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Spec Mapping — OCR H432 Module 5.1.3 — Acids, bases and buffers, opening sub-topic on the Brønsted–Lowry definition: acids as proton donors, bases as proton acceptors, conjugate acid–base pairs, monobasic / dibasic / tribasic acids, the amphoteric behaviour of water, and the proton-transfer framework that underpins every quantitative calculation in the rest of the module (refer to the official OCR H432 specification document for exact wording). This is the conceptual gateway lesson — every subsequent topic (Kw, strong-acid pH, Ka, weak-acid pH, buffers, titration curves, indicators) is built on the Brønsted–Lowry vocabulary established here.
The Brønsted–Lowry framework is deceptively easy to state and notoriously easy to misapply in exam scripts. Year-13 candidates often arrive at this lesson confident that they understand acids from GCSE — "they're sour, they turn litmus red, they react with metals to make hydrogen" — and then lose marks because the OCR mark scheme rewards the operational definition (acid = proton donor; base = proton acceptor) and the structural consequence (every acid-base reaction generates two conjugate pairs). This lesson sets up that vocabulary precisely so that the quantitative work in lessons 2–10 has a sure foundation. The proton-transfer logic introduced here is also the bridge to the redox half-equation logic from the ocr-alevel-chemistry-acids-redox-bonding course — both topics ask you to track a single transferred particle (proton or electron) between two reacting species.
Key Definitions:
- Acid — a species that donates a proton (H⁺) to another species (Brønsted–Lowry, 1923).
- Base — a species that accepts a proton (Brønsted–Lowry).
- Conjugate acid–base pair — two species differing by exactly one H⁺.
- Amphoteric / amphiprotic — a species that can act as either acid or base depending on its partner.
- Basicity (of an acid) — the number of replaceable acidic protons per molecule (1 = monobasic / monoprotic, 2 = dibasic / diprotic, 3 = tribasic / triprotic).
Before 1923 the Arrhenius definition (acids release H⁺ in water; bases release OH⁻ in water) was the operational standard. It worked well for textbook examples like HCl, HNO₃ and NaOH, but it had two serious deficiencies that OCR explicitly examines:
The Danish chemist Brønsted and the English chemist Lowry independently proposed the proton-transfer definition in 1923. Their reformulation discards the requirement that water be the solvent. An acid is anything that donates H⁺; a base is anything that accepts H⁺. Every acid-base reaction is now a single proton-transfer event:
HA+B⇌A−+BH+
The acid HA loses a proton to become its conjugate base A⁻. The base B gains the proton to become its conjugate acid BH⁺. The equilibrium arrow is mandatory in the general case because the products themselves are an acid (BH⁺) and a base (A⁻) and can react in reverse. The position of equilibrium is determined by the relative strengths of the two acids on either side — equilibrium lies on the side of the weaker acid (and weaker base).
A still more general definition was proposed by G. N. Lewis the same year (acid = electron-pair acceptor; base = electron-pair donor). Lewis acids include species with no proton to donate, such as BF₃ accepting the lone pair of NH₃ to form an adduct. At OCR A-Level the Lewis definition appears at the fringes — in dative covalent bonding and transition-metal complex formation — but the operational definition for Module 5.1.3 is Brønsted–Lowry throughout.
When hydrogen chloride dissolves in water the equilibrium
HCl(aq)+H2O(l)→H3O+(aq)+Cl−(aq)
is set up. For HCl the equilibrium lies far to the right (essentially complete dissociation), which is why we use a single forward arrow.
The product H₃O⁺ is the oxonium ion (sometimes called hydronium). It is the species that gives acidic solutions their characteristic chemistry — bare H⁺ does not exist free in water, because the electric field at the surface of a proton is enormous and it is immediately solvated by surrounding water molecules. At A-Level we routinely abbreviate H₃O⁺(aq) to H⁺(aq), but you should recognise the abbreviation for what it is: the proton in water is always at minimum H₃O⁺, with several additional waters loosely associated in a hydration shell.
NH3(aq)+H2O(l)⇌NH4+(aq)+OH−(aq)
Ammonia is a weak base, so this equilibrium lies far to the left — most NH₃ molecules in solution remain un-protonated, and the OH⁻ concentration is much lower than the total ammonia concentration. Notice that the same molecule (water) acted as base in worked example 1 and as acid in worked example 2. This dual behaviour is called amphoteric (general) or amphiprotic (specifically when proton transfer is what flips the role).
Identify the conjugate acid–base pairs in:
HSO4−(aq)+H2O(l)⇌SO42−(aq)+H3O+(aq)
Solution: HSO₄⁻ loses a proton to become SO₄²⁻ — HSO₄⁻ is acid 1, SO₄²⁻ is its conjugate base. H₂O gains a proton to become H₃O⁺ — H₂O is base 2, H₃O⁺ is its conjugate acid.
