OCR A-Level Chemistry: Acids, Bases and Buffers — Complete Revision Guide (H432)
OCR A-Level Chemistry: Acids, Bases and Buffers
Acids, bases and buffers is the quantitative A2 extension of the AS Brønsted-Lowry framework introduced in Acids, Redox, Electrons and Bonding. At AS, acids and bases were classified qualitatively and titration arithmetic was developed; at A2, the pH scale is introduced, weak acids are characterised by their acid dissociation constant Ka, and the buffering principle is put to work. The four titration-curve shapes (strong-strong, strong-weak, weak-strong, weak-weak) and the corresponding indicator-selection logic are routine items on Paper 1 (Periodic Table, Elements and Physical Chemistry) at A2 and underpin Paper 3 (Unified Chemistry) synoptic items where buffers connect to biological context.
H432 examiners gravitate towards this module because it is genuinely discriminating: the arithmetic is short, but every step encodes a conceptual decision (strong or weak? mono- or diprotic? base attacking acid or acid attacking base?) that separates fluent candidates from those who have memorised a single template. The same physical apparatus — pH probe, burette, conical flask, magnetic stirrer — produces qualitatively different curves depending on which species are reacting, so a single experimental skill is leveraged into four exam scenarios. Module 5.1.3 also acts as the gateway to the Module 5.2.3 redox-titration arithmetic and to the Module 6.3 biological-buffer commentary on Paper 3. Candidates who internalise the Ka expression as a rearrangeable identity rather than as a memorised formula find the buffer, half-equivalence-point and weak-acid pH calculations collapse into a single workflow.
Course 8 of the H432 Chemistry learning path on LearningBro, Acids, Bases and Buffers, develops the pH calculation toolkit that turns the AS-level Brønsted-Lowry framework into predictive chemistry. It builds in three phases: the pH scale, Kw and strong acid/base pH calculations; the Ka framework for weak acids and the calculation of pH from a weak-acid concentration; and the buffer-design framework with Henderson-Hasselbalch logic, four titration curves, and indicator choice. It sits adjacent to Quantitative Rates and Equilibrium and feeds into Energetics and Electrode Potentials on the OCR A-Level Chemistry learning path.
Guide Overview
The Acids, Bases and Buffers course is built as a sequence of lessons that move from Brønsted definitions through pH calculations into buffers and titration curves.
- Brønsted-Lowry Acid-Base Theory Recap
- Kw and the pH Scale
- pH of Strong Acids
- Weak Acids and Ka
- pH of Weak Acids
- pH of Strong Bases
- Buffer Principle and Composition
- Buffer pH Calculations
- Titration Curves: Strong-Strong
- Titration Curves: Strong-Weak and Weak-Strong
- Titration Curves: Weak-Weak
- Indicators and Indicator Selection
OCR H432 Specification Coverage
This course addresses OCR H432 Module 5.1.3 (acids, bases and buffers). The specification organises the topic into the pH scale grounded in Kw, the Ka treatment of weak acids, buffer theory and titration curves with indicator selection (refer to the official OCR specification document for exact wording).
| Sub-topic | Spec area | Primary lesson(s) |
|---|---|---|
| Brønsted-Lowry conjugate pairs; acid/base reactions | OCR H432 Module 5.1.3 | Brønsted-Lowry Acid-Base Theory Recap |
| Kw; pH = -log[H⁺]; pOH | OCR H432 Module 5.1.3 | Kw and the pH Scale |
| Strong acid pH calculations | OCR H432 Module 5.1.3 | pH of Strong Acids |
| Ka and weak acid dissociation | OCR H432 Module 5.1.3 | Weak Acids and Ka |
| pH of weak acids from Ka | OCR H432 Module 5.1.3 | pH of Weak Acids |
| Strong base pH via Kw | OCR H432 Module 5.1.3 | pH of Strong Bases |
| Buffer principle; buffer pH | OCR H432 Module 5.1.3 | Buffer Principle and Composition; Buffer pH Calculations |
| Titration curve shapes for the four acid-base combinations | OCR H432 Module 5.1.3 | Titration Curves: Strong-Strong; Strong-Weak and Weak-Strong; Weak-Weak |
| Indicator endpoint selection | OCR H432 Module 5.1.3 | Indicators and Indicator Selection |
Module 5.1.3 is examined on Paper 1 with routine pH and buffer calculations, and on Paper 3 with titration-curve sketching and indicator selection in unfamiliar contexts (biological buffers, environmental water chemistry, food chemistry).
