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Acids and bases pervade chemistry — from gastric hydrochloric acid to the carbonate buffering that keeps blood pH at 7.40. A-Level Chemistry adopts the Brønsted-Lowry definition (1923): an acid is a proton (H⁺) donor; a base is a proton acceptor. This supersedes the older Arrhenius picture and underpins everything in this course: pH, weak-acid equilibria, buffers, indicators, titration curves. This lesson establishes the machinery — conjugate pairs, water as amphoteric, polyprotic acids, and the strength/concentration distinction. The Lewis definition (also 1923) is signposted as undergraduate but not examined at A-Level.
Spec mapping (AQA 7405): This lesson maps to §3.1.12.1 (Brønsted-Lowry acid-base equilibria in aqueous solution) and introduces the conceptual foundation for lessons 1 (pH calculations), 2 (Ka, Kw, and pKa for weak acids), and onwards through buffers (§3.1.12.5) and indicators (§3.1.12.6). It draws on prior knowledge of covalent and dative bonding from §3.1.3 — proton transfer is mechanistically the formation of a new dative covalent bond from the base's lone pair to the H⁺. It also connects to §3.2 (Group 1 and Group 2 hydroxides as familiar examples of strong bases). Refer to the official AQA 7405 specification for the exact wording of each subsection.
Assessment objectives: AO1 covers recall of the Brønsted-Lowry definitions and the recognition of conjugate acid-base pairs. AO2 requires students to identify the acid, base, and conjugate species in given equations, and to write balanced ionic equations for proton-transfer reactions. AO3 — the discriminator at the top of the grade band — tests rationalisation of amphoteric behaviour (water, hydrogencarbonate, amino acids) and reasoning about why conjugate strengths are inversely related to parent strengths. Expect Brønsted-Lowry questions on every Paper 2 and synoptically on Paper 3.
The conceptual history of acid-base chemistry runs through three increasingly general definitions. Knowing all three — and knowing which one AQA examines — saves marks.
Arrhenius defined an acid as a substance that produces H⁺ ions in aqueous solution, and a base as a substance that produces OH⁻ ions in aqueous solution. Familiar examples: HCl(aq) → H⁺(aq) + Cl⁻(aq); NaOH(aq) → Na⁺(aq) + OH⁻(aq); neutralisation is H⁺ + OH⁻ → H₂O.
The Arrhenius picture is intuitive and works well for simple aqueous reactions, but it fails as soon as the chemistry moves outside water. Ammonia (NH₃) is clearly basic — it reacts with HCl to form NH₄Cl — but contains no OH⁻ ions. Arrhenius cannot explain this. It also fails for gas-phase proton transfer (HCl(g) + NH₃(g) → NH₄Cl(s), the classic "ammonia and concentrated HCl" demonstration that produces a white smoke ring without any solvent present).
Brønsted in Denmark and Lowry in England independently published essentially the same generalisation in the same year. An acid is a proton donor; a base is a proton acceptor. The reaction need not involve water, and no OH⁻ ion need be present. The white smoke in the HCl/NH₃ demonstration is now explained as proton transfer from HCl (acid) to NH₃ (base) forming NH₄⁺Cl⁻ — Arrhenius could not handle this case but Brønsted-Lowry handles it trivially.
This is the definition AQA examines. Throughout this course, "acid" and "base" mean "Brønsted-Lowry acid" and "Brønsted-Lowry base" unless otherwise stated.
In the same year as Brønsted and Lowry, Lewis proposed an even more general definition. A Lewis acid is an electron-pair acceptor; a Lewis base is an electron-pair donor. Every Brønsted-Lowry acid-base reaction is also a Lewis acid-base reaction (H⁺ accepts a lone pair from the base), but the converse is not true: BF₃ + NH₃ → F₃B-NH₃ is a Lewis acid-base reaction with no proton transfer.
Lewis theory is the undergraduate framework. It is essential for transition-metal complex chemistry (ligands are Lewis bases donating lone pairs to metal-ion Lewis acids) and for organic mechanisms (electrophiles are Lewis acids; nucleophiles are Lewis bases). It is not examined at AQA A-Level — but knowing it exists is useful for the "Going Further" extension and for A* synoptic answers.