Pair 1: HSO₄⁻ / SO₄²⁻. Pair 2: H₂O / H₃O⁺.
The interesting subtlety: HSO₄⁻ itself is the conjugate base of H₂SO₄ in the first dissociation step (H₂SO₄ → H⁺ + HSO₄⁻), yet here it is acting as an acid. The same species can be an acid in one reaction and a base in another — what it does depends on the partner, not on the species in isolation. Species capable of either role are called amphiprotic.
Every Brønsted–Lowry reaction contains two conjugate pairs. To identify them:
graph LR
A["HA<br/>(acid 1)"] -->|"loses H+"| B["A−<br/>(conjugate base 1)"]
C["B<br/>(base 2)"] -->|"gains H+"| D["BH+<br/>(conjugate acid 2)"]
The graphic above is the universal Brønsted–Lowry template. Every proton-transfer reaction you ever write at A-Level fits this template.
Water can act as either acid or base depending on its partner. With a stronger acid (HCl) water behaves as a base; with a base (NH₃) water behaves as an acid:
| Partner | Reaction | Water's role |
|---|---|---|
| HCl | H2O+HCl→H3O++Cl− | base |
| NH₃ | H2O+NH3⇌OH−+NH4+ | acid |
| Itself | 2H2O⇌H3O++OH− | both! |
The third row is the self-ionisation of water, which is the topic of lesson 2 (Kw). Water is the only common molecule that ionises itself — and the equilibrium constant for that self-ionisation gives us pH, the central observable of acid-base chemistry.
Amphoteric behaviour also explains why some inorganic oxides react with both acids and alkalis. Aluminium oxide is a classic example: Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O and Al₂O₃ + 2NaOH + 3H₂O → 2NaAl(OH)₄. We classify Al₂O₃ as an amphoteric oxide, although strictly it is not Brønsted-amphiprotic (it does not transfer a proton; it reacts ionically). This subtle distinction — amphoteric (general) vs amphiprotic (proton-specific) — is one of the A-vs-A* discriminators in OCR mark schemes.
Some acids can donate more than one proton per molecule. The number of replaceable protons is the acid's basicity (sometimes called proticity). A common A-Level error is to count all hydrogens in the molecular formula instead of only the acidic ones.
| Acid | Formula | Basicity | Ionisation steps | Notes |
|---|---|---|---|---|
| Hydrochloric | HCl | 1 (monobasic) | HCl→H++Cl− | Strong; one-step |
| Nitric | HNO₃ | 1 (monobasic) | HNO3→H++NO3− | Strong; one-step |
| Ethanoic | CH₃COOH | 1 (monobasic) | CH3COOH⇌H++CH3COO− | Weak; only the COOH hydrogen is acidic |
| Sulfuric | H₂SO₄ | 2 (dibasic) | H2SO4→H++HSO4−; HSO4−⇌H++SO42− | First strong; second weak |
| Carbonic | H₂CO₃ | 2 (dibasic) | H2CO3⇌H++HCO3−; HCO3−⇌H++CO32− | Both weak |
| Phosphoric(V) | H₃PO₄ | 3 (tribasic) | Three successive ionisations | All weak; pKa≈2.1,7.2,12.4 |
The acidic-hydrogen rule: a hydrogen atom is acidic when it is bonded to a strongly electronegative atom (O, halogen, sometimes N) such that the bond is polarised enough for H⁺ to leave in water. In ethanoic acid CH₃COOH the molecule has four hydrogens but only one is acidic — the one bonded to oxygen. The three methyl hydrogens are bonded to carbon (electronegativity 2.5 vs hydrogen 2.2 — barely polar) and cannot leave under aqueous conditions. So ethanoic acid is monobasic, not tetrabasic.
The pattern extends to other carboxylic acids: butanedioic acid HOOC–CH₂–CH₂–COOH is dibasic (two COOH groups); citric acid is tribasic; ethylenediaminetetraacetic acid (EDTA) is tetrabasic. The classification reflects the number of O–H protons, not the total H count.
All the familiar acid reactions you met at GCSE can now be reframed as proton-transfer events:
Mg(s)+2HCl(aq)→MgCl2(aq)+H2(g)
The metal here is not strictly a Brønsted–Lowry base — this is a redox reaction (Mg is oxidised; H⁺ is reduced). But the acid is still behaving as a proton donor at the metal surface, and the rate depends on [H⁺].