Topic-by-Topic Walkthrough
Brønsted Recap, Kw and the pH Scale
The Brønsted-Lowry recap lesson reaffirms the conjugate-pair view: acid HA → A⁻ + H⁺, with HA and A⁻ as the acid and its conjugate base. Strong conjugate acids have weak conjugate bases and vice versa. The Kw lesson develops the autoionisation of water (2H₂O ⇌ H₃O⁺ + OH⁻) with Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ mol² dm⁻⁶ at 298 K, and the consequent definitions pH = -log[H⁺], pOH = -log[OH⁻], and pH + pOH = 14. Worked check: pure water has [H⁺] = [OH⁻] = 10⁻⁷, so pH = 7 at 298 K. Kw increases with temperature (autoionisation is endothermic), so the pH of pure water decreases slightly above 25 °C — but the water remains neutral because [H⁺] still equals [OH⁻]. This nuance is a Paper 1 mark-bearing distinction.
Strong Acid pH and Weak Acid Ka
The strong acid pH lesson develops the assumption that strong acids fully dissociate, so [H⁺] equals the initial acid concentration and pH = -log[concentration]. Worked example: 0.0500 mol dm⁻³ HCl gives [H⁺] = 0.0500, pH = -log(0.0500) = 1.30. Common pitfalls include forgetting that a diprotic strong acid (H₂SO₄) gives two H⁺ per formula unit, and that very dilute strong acids (≤ 10⁻⁶ mol dm⁻³) require the water-autoionisation contribution to be included.
The weak acids and Ka lesson develops the dissociation HA ⇌ H⁺ + A⁻ with Ka = [H⁺][A⁻]/[HA]. Strong acids have Ka much greater than 1 (essentially complete dissociation); weak acids have Ka much less than 1 (typical organic acids have Ka ~ 10⁻⁵, giving pKa ~ 5). pKa = -log Ka, so the smaller the pKa, the stronger the acid. The pH of weak acids lesson develops the standard simplifying assumptions: [H⁺] = [A⁻] (so the dissociation is the only source of each), and [HA]_eq ≈ [HA]_initial (so very little dissociates). This gives [H⁺] = √(Ka × [HA]). Worked example: 0.100 mol dm⁻³ ethanoic acid (Ka = 1.8 × 10⁻⁵) gives [H⁺] = √(1.8 × 10⁻⁵ × 0.100) = 1.34 × 10⁻³, pH = 2.87.
pH of Strong Bases and Buffer Calculations
The pH of strong bases lesson develops the route via Kw: strong bases fully dissociate to give [OH⁻] = initial concentration, then [H⁺] = Kw/[OH⁻], then pH = -log[H⁺]. Worked example: 0.0500 mol dm⁻³ NaOH gives [OH⁻] = 0.0500, [H⁺] = 10⁻¹⁴/0.0500 = 2.0 × 10⁻¹³, pH = 12.70.
The buffer principle lesson develops the two routes to a buffer: a weak acid plus its conjugate base in similar concentrations (acidic buffer, e.g. CH₃COOH/CH₃COO⁻Na⁺), or a weak base plus its conjugate acid (basic buffer, e.g. NH₃/NH₄Cl). When acid is added, the conjugate base accepts H⁺; when base is added, the weak acid donates H⁺ — both neutralise the added species and leave [H⁺] approximately unchanged. The buffer pH calculations lesson develops the route via the Ka expression rearranged: [H⁺] = Ka × [HA]/[A⁻]. Worked example: a buffer made from 0.20 mol dm⁻³ ethanoic acid and 0.30 mol dm⁻³ sodium ethanoate gives [H⁺] = 1.8 × 10⁻⁵ × 0.20/0.30 = 1.2 × 10⁻⁵, pH = 4.92. Adding small amounts of acid or base shifts the ratio but not enough to change the pH appreciably.
