) | 2-4 marks |
| Required Practical 9 | §3.1.12 / RP9 | 4-8 marks |
Weights are estimates modelled on recent 7405 paper structures, but the reliable observation is that an extended-response buffer or pH-after-titration question — typically eight to ten marks — appears on essentially every Paper 1, and Required Practical 9 surfaces on most Paper 3s.
Brønsted-Lowry Acids and Bases
The Brønsted-Lowry theory defines an acid as a proton donor and a base as a proton acceptor. This definition is broader than the Arrhenius picture (acids release H+ in water, bases release OH-) and applies in non-aqueous solvents and to species without explicit OH groups, which AQA exploits in synoptic questions on amines and amino acids.
Every Brønsted acid has a corresponding conjugate base — what is left after the proton has been donated — and every Brønsted base has a conjugate acid. These come in pairs: HCl/Cl-, CH3COOH/CH3COO-, NH4+/NH3, H2O/OH- (water acting as an acid gives H3O+/H2O; water acting as a base gives H2O/OH-). Water is amphoteric because it can do either, which is why its autoionisation underpins the whole pH scale.
Worked example. Identify the conjugate acid-base pairs in NH3 + H2O ⇌ NH4+ + OH-. NH3 (base) and NH4+ (its conjugate acid) form one pair; H2O (acid) and OH- (its conjugate base) form the other. The proton has been transferred from H2O to NH3.
A common pitfall is to write the conjugate base as the original acid plus an extra proton rather than minus one proton. Another is to call something a Brønsted-Lowry acid without specifying that it is a proton donor — AQA mark schemes reward the specific phrase.
See the Brønsted-Lowry lesson.
Acids, Bases and pH
The pH scale converts hydrogen-ion concentration to a logarithmic value: pH = -log10 [H+], with [H+] in mol dm^-3. The inverse is [H+] = 10^-pH. The autoionisation of water gives Kw = [H+][OH-] = 1.0 × 10^-14 mol^2 dm^-6 at 298 K, so in any aqueous solution at 298 K, pH + pOH = 14.
For a strong acid that fully dissociates in water, [H+] equals the initial acid concentration. For a strong base that fully dissociates, [OH-] equals the initial base concentration and [H+] is recovered through Kw.
Worked example. Calculate the pH of 0.050 mol dm^-3 HNO3. HNO3 is a strong monoprotic acid, so [H+] = 0.050 mol dm^-3 and pH = -log10 (0.050) = 1.30.
Worked example. Calculate the pH of 0.10 mol dm^-3 NaOH. [OH-] = 0.10, [H+] = 1.0 × 10^-14 / 0.10 = 1.0 × 10^-13, pH = 13.0.
Common pitfalls: forgetting the negative sign in pH = -log10 [H+]; quoting too many decimals (three significant figures of [H+] gives two decimals of pH); and treating Kw as temperature-independent when AQA explicitly tests Kw values at temperatures other than 298 K.
See the acids, bases and pH lesson.
Ka, Kw and pKa
For a weak acid HA dissociating in water, HA ⇌ H+ + A-, the acid dissociation constant is Ka = [H+][A-] / [HA]. The smaller Ka, the weaker the acid. The logarithmic form, pKa = -log10 Ka, gives a more convenient scale: larger pKa, weaker acid.
| Acid | Ka / mol dm^-3 | pKa |
|---|---|---|
| HCl (strong) | very large | very negative |
| HNO3 (strong) | ~24 | ~-1.4 |
| CCl3COOH | 0.20 | 0.7 |
| HF | 6.6 × 10^-4 | 3.2 |
| CH3COOH | 1.8 × 10^-5 | 4.76 |
| HCN | 4.9 × 10^-10 | 9.3 |
Strong acids essentially fully dissociate (Ka very large). Weak acids only partially dissociate; the equilibrium lies well to the left, an observation the German chemist Ostwald formalised in his dilution law.
Kw itself has the dimensions and structure of an equilibrium constant (Kw = [H+][OH-]), and for any conjugate acid-base pair Ka × Kb = Kw, so pKa + pKb = 14 at 298 K. This is heavily used in synoptic questions linking weak acids and their conjugate bases.
A common pitfall is to write pKa = log Ka instead of -log Ka. Another is to forget that Ka, like all equilibrium constants, depends on temperature — AQA can give a value at a non-standard temperature and expect candidates to use it directly.
See the Ka, Kw and pKa lesson.
pH of Weak Acids and Bases
For a weak acid HA at initial concentration c, only a small fraction dissociates. Two AQA-standard assumptions apply: [H+] = [A-] (each dissociation event gives one of each) and [HA] ≈ c (the small loss to dissociation is negligible). Substituting into Ka gives Ka = [H+]^2 / c, so [H+] = √(Ka × c) and pH = ½ (pKa - log10 c).