HCl(aq)+NaOH(aq)→NaCl(aq)+H2O(l)
Stripping the spectator Na⁺ and Cl⁻:
H+(aq)+OH−(aq)→H2O(l),ΔHneut≈−57 kJ mol−1
The hydroxide ion accepts a proton from H₃O⁺ (or from H⁺) to form water. This is the archetypal Brønsted neutralisation, and explains why the enthalpy of neutralisation for any strong acid + strong alkali pair is the same value — the underlying reaction is identical.
2HCl(aq)+Na2CO3(aq)→2NaCl(aq)+H2O(l)+CO2(g)
The carbonate ion is the base; it accepts two protons in succession to form H₂CO₃, which decomposes rapidly to H₂O and CO₂, hence the characteristic effervescence. Net ionic: CO32−+2H+→H2O+CO2.
HCl(aq)+NH3(aq)→NH4Cl(aq)
Net ionic: H+(aq)+NH3(aq)→NH4+(aq). NH₃ is the base.
In every case one species donates a proton and another accepts it. Identifying the pairs is the first step to writing any acid-base equation correctly.
Identify the conjugate pairs in the hydrolysis of the ammonium ion:
NH4+(aq)+H2O(l)⇌H3O+(aq)+NH3(aq)
Solution: NH₄⁺ donates a proton → NH₄⁺ is the acid; NH₃ is its conjugate base. H₂O accepts a proton → H₂O is the base; H₃O⁺ is its conjugate acid. Pairs: NH₄⁺ / NH₃ and H₂O / H₃O⁺. Notice this is the exact reverse role-assignment from worked example 2 (NH₃ + H₂O): here the ammonium ion acts as an acid because it has a proton to donate, whereas in worked example 2 ammonia acted as a base because it had a lone pair to accept one. The equilibrium for NH₄⁺ + H₂O lies far to the left (NH₄⁺ is only a very weak acid, Ka≈5.6×10−10), which is why a solution of pure ammonium chloride is only mildly acidic.
Write the first-step Brønsted–Lowry equation for the reaction between carbonate ion and ethanoic acid:
CH3COOH(aq)+CO32−(aq)⇌CH3COO−(aq)+HCO3−(aq)
Here CH₃COOH is acid 1 (donates H⁺) and CO₃²⁻ is base 2 (accepts H⁺). The conjugate pairs are CH₃COOH / CH₃COO⁻ and HCO₃⁻ / CO₃²⁻. A second proton transfer can then produce H₂CO₃, which decomposes to CO₂ and H₂O — hence the familiar effervescence observed when ethanoic acid is added to sodium carbonate.
Synoptic Links — Connects to:
ocr-alevel-chemistry-acids-redox-bonding / acids-and-bases(the Year-12 introduction to Brønsted–Lowry — this lesson is the Year-13 quantitative continuation).ocr-alevel-chemistry-acids-redox-bonding / acid-base-titrations(the titration calculations from PAG 2 are extended in lesson 9 to construct full pH curves).ocr-alevel-chemistry-quantitative-rates-equilibrium(the equilibrium concept and Kc treatment underpin every Ka, Kw and Henderson–Hasselbalch calculation in lessons 2–8).ocr-alevel-chemistry-acids-bases-buffers / ionic-product-of-water-and-ph(the immediate next step — Kw — applies the self-ionisation idea quantitatively).
Practical Activity Group anchor: PAG 2 — Acid–base titration. Although the volumetric calculations are the focus of PAG 2 (covered in the prerequisite acid-base-titrations lesson), the Brønsted–Lowry vocabulary established here is the language used in every PAG 2 write-up — "neutralisation is proton transfer from the strong acid in the burette to the weak base in the conical flask," etc. PAG 11 — pH measurement also uses this framework, as the pH meter responds to [H₃O⁺] which is generated by proton donation from the analyte.
Question (6 marks): Aqueous ammonia is a weak base. (a) Define the term Brønsted–Lowry base. (b) Write a balanced equation, including state symbols and an appropriate arrow, for the reaction of ammonia with water. (c) Identify the two conjugate acid–base pairs in this reaction. (d) Explain why aqueous ammonia is described as a "weak" base in terms of the position of equilibrium.
| Mark | AO | Awarded for |
|---|---|---|
| 1 | AO1 | Brønsted–Lowry base = proton acceptor |
| 2 | AO1 | Balanced equation NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq) |
| 3 | AO1 | Correct use of equilibrium arrow (⇌) and state symbols |
| 4 | AO2 | Pair 1 identified: NH₃ / NH₄⁺ (base / conjugate acid) |
| 5 | AO2 | Pair 2 identified: H₂O / OH⁻ (acid / conjugate base) |
| 6 | AO3 | "Weak" justified — equilibrium lies far to the left; only a small fraction of NH₃ molecules are protonated at any instant |
AO split: AO1 = 3, AO2 = 2, AO3 = 1.
(a) A Brønsted–Lowry base accepts a proton.