Titration Curves and Indicators
The titration-curve lessons develop the four canonical shapes. Strong-strong starts at low pH, rises sharply through pH 7 at equivalence (vertical rise from about pH 3 to pH 11), then plateaus at high pH. Strong-weak — strong acid into weak base — has equivalence below pH 7 because the conjugate acid of the weak base is mildly acidic; weak-strong (weak acid into strong base) has equivalence above pH 7 for the symmetric reason. Crucially, the weak-strong curve has a buffer region around pH = pKa of the weak acid where the slope is shallow, and the half-equivalence point gives pH = pKa exactly — a routine experimental technique for measuring Ka. Weak-weak curves have no sharp inflection and cannot be titrated to a precise endpoint.
The indicators and indicator selection lesson develops indicators as weak acids whose conjugate forms have different colours, with a working range of about ±1 pH unit around their pKa. Methyl orange (pKa ~ 3.7, red-to-yellow at 3.1-4.4) is appropriate for strong-strong and strong-weak titrations because its colour change falls within the steep vertical inflection. Phenolphthalein (pKa ~ 9.3, colourless-to-pink at 8.3-10) is appropriate for strong-strong and weak-strong titrations. Neither works for weak-weak. The match of indicator working range to titration inflection range is the routine selection logic and a common Paper 1 explanation item.
A Typical H432 Paper 1 Question
A standard Paper 1 prompt on this material gives candidates a weak acid by name (often a named carboxylic acid or biological acid such as lactic acid), a numerical Ka, and a concentration, then asks for the pH of (a) the pure weak acid, (b) a buffer made by partial neutralisation with NaOH, and (c) the pH after a small further addition of strong acid. The route is fixed: for (a), use [H⁺] = √(Ka × [HA]) and convert to pH; for (b), recognise that partial neutralisation generates the conjugate base in known moles while leaving the unreacted weak acid, then apply [H⁺] = Ka × n(HA)/n(A⁻); for (c), adjust the moles of HA and A⁻ by the moles of strong acid added, then recompute. The discriminator at the top band is the explicit statement that volumes cancel in the buffer ratio (so concentrations can be replaced by moles), and the explicit assumption that the weak acid dissociation is suppressed by the common-ion presence of the conjugate base.
Synoptic Links
Acids, bases and buffers connect through the rest of the spec. The biological buffering of blood plasma at pH 7.4 by the carbonate/bicarbonate system is the canonical synoptic application revisited in transition elements and aromatic when haemoglobin's pH dependence is discussed. The amphoteric nature of amino acids and the resulting isoelectric point connects to carbonyls, polymers and spectroscopy. The Ka-style equilibrium expression generalises directly from the Kc and Kp framework of quantitative rates and equilibrium — buffers are simply equilibria in which a Le Chatelier shift counters added perturbation. The mole arithmetic of titration calculations reuses the framework introduced in atoms, moles and equations.
Paper 3 'Unified chemistry' items routinely deploy this module against an unfamiliar context. A biological scenario might present the buffering of intracellular fluid by the dihydrogen-phosphate/monohydrogen-phosphate pair, with candidates expected to identify pKa2 of phosphoric acid as the relevant constant and then perform a Henderson-Hasselbalch-style calculation. An environmental scenario might present the carbonate buffering of natural waters and the pH consequence of dissolved CO₂ from acid deposition. An industrial scenario might present a fermentation broth whose pH must be held inside the optimum window for the active enzymes, with candidates asked to choose between a citrate and a phosphate buffer on the grounds of pKa proximity. The chemistry is identical to the Paper 1 calculations; the synoptic mark comes from naming the species, sourcing the right pKa, and explaining qualitatively how each addition shifts the equilibrium.
What Examiners Reward
Top-band marks on this module hinge on the explicit handling of assumptions rather than on raw arithmetic. For weak-acid pH calculations, examiners want the line "we assume [H⁺] = [A⁻] (no other source of either) and [HA]_eq ≈ [HA]_initial (negligible dissociation)" before the square-root substitution is taken. For buffer calculations, they want the explicit statement that volumes cancel in the [HA]/[A⁻] ratio. For titration curves, they want the identification of the buffer region around pH = pKa as the shallow-slope plateau before the steep equivalence inflection, and the identification of the half-equivalence point as the place where [HA] = [A⁻] and pH = pKa.