Key Point: AQA A-Level Chemistry uses the Brønsted-Lowry definition exclusively. Mentioning Lewis acid-base theory in an exam answer earns no marks (and risks confusion if mis-applied), but understanding it conceptually helps with synoptic questions on transition-metal complexes (§3.2.5) and organic mechanisms.
A Brønsted-Lowry acid-base reaction is a proton transfer from the acid to the base. Generic equation:
HA + B ⇌ A⁻ + BH⁺
Here HA is the acid (it donates its proton) and B is the base (it accepts the proton). The products are A⁻ — what HA has become after losing its proton — and BH⁺ — what B has become after gaining a proton.
The species formed when an acid loses its proton is called the conjugate base of that acid. The species formed when a base gains a proton is called the conjugate acid of that base. So in the generic equation:
HA and A⁻ together form a conjugate acid-base pair. So do B and BH⁺. Every Brønsted-Lowry reaction involves exactly two conjugate pairs — and identifying them is one of the most common AO2 exam tasks in this section of the specification.
Consider the reaction of hydrogen chloride dissolving in water:
HCl(aq) + H₂O(l) ⇌ H₃O⁺(aq) + Cl⁻(aq)
The two conjugate pairs are (HCl / Cl⁻) and (H₃O⁺ / H₂O).
NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq)
Now water plays the opposite role.
The two conjugate pairs are (H₂O / OH⁻) and (NH₄⁺ / NH₃).
Notice that water behaves as an acid in one reaction and a base in the other. This dual behaviour is called amphoteric and is the subject of the next section.
At the orbital level, proton transfer is the formation of a new dative covalent bond from a lone pair on the base to the H⁺. In HCl + H₂O, an oxygen lone pair attacks the proton end of the polar H-Cl bond; the H-Cl bond heterolyses (both electrons leave on Cl); and a new O-H bond forms. The arrow-pushing curly arrows used in organic chemistry (§3.3) are the same machinery — Brønsted-Lowry proton transfer is the simplest possible nucleophile-electrophile reaction, with the base as nucleophile and the acidic proton as electrophile.
This is why §3.1.3 (covalent and dative bonding) is prerequisite knowledge: the formation of H₃O⁺ from H⁺ + H₂O is the textbook example of a dative covalent bond, with both bonding electrons supplied by the oxygen.
The following table catalogues conjugate pairs that recur throughout A-Level and into undergraduate chemistry. Memorising the relationship — drop or add one H⁺, adjust the charge by −1 or +1 — makes identifying conjugate species mechanical.
| Acid | Conjugate base | Notes |
|---|---|---|
| HCl | Cl⁻ | Strong acid; fully dissociated in water |
| HNO₃ | NO₃⁻ | Strong acid |
| H₂SO₄ | HSO₄⁻ | First ionisation; strong |
| HSO₄⁻ | SO₄²⁻ | Second ionisation; weak (Ka₂ ≈ 1.2 × 10⁻²) |
| CH₃COOH | CH₃COO⁻ | Weak acid; Ka ≈ 1.8 × 10⁻⁵ |
| HF | F⁻ | Weak acid in water (Ka ≈ 6.6 × 10⁻⁴) |
| HCN | CN⁻ | Very weak acid (Ka ≈ 6.2 × 10⁻¹⁰) |
| NH₄⁺ | NH₃ | Conjugate acid of ammonia |
| H₃O⁺ | H₂O | Conjugate acid of water |
| H₂O | OH⁻ | Conjugate base of water |
| H₂CO₃ | HCO₃⁻ | Carbonic acid (mostly CO₂(aq); Ka₁ ≈ 4.5 × 10⁻⁷) |
| HCO₃⁻ | CO₃²⁻ | Hydrogencarbonate; Ka₂ ≈ 4.7 × 10⁻¹¹ |
Exam Tip: When asked to write a conjugate species, change the formula by exactly one H and one unit of charge — never more, never less. The conjugate base of H₂SO₄ is HSO₄⁻, not SO₄²⁻ (that would be losing two protons). The conjugate acid of CO₃²⁻ is HCO₃⁻, not H₂CO₃.
