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Spec Mapping — OCR H432 Module 2.1.4 — Acids, covering the Brønsted–Lowry definition of acids and bases, common laboratory acids and alkalis, the distinction between strong and weak acids, neutralisation reactions (acid + metal, acid + metal oxide, acid + metal hydroxide, acid + metal carbonate), salt formation, and the conceptual difference between acid strength and concentration (refer to the official OCR H432 specification document for exact wording). This lesson is the conceptual bridge between GCSE acid-base ideas and the quantitative pH/Ka/Kw work in Year 13 Module 5.1.3.
The acid–base chemistry you study at A-Level is not just a tidier version of the GCSE story — it is a complete reframing built on the Brønsted–Lowry definition of acids as proton donors and bases as proton acceptors. That single shift in vocabulary opens up conjugate acid–base pair logic, lets you describe weak acid equilibria sensibly, and underpins every titration calculation, indicator choice, and buffer derivation in the rest of the course. This lesson establishes the qualitative foundation — what counts as an acid or base, the difference between strong and concentrated, salt nomenclature, and the four classes of neutralisation reaction OCR can examine — so that the next lesson can layer on stoichiometric titration arithmetic and Module 5.1.3 can layer on quantitative pH derivation.
Key Definitions:
- Acid — a species that donates a proton (H⁺) to another species (Brønsted–Lowry).
- Base — a species that accepts a proton (Brønsted–Lowry).
- Alkali — a soluble base that releases OH⁻ ions in aqueous solution.
- Conjugate acid–base pair — two species that differ by a single proton.
The Danish chemist Brønsted and the English chemist Lowry proposed the proton-transfer definition of acids and bases independently in 1923. Their formulation displaced the earlier Arrhenius picture (acids release H⁺, bases release OH⁻) because it survives in non-aqueous solvents — the same logic explains the behaviour of NH₃ + HCl in the gas phase, or HCl in glacial ethanoic acid, where the Arrhenius framework breaks down. The Brønsted–Lowry definition is the one OCR examines.
When an acid dissolves in water it does not produce a "free" proton — bare protons cannot exist in aqueous solution because they are vanishingly small and the electric field at their surface is enormous. Instead the proton is immediately accepted by a water molecule:
HCl(aq)+H2O(l)→H3O+(aq)+Cl−(aq)
The product H₃O⁺ is the oxonium ion (sometimes called the hydronium ion). For shorthand we write H⁺(aq), but every reference to H⁺(aq) in this course should be read as H₃O⁺(aq).
A wider definition exists — the Lewis definition (G. N. Lewis, 1923) frames an acid as an electron-pair acceptor and a base as an electron-pair donor, which extends acid-base behaviour to species with no proton at all (e.g. BF₃ accepting the lone pair of NH₃). Lewis acid-base ideas appear at the fringes of the OCR H432 specification when you meet dative covalent bonds and transition-metal complex formation later in this course and in Module 6. For Module 2.1.4, however, the Brønsted–Lowry definition is the operational one.
You must know the formulae of common laboratory acids, their basicity (number of acidic protons per molecule), and whether they are strong or weak.
| Acid | Formula | Basicity | Type | Notes |
|---|---|---|---|---|
| Hydrochloric acid | HCl | 1 | Strong | Monoprotic; near-100 % dissociation |
| Nitric acid | HNO₃ | 1 | Strong | Strong oxidising agent in concentrated form |
| Sulfuric acid | H₂SO₄ | 2 | Strong (1st) / weak (2nd) | Diprotic; first H is strong, HSO₄⁻ is a weak acid |
| Phosphoric(V) acid | H₃PO₄ | 3 | Weak | Triprotic; first Ka≈7×10−3 |
| Ethanoic acid | CH₃COOH | 1 | Weak | Carboxylic acid; Ka≈1.7×10−5 |
| Methanoic acid | HCOOH | 1 | Weak | Strongest of the simple carboxylic acids |
| Carbonic acid | H₂CO₃ | 2 | Weak | Aqueous CO₂; first Ka≈4×10−7 |
Basicity is the number of replaceable hydrogens per acid molecule, not a measure of strength. H₂SO₄ is diprotic because each molecule can in principle donate two protons:
H2SO4→H++HSO4−(strong) HSO4−⇌H++SO42−(weak; partial)
Notice the asymmetry — the first dissociation is strong (single