Worked example. Calculate the pH of 0.10 mol dm^-3 ethanoic acid (Ka = 1.8 × 10^-5 mol dm^-3). [H+] = √(1.8 × 10^-5 × 0.10) = √(1.8 × 10^-6) = 1.34 × 10^-3 mol dm^-3. pH = -log10 (1.34 × 10^-3) = 2.87.
The assumption [HA] ≈ c is valid when less than around 5 percent of the acid dissociates. For very dilute weak acids this fails and a quadratic is required; AQA rarely sets such cases at A-Level but you should be able to spot when the approximation breaks.
For a weak base B + H2O ⇌ BH+ + OH-, Kb = [BH+][OH-] / [B], analogous to Ka. Use Kw to convert to [H+] and then to pH, or use pKa + pKb = 14 if you already have pKa of the conjugate acid.
A common pitfall is to forget the square root and write [H+] = Ka × c. Another is to use the initial weak-base concentration as [OH-] — which is only true for strong bases.
See the pH of weak acids and bases lesson.
Buffer Solutions and pH Curves
A buffer is a solution that resists changes in pH when small amounts of acid or base are added. AQA emphasises two buffer compositions: a weak acid with its conjugate base (acidic buffer, e.g. ethanoic acid plus sodium ethanoate), and a weak base with its conjugate acid (basic buffer, e.g. ammonia plus ammonium chloride).
The pH of a buffer is calculated from a rearrangement of Ka, sometimes called the Henderson-Hasselbalch equation: pH = pKa + log10 ([A-] / [HA]). When [A-] = [HA], pH = pKa, so the buffer's working pH is set by the choice of weak acid. The German biochemist Henderson and the Danish chemist Hasselbalch arrived at the logarithmic form independently in the early twentieth century.
Worked example. Calculate the pH of a buffer made from 0.20 mol dm^-3 ethanoic acid (pKa 4.76) and 0.10 mol dm^-3 sodium ethanoate. pH = 4.76 + log10 (0.10 / 0.20) = 4.76 - 0.30 = 4.46.
The buffer capacity — the amount of acid or base that can be absorbed before pH changes significantly — is largest when [HA] = [A-] (pH = pKa) and decreases as the ratio moves away from 1. When acid is added, the conjugate base A- mops it up: A- + H+ → HA. When base is added, HA neutralises the OH-: HA + OH- → A- + H2O.
A common pitfall is inverting the ratio in the Henderson-Hasselbalch expression — the conjugate base is on top.
See the buffers and titrations lesson.
pH Curves and Indicators
A pH curve plots the pH of the analyte solution against the volume of titrant added. The shape depends on whether the acid and base are strong or weak, and AQA expects you to recognise all four combinations.
| Combination | Initial pH | Equivalence pH | Steep portion | Suitable indicator |
|---|---|---|---|---|
| Strong acid + strong base | ~1 | 7 | pH 3-11 | Methyl orange or phenolphthalein |
| Strong acid + weak base | ~1 | <7 | pH 3-7 | Methyl orange |
| Weak acid + strong base | ~3 | >7 | pH 7-11 | Phenolphthalein |
| Weak acid + weak base | ~3 | ~7 | No clear steep section | pH meter — indicators unsuitable |
An indicator is itself a weak acid (HIn) whose conjugate base In- has a different colour. The colour change spans approximately pKa(In) ± 1, and the indicator's pKa must lie within the steep section of the pH curve to give a sharp end point. The three AQA-standard indicators are methyl orange (pKa 3.7, red to yellow), bromothymol blue (pKa 7.0, yellow to blue) and phenolphthalein (pKa 9.3, colourless to pink).
Worked example. For 0.10 mol dm^-3 ethanoic acid against 0.10 mol dm^-3 NaOH, equivalence pH is around 8.7 with a steep section between pH 7 and pH 11. Phenolphthalein (pKa 9.3) sits inside this band and gives a sharp end point; methyl orange changes too early and is unsuitable.
A common pitfall is choosing an indicator just because its pKa lies between 4 and 10, without checking the curve. Another is forgetting that the equivalence point is not always pH 7.
See the pH curves and indicators lesson.