Common pitfalls cluster around five recurring mistakes. First, forgetting that diprotic strong acids deliver two H⁺ per formula unit, so 0.10 mol dm⁻³ H₂SO₄ has [H⁺] = 0.20 mol dm⁻³ and pH = 0.70, not pH = 1.00. Second, applying the weak-acid pH formula to a buffer (where the conjugate base is no longer at trace concentration and [H⁺] ≠ [A⁻]). Third, omitting the Kw step for strong-base pH and writing pH = -log(0.05) = 1.30 instead of pH = 12.70. Fourth, selecting phenolphthalein for a strong-acid-weak-base titration where the equivalence sits well below pH 7 and phenolphthalein has already changed colour by then. Fifth, choosing indicators whose colour change misses the vertical inflection entirely on weak-weak titrations (where no sharp inflection exists and no indicator is appropriate). Each of these is a one- or two-mark deduction that compounds quickly across a multi-part question.
Practical Activity Groups (PAGs)
This course anchors PAG 11 (pH measurement) through the experimental titration of strong and weak acids, the determination of pKa from the half-equivalence point of a weak-acid-strong-base titration, and the preparation and characterisation of buffer solutions. The procedure — calibrate the pH probe with two standard buffers, titrate slowly through the equivalence region with continuous pH readout, identify equivalence and half-equivalence — is the routine practical-skills item that ties this course to assessed lab work.
Going Further
Undergraduate analogues of this material extend in several directions. First, the Brønsted-Lowry framework generalises into Lewis acid-base theory (electron-pair donors and acceptors) and hard-soft acid-base theory, foundational for inorganic chemistry. Second, the buffer-design framework generalises into biological buffering systems (haemoglobin, phosphate, bicarbonate) which together maintain blood plasma pH at 7.40 ± 0.05 under enormous metabolic perturbation. Third, pH and Ka generalise into the activity coefficients and ionic-strength corrections needed for accurate work in concentrated electrolytes. Fourth, the polyprotic-acid speciation diagrams that undergraduates draw for amino acids, citrate and EDTA generalise the half-equivalence-point pKa logic of this course into multi-step distribution curves, and the same diagrams reappear in undergraduate biochemistry as the basis for explaining enzyme pH optima. Fifth, the thermodynamic underpinning of acid strength — the standard Gibbs energy of dissociation linked through ΔG = -RT ln K to Ka — bridges this module with the energetics material in Energetics and Electrode Potentials and is the conceptual jump that distinguishes a chemistry undergraduate's treatment from an A-Level treatment.
Oxbridge-style interview prompts on this material include: "Why does the pH of pure water decrease as temperature rises, while the water itself remains neutral?" "Sketch the titration curve of a diprotic weak acid (e.g. carbonic acid) against a strong base, identifying the two equivalence and two half-equivalence points." "What concentration of ethanoic acid and sodium ethanoate would you need to mix to make a buffer at pH 5.00, and what is its buffering capacity?" "Glycine has pKa1 ≈ 2.3 (carboxyl) and pKa2 ≈ 9.6 (ammonium). Predict its isoelectric point and explain why it is the average of the two pKa values." "A pH meter is calibrated at 25 °C but the unknown solution is at 50 °C. Will the reported pH be too high or too low for a sample of pure water, and roughly by how much?"
Authorship and Sign-off
This guide was authored independently by John Haigh, paraphrasing OCR H432 Module 5.1.3 as descriptive use. No verbatim spec text, mark-scheme phrasing, examiner-report quotation, or past-paper question reference appears. The worked examples are original.
Start at the Acids, Bases and Buffers course and work through every lesson in sequence. Once the pH scale, the Ka framework, the buffer-pH calculation and the four titration-curve shapes are automatic, every acid-base item across the H432 series resolves into a substitution-and-rearrangement routine rather than a guessing exercise.