A useful sanity check: charge difference between an acid and its conjugate base is always exactly +1 (the acid is +1 higher because it carries the extra H⁺). So:
If your candidate conjugate species violates this rule, you have made an arithmetic error.
A species that can act as either an acid or a base is called amphoteric (or amphiprotic when the dual behaviour specifically involves donating and accepting protons). Water is the canonical example.
With HCl, water accepts a proton from the strong acid:
HCl + H₂O → H₃O⁺ + Cl⁻ (water is the base, H₃O⁺ is its conjugate acid)
With NH₃, water donates a proton to the base:
H₂O + NH₃ ⇌ OH⁻ + NH₄⁺ (water is the acid, OH⁻ is its conjugate base)
Because water is amphoteric, two water molecules can react with each other — one donating a proton, one accepting. This is the autoionisation (or "self-ionisation") of water:
2H₂O(l) ⇌ H₃O⁺(aq) + OH⁻(aq)
(Equivalently, dropping the explicit hydration of H⁺: H₂O ⇌ H⁺ + OH⁻.)
The equilibrium lies far to the left — only about 1 in 5.5 × 10⁸ water molecules is dissociated at any moment — but it is the reason that pure water has [H⁺] ≠ 0. The equilibrium constant for this reaction, after the (constant) activity of water has been absorbed into the constant, is the ionic product of water, Kw:
Kw = [H₃O⁺][OH⁻] = 1.00 × 10⁻¹⁴ mol² dm⁻⁶ at 298 K
This relation underlies every pH calculation in the next two lessons. In pure water [H₃O⁺] = [OH⁻] = 1.00 × 10⁻⁷ mol dm⁻³, giving pH = 7. The value of Kw is temperature-dependent (it doubles roughly every 10 K) — the dissociation is endothermic — so neutral pH is exactly 7 only at 298 K. At body temperature (310 K) neutral pH is about 6.81; biological "neutrality" is therefore not pH 7 but a slightly lower value, with arterial blood maintained at 7.40 by the carbonate buffer.
Common Misconception: Students sometimes assume "neutral" always means "pH 7". Strictly, neutral means [H⁺] = [OH⁻], which equals 10⁻⁷ only at 298 K. At higher temperatures, neutral pH is lower; at lower temperatures, higher. AQA expects the value Kw = 1.00 × 10⁻¹⁴ at 298 K unless otherwise stated.
The following examples show the standard exam format. In each case, identify the acid, the base, and the two conjugate acid-base pairs.
HNO₃(aq) + H₂O(l) → H₃O⁺(aq) + NO₃⁻(aq)
CH₃COOH(aq) + H₂O(l) ⇌ H₃O⁺(aq) + CH₃COO⁻(aq)
HCO₃⁻(aq) + H₂O(l) ⇌ H₂CO₃(aq) + OH⁻(aq)
HCO₃⁻(aq) + H₂O(l) ⇌ CO₃²⁻(aq) + H₃O⁺(aq)
An acid that can donate more than one proton per molecule is called polyprotic (or specifically diprotic for two protons, triprotic for three). The protons leave sequentially, and each ionisation has its own acid dissociation constant Ka.
Step 1: H₂SO₄(aq) + H₂O(l) → H₃O⁺(aq) + HSO₄⁻(aq) (Ka₁ very large — strong acid; fully ionised)
Step 2: HSO₄⁻(aq) + H₂O(l) ⇌ H₃O⁺(aq) + SO₄²⁻(aq) (Ka₂ ≈ 1.2 × 10⁻², pKa₂ ≈ 1.92 — only moderately strong)
The first proton comes off completely; the second is held more tightly because removing a positive charge (H⁺) from a species that is already negative (HSO₄⁻) is harder than removing it from a neutral molecule (H₂SO₄). This is a general pattern: for any polyprotic acid, Ka₁ > Ka₂ > Ka₃, typically by 4-6 orders of magnitude per step.