Polyprotic Acids and Salt Hydrolysis
Polyprotic acids release more than one proton per molecule. Sulfuric acid is diprotic; its first dissociation (H2SO4 → H+ + HSO4-) is essentially complete, and the second (HSO4- ⇌ H+ + SO4^2-) has Ka around 1.2 × 10^-2, so it is weak but not negligible. For typical A-Level dilute H2SO4 problems, AQA usually treats both protons as fully dissociated; for 0.050 mol dm^-3 H2SO4 the working [H+] is 0.10 mol dm^-3, giving pH 1.0.
Phosphoric acid (H3PO4) is triprotic, with pKa1 around 2.1, pKa2 around 7.2 and pKa3 around 12.4. Each successive dissociation is weaker because removing a proton from an increasingly negative species is harder. AQA examines this mainly through qualitative pH-curve recognition — three buffer regions and three equivalence steps.
Salt hydrolysis is the flip side: when a salt of a weak acid (e.g. sodium ethanoate) dissolves in water, the conjugate base partially reprotonates water, giving a slightly basic solution. Salts of weak bases (e.g. ammonium chloride) give slightly acidic solutions. This is why a weak-acid/strong-base titration has an equivalence pH above 7 — the salt solution at equivalence is itself basic.
A common pitfall is forgetting that the second proton of H2SO4 is not strictly strong; for very accurate work the HSO4- equilibrium matters.
See the polyprotic acids and salt hydrolysis lesson.
Required Practical 9: Titration and pH Curves
Required Practical 9 asks you to measure a pH curve for the addition of a base to an acid. The standard AQA method uses a calibrated pH meter rather than indicator drops, because the goal is the full curve, not a single end point.
Method outline. Calibrate the pH meter with buffers at pH 4 and pH 7 (and pH 10 if available). Pipette a known volume of acid (typically 25.0 cm^3) into a beaker with a magnetic stirrer. Record the initial pH. Add base from a burette in small increments — large increments far from equivalence, very small (0.1-0.2 cm^3) increments near the steep section — and record pH after each addition until the curve has clearly flattened past equivalence. Plot pH against volume added.
Key analyses. (1) The equivalence point is the centre of the steep section, not the inflection where pH crosses 7. (2) The half-equivalence volume — half the volume needed to reach equivalence — corresponds to [HA] = [A-], so pH at that point equals pKa of the weak acid. This is the standard AQA experimental route to pKa.
Required-practical pitfalls. Not stirring continuously means the recorded pH lags the true value; failing to take small increments near equivalence smears out the steep section; reading the half-equivalence volume from the wrong axis is a classic mark-loser. AQA can also ask you to evaluate the precision of the pH-meter reading versus the burette reading and to suggest improvements.
See the Required Practical 9 lesson.
Required Practical 9 — Quick Reference
- Calibrate the pH meter against standard buffers before use.
- Use a magnetic stirrer; record initial pH before adding titrant.
- Increment finely (0.1-0.2 cm^3) near the steep section; coarsely elsewhere.
- Equivalence is the steep-section midpoint; half-equivalence pH = pKa.
- Compare your experimental pKa to the data-book value as evaluation.
Exam Strategy and Common Mark-Loss Patterns
- Forgetting the negative sign in pH = -log10 [H+] or in pKa = -log10 Ka.
- Treating [H+] of a weak acid as equal to its initial concentration.
- Forgetting the doubling for diprotic strong acids (H2SO4 at A-Level).
- Inverting the ratio in the Henderson-Hasselbalch equation.
- Using base concentration as [OH-] for weak bases.
- Choosing an indicator whose pKa lies outside the steep section.
- Assuming the equivalence point is always pH 7.
- Misreading the half-equivalence volume on an RP9 curve.
- Failing to quote three significant figures on Ka-derived [H+] values.
How to revise. Drill pH conversions both ways daily, twenty conversions for a fortnight, until they become automatic. Practise weak-acid pH problems in batches of ten until √(Ka × c) feels obvious. Build a four-card flashcard set for the pH-curve combinations and a three-card set for indicators. Run through one RP9 dataset a week and extract pKa from the half-equivalence volume.
Related Reading and Next Steps
Acids and buffers builds on the kinetics and equilibrium framework — Ka, Kb and Kw are equilibrium constants and inherit every property of Kc. The titration arithmetic is a direct extension of the mole and concentration work in atomic structure and amounts. The acid-base behaviour of carboxylic acids, phenols and amino-acid zwitterions in organic advanced is exactly the buffer chemistry of weak acids and their conjugate bases.
If you want a structured, exam-paced walkthrough of every sub-topic above with worked examples, AI tutor feedback and timed quizzes, the full LearningBro Acids, Bases and Buffers course is the most direct way to convert this guide into reliable marks. Get this section fluent and a substantial fraction of the physical chemistry on every AQA 7405 paper becomes routine.