(b) NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq).
(c) NH₃ is the base and NH₄⁺ is its conjugate acid. H₂O is the acid and OH⁻ is its conjugate base.
(d) Ammonia is a weak base because not all of it reacts with water — only some forms NH₄⁺ and OH⁻.
Examiner commentary: The next-band move is precision. M1 is awarded for the definition. M1 for the balanced equation. M1 for the equilibrium arrow with state symbols. M1 for pair 1, M1 for pair 2. Part (d) loses a mark because "only some" is too vague — the mark scheme expects an explicit statement that the equilibrium lies far to the left so the equilibrium concentrations of NH₄⁺ and OH⁻ are much smaller than the original ammonia concentration. The candidate also misses the synoptic link to the position-of-equilibrium reasoning from quantitative-rates-equilibrium. To lift to top band, add the equilibrium-arrow logic explicitly (left-leaning ⇌) and quantify "small fraction" (around 1 % for 0.1 mol dm⁻³ aqueous ammonia at 298 K).
(a) A Brønsted–Lowry base is a species that accepts a proton (H⁺) from another species.
(b) NH3(aq)+H2O(l)⇌NH4+(aq)+OH−(aq) — equilibrium arrow because the reaction is reversible and partial.
(c) Pair 1: NH₃ (base) and NH₄⁺ (its conjugate acid) — they differ by a single H⁺. Pair 2: H₂O (acid) and OH⁻ (its conjugate base) — they also differ by a single H⁺.
(d) Aqueous ammonia is a weak base because the equilibrium lies far to the left. At 298 K and 0.1 mol dm⁻³, only about 1 % of NH₃ molecules accept a proton from water at any instant; the bulk of the dissolved ammonia remains as un-protonated NH₃. This is a position-of-equilibrium statement: Kb for ammonia is small (≈1.8×10−5 mol dm⁻³), so [OH−] at equilibrium is much smaller than [NH3]initial.
Examiner commentary: Full 6/6. The candidate states the definition precisely, balances the equation with the correct arrow and state symbols, identifies both pairs with the "differ by one H⁺" rule made explicit, and grounds the "weak" descriptor in the position-of-equilibrium framework. The discriminators that lift this to A*: quantifying "about 1 %" (showing the answer is calibrated, not vague); naming Kb even though the formal definition does not arrive until lesson 4; and the synoptic gesture to "position of equilibrium" — language the candidate carries forward from quantitative-rates-equilibrium. The neat structuring (a), (b), (c), (d) also signals exam discipline.
Pedagogical observations — not fabricated statistics:
The Brønsted–Lowry framework was extended in 1923 by G. N. Lewis to acids as electron-pair acceptors and bases as electron-pair donors, generalising acid-base chemistry to species with no proton at all. Lewis acid-base chemistry dominates inorganic chemistry beyond A-Level: BF₃ + NH₃ → F₃B–NH₃ (an adduct, no proton transferred); transition-metal Lewis acidity drives the formation of all classical coordination complexes ([Fe(H₂O)₆]³⁺, [Cu(NH₃)₄]²⁺); the recent Frustrated Lewis Pair catalysis (Stephan and Erker, 2006 onwards) uses sterically hindered Lewis acid-base pairs to activate H₂ and CO₂ at room temperature — a major topic in green-chemistry research. Recommended reading: Atkins, Physical Chemistry, chapter on chemical equilibrium; Housecroft & Sharpe, Inorganic Chemistry, on superacid behaviour (the Magic Acid HF·SbF₅ system has Hammett acidity H0≈−25, strong enough to protonate methane). Oxbridge interview-style prompt: "Carborane acids are stronger than fluoroantimonic acid but are described as 'gentle' — how can a superacid be gentle?" The answer hinges on the conjugate base being extraordinarily non-nucleophilic, so the acid donates a proton vigorously but the conjugate base does nothing further. Strength and reactivity are not the same property.
The errors that distinguish A from A*:
The Brønsted–Lowry framework defines acid-base chemistry by proton transfer: acids donate H⁺, bases accept H⁺. Every reaction generates two conjugate pairs differing by one proton. Water is amphoteric — it acts as base toward HCl and as acid toward NH₃. Basicity (mono-, di-, tri-) counts only the protons an acid can actually donate, not the total H atoms. The Lewis framework generalises further (electron-pair donor / acceptor), but OCR H432 examines Brønsted–Lowry throughout Module 5.1.3. The next lesson layers quantitative pH onto this foundation via the ionic product of water, Kw.
Reference: OCR A-Level Chemistry A (H432) Module 5.1.3 — Acids, Bases and Buffers, sub-section on the Brønsted–Lowry definition (refer to the official OCR H432 specification document for exact wording).