Step 1: H₃PO₄ + H₂O ⇌ H₃O⁺ + H₂PO₄⁻ (pKa₁ ≈ 2.15)
Step 2: H₂PO₄⁻ + H₂O ⇌ H₃O⁺ + HPO₄²⁻ (pKa₂ ≈ 7.20)
Step 3: HPO₄²⁻ + H₂O ⇌ H₃O⁺ + PO₄³⁻ (pKa₃ ≈ 12.35)
The five-decade spacing of the pKa values explains why phosphate is such a useful biological buffer: in the physiological pH range around 7.4, the H₂PO₄⁻/HPO₄²⁻ pair is roughly half-protonated and resists pH change. The same logic underlies the carbonate buffer in blood.
Step 1: H₂CO₃ + H₂O ⇌ H₃O⁺ + HCO₃⁻ (pKa₁ ≈ 6.35)
Step 2: HCO₃⁻ + H₂O ⇌ H₃O⁺ + CO₃²⁻ (pKa₂ ≈ 10.33)
H₂CO₃ itself is a transient species — in aqueous solution it equilibrates rapidly with dissolved CO₂: CO₂(aq) + H₂O ⇌ H₂CO₃. The combined system CO₂/HCO₃⁻ is the principal blood-pH buffer and is treated quantitatively in the buffers lesson later in this course.
This is the single most-tested conceptual distinction in the Brønsted-Lowry section of the AQA specification. Strength refers to the extent of ionisation; concentration refers to the number of moles of acid dissolved per unit volume. The two are independent.
A 0.10 mol dm⁻³ solution of HCl (strong) and a 0.10 mol dm⁻³ solution of CH₃COOH (weak) have the same concentration but very different pH values: pH 1.0 for the HCl and pH 2.87 for the ethanoic acid. Conversely, a 0.0010 mol dm⁻³ solution of HCl and a 0.10 mol dm⁻³ solution of CH₃COOH happen to have similar pH (both about 3) — same pH, different strengths, different concentrations.
Common Misconception: "Concentrated" and "strong" are not synonyms; "dilute" and "weak" are not synonyms. A 12 mol dm⁻³ solution of ethanoic acid is concentrated but still weak. A 10⁻⁵ mol dm⁻³ solution of HCl is dilute but still strong. Confusing these is the most-marked-down error in Paper 2 acid-base questions.
The same distinction applies to bases.
The hydrogencarbonate ion HCO₃⁻ is amphoteric in the same way water is. It is the middle species in the H₂CO₃ → HCO₃⁻ → CO₃²⁻ ladder, so it can either accept a proton (becoming H₂CO₃) or donate one (becoming CO₃²⁻).
HCO₃⁻ acting as a base: HCO₃⁻ + H⁺ → H₂CO₃ (which rapidly converts to CO₂ + H₂O — this is why dropping a vitamin C tablet into bicarbonate solution fizzes).
HCO₃⁻ acting as an acid: HCO₃⁻ + OH⁻ → CO₃²⁻ + H₂O.
In the blood, the equilibrium CO₂(aq) + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻ holds the arterial pH at 7.40 ± 0.05 against metabolic acid load. If too much CO₂ is dissolved (respiratory acidosis), the kidneys reabsorb more bicarbonate; if too little (respiratory alkalosis from hyperventilation), the kidneys excrete more. This is the standard physiology example of an amphoteric buffer and may appear synoptically in §3.1.12.5 (buffers) and indirectly in biology cross-curricular questions.
Other common amphoteric species you should recognise: H₂O (the canonical case), HSO₄⁻ (donates H⁺ to give SO₄²⁻; accepts to give H₂SO₄), H₂PO₄⁻ and HPO₄²⁻ (the middle members of the phosphate ladder), and amino acids in zwitterionic form (NH₃⁺-CHR-COO⁻; treated in §3.3.13 organic).
A simple practical exercise illustrates the Brønsted-Lowry framework directly. Prepare five 0.10 mol dm⁻³ solutions: HCl, CH₃COOH, NH₃, NaOH, and NaHCO₃. Measure the pH of each using a calibrated pH meter or universal indicator paper.
| Solution | Expected pH (0.10 mol dm⁻³) | Brønsted-Lowry interpretation |
|---|---|---|
| HCl | 1.0 | Strong acid; fully donates H⁺ to water |
| CH₃COOH | 2.87 | Weak acid; partially donates H⁺ |
| NaHCO₃ | ~8.3 | Amphoteric anion; slightly basic in water |
| NH₃ | 11.13 | Weak base; partially accepts H⁺ from water |
| NaOH | 13.0 | Strong base; fully provides OH⁻ |
Comparing the HCl and CH₃COOH solutions at identical concentration (pH 1.0 vs 2.87) directly demonstrates the strength/concentration distinction. Comparing NH₃ and NaOH at identical concentration (pH 11.13 vs 13.0) does the same on the basic side. This is the recommended Brønsted-Lowry demonstration in most A-Level practical handbooks.
Practical Tip: Always calibrate the pH meter using standard buffer solutions (typically pH 4.00 and pH 7.00 for acids; pH 7.00 and pH 10.00 for bases) immediately before measurement, and rinse the electrode with deionised water between samples. Stale calibration is the largest source of systematic error in pH measurements.
The Brønsted-Lowry framework threads through almost every remaining topic in A-Level Chemistry:
Question 1. [12 marks total]
(a) Define a Brønsted-Lowry acid and a Brønsted-Lowry base. [2 marks]
(b) For each of the following reactions, identify the Brønsted-Lowry acid and base on the left-hand side, and write down the two conjugate acid-base pairs.
(i) HBr(aq) + H₂O(l) → H₃O⁺(aq) + Br⁻(aq)
(ii) NH₃(aq) + CH₃COOH(aq) ⇌ NH₄⁺(aq) + CH₃COO⁻(aq)
(iii) HSO₄⁻(aq) + OH⁻(aq) → SO₄²⁻(aq) + H₂O(l)
[4 marks total]
(c) Explain, with reference to two specific reactions, why water is described as amphoteric. [3 marks]
(d) Sulfuric acid is described as a much stronger acid than the hydrogensulfate ion HSO₄⁻. Explain this difference in terms of the structures of the species involved and the energetics of proton loss. [3 marks]
(a) Brønsted-Lowry definitions [2 marks, AO1]
Award full marks for clear, complete definitions. Lose a mark for stating "produces H⁺" without "donates" (that's the Arrhenius definition). Lose a mark if "H⁺" is replaced with "H atom" or "H₂".
(b) Identifying acids, bases, and conjugate pairs [4 marks, AO2]
(i) [1 mark]: acid HBr; base H₂O; conjugate pairs (HBr / Br⁻) and (H₃O⁺ / H₂O). All three elements required for the mark.
(ii) [1 mark]: acid CH₃COOH; base NH₃; conjugate pairs (CH₃COOH / CH₃COO⁻) and (NH₄⁺ / NH₃).
(iii) [1 mark]: acid HSO₄⁻; base OH⁻; conjugate pairs (HSO₄⁻ / SO₄²⁻) and (H₂O / OH⁻).
[1 mark]: across all three, conjugate pairs are written with the acid on the left and the base on the right (or labelled). Marker discretion — if all three are correct, award.
Common error: students sometimes label H₂O as an acid in (i) because it appears on the left. Read the question — water accepts a proton in (i), so it is the base.
(c) Water amphoteric [3 marks, AO3]
Either reaction in either role can be drawn from the autoionisation 2H₂O ⇌ H₃O⁺ + OH⁻ (where one water acts as acid and the other as base). Accept this single example for 2 of the 3 marks provided both roles are explicitly identified.
(d) H₂SO₄ vs HSO₄⁻ [3 marks, AO3]
Award for any clear electrostatic reasoning. Do not require numerical Ka values.
The three responses below cover the meaningful A-Level range: Grade C (borderline pass), Grade B (solid mark-scheme coverage), and Grade A* (top-band synoptic insight). No D or E responses are shown — no A-Level student should be aiming for those bands, and modelling failure adds nothing pedagogically. The editorial commentary after each response names the marks earned and the moves that differentiate from the adjacent bands.
(a) A Brønsted-Lowry acid is a proton donor — it gives away an H⁺ ion. A Brønsted-Lowry base is a proton acceptor — it takes an H⁺ ion.
(b) (i) Acid HBr, base H₂O. Conjugate pairs: HBr / Br⁻ and H₃O⁺ / H₂O.
(ii) Acid CH₃COOH, base NH₃. Conjugate pairs: CH₃COOH / CH₃COO⁻ and NH₄⁺ / NH₃.
(iii) Acid HSO₄⁻, base OH⁻. Conjugate pairs: HSO₄⁻ / SO₄²⁻ and H₂O / OH⁻.
(c) Water is amphoteric because it can act as both an acid and a base. With HCl it acts as a base because it accepts a proton: HCl + H₂O → H₃O⁺ + Cl⁻. With NH₃ it acts as an acid because it donates a proton: H₂O + NH₃ → OH⁻ + NH₄⁺.
(d) H₂SO₄ is a stronger acid than HSO₄⁻ because H₂SO₄ is neutral and HSO₄⁻ is already negatively charged. It is harder to take an H⁺ away from a negative ion (HSO₄⁻) than from a neutral molecule (H₂SO₄), because the H⁺ is more strongly attracted to the negative charge that is being left behind. So the second proton is harder to remove than the first, which is why HSO₄⁻ is a weaker acid.
Editorial commentary (Grade C): Hits all the basic mark-scheme points and would secure a borderline C. The conjugate pair identifications in (b) are clean. Part (c) names amphoteric and gives both roles with reactions — full marks. Part (d) gets the electrostatic reasoning correct but does not quantify with Ka values or generalise to other polyprotic acids. To progress to B, the answer should cite Ka₁ and Ka₂ values, generalise that Ka₁ > Ka₂ for any polyprotic acid, and link the structural reasoning to bond polarity or charge delocalisation in the conjugate base.
(a) A Brønsted-Lowry acid is a species that donates a proton (an H⁺ ion). A Brønsted-Lowry base is a species that accepts a proton. The Brønsted-Lowry definition (1923) supersedes the older Arrhenius definition because it does not require water as a solvent, and it explains gas-phase reactions such as HCl(g) + NH₃(g) → NH₄Cl(s).
(b) (i) Acid HBr, base H₂O; conjugate pairs (HBr / Br⁻) and (H₃O⁺ / H₂O). (ii) Acid CH₃COOH, base NH₃; conjugate pairs (CH₃COOH / CH₃COO⁻) and (NH₄⁺ / NH₃). (iii) Acid HSO₄⁻, base OH⁻; conjugate pairs (HSO₄⁻ / SO₄²⁻) and (H₂O / OH⁻).
(c) An amphoteric species can act as both a Brønsted-Lowry acid and a Brønsted-Lowry base, depending on its reaction partner. Water with HCl acts as a base (H₂O + HCl → H₃O⁺ + Cl⁻ — water accepts H⁺). Water with NH₃ acts as an acid (H₂O + NH₃ ⇌ OH⁻ + NH₄⁺ — water donates H⁺). The autoionisation 2H₂O ⇌ H₃O⁺ + OH⁻ shows both roles simultaneously, with Kw = [H₃O⁺][OH⁻] = 1.00 × 10⁻¹⁴ at 298 K.
(d) H₂SO₄ loses H⁺ from a neutral molecule (charge changes 0 → −1), giving HSO₄⁻. HSO₄⁻ then has to lose H⁺ from an already-negative species (charge changes −1 → −2), which is electrostatically much less favourable: the leaving H⁺ is more strongly attracted back to the higher negative charge on SO₄²⁻. Quantitatively, Ka₁ of H₂SO₄ is very large (effectively infinite for A-Level — fully ionised), while Ka₂ ≈ 1.2 × 10⁻², so pKa₂ ≈ 1.92. This is the general pattern for any polyprotic acid: Ka₁ ≫ Ka₂ ≫ Ka₃, typically by 4–6 orders of magnitude per ionisation.
Editorial commentary (Grade B): Full mark-scheme coverage with appropriate detail. The reference to Arrhenius vs Brønsted-Lowry in (a) is gracious. Part (c) adds the autoionisation equation and Kw — a clean synoptic move into the next lesson. Part (d) quotes Ka and pKa values and generalises to other polyprotic acids. To reach A*, the response should link the conjugate-strength inverse relationship explicitly (a stronger acid has a weaker conjugate base, with Ka × Kb = Kw at 298 K), or comment on charge delocalisation in SO₄²⁻ as the deeper structural explanation, or flag the Lewis-vs-Brønsted distinction for complex-ion equilibria.
(a) A Brønsted-Lowry acid is a proton (H⁺) donor; a Brønsted-Lowry base is a proton acceptor. The 1923 Brønsted-Lowry definition generalises the earlier Arrhenius (1884) H⁺/OH⁻-in-water picture: it removes the requirement for an aqueous solvent and reduces neutralisation to a single mechanistic primitive — proton transfer. (Lewis, also 1923, generalises further to electron-pair acceptor / donor; this is not on the AQA A-Level specification but is essential for d-block ligand chemistry and is the basis for the curly-arrow mechanism in §3.3.)
(b) (i) Acid HBr; base H₂O; pairs (HBr / Br⁻) and (H₃O⁺ / H₂O). (ii) Acid CH₃COOH; base NH₃; pairs (CH₃COOH / CH₃COO⁻) and (NH₄⁺ / NH₃). (iii) Acid HSO₄⁻; base OH⁻; pairs (HSO₄⁻ / SO₄²⁻) and (H₂O / OH⁻).
(c) Amphoteric = both acid and base. Water with HCl: H₂O + HCl → H₃O⁺ + Cl⁻ (water is the base). Water with NH₃: H₂O + NH₃ ⇌ OH⁻ + NH₄⁺ (water is the acid). The autoionisation 2H₂O ⇌ H₃O⁺ + OH⁻ shows both roles in a single reaction with Kw = [H₃O⁺][OH⁻] = 1.00 × 10⁻¹⁴ mol² dm⁻⁶ at 298 K (Kw increases with temperature, so the autoionisation is endothermic — ΔH ≈ +55.8 kJ mol⁻¹ — and the "neutral" pH falls below 7 at body temperature).
(d) For any conjugate pair, Ka × Kb = Kw at 298 K — so a stronger acid must have a weaker conjugate base. Ka₁(H₂SO₄) is so large that H₂SO₄ is treated as fully ionised; the conjugate base HSO₄⁻ is correspondingly weakly basic. Ka₂(HSO₄⁻) ≈ 1.2 × 10⁻² (pKa₂ ≈ 1.92); the conjugate base SO₄²⁻ is therefore more basic than HSO₄⁻. Structurally, removing H⁺ from neutral H₂SO₄ produces HSO₄⁻ (single −1 charge, delocalised over three oxygens); removing H⁺ from already-negative HSO₄⁻ produces SO₄²⁻ (double −2 charge, delocalised over four oxygens). The electrostatic penalty for the second ionisation arises because H⁺ must be separated from a higher negative charge density; partial offset by the increased delocalisation in SO₄²⁻ (four equivalent S-O bonds of bond order 1.5) is not enough to overcome the electrostatic cost. This is the general pattern: each successive Ka in a polyprotic acid is ~10⁴–10⁶ times smaller than the previous one.
Editorial commentary (Grade A):* Genuinely A* throughout. Part (a) cites Arrhenius and Lewis without quoting and frames Brønsted-Lowry as the mechanistic primitive. Part (c) adds the temperature dependence of Kw and the enthalpy of autoionisation — a clean synoptic move into A2 thermodynamics. Part (d) introduces the Ka × Kb = Kw relation (the conjugate-strength inverse — the key A* insight), gives the structural argument with delocalisation in SO₄²⁻ and HSO₄⁻, and generalises quantitatively to all polyprotic acids. This is the calibre of answer that distinguishes the top 8% of the cohort.
Three undergraduate-adjacent extensions worth knowing about, none examined at A-Level:
Brønsted-Lowry proton transfer is the mechanistic primitive of the next several lessons in this course: pH calculations, Ka and Kb, buffers, indicators, and titration curves all reduce to identifying which species is donating H⁺ and which is accepting. Master conjugate pairs, the strength/concentration distinction, and the amphoteric character of water here, and the rest of §3.1.12 unfolds as a quantitative elaboration of a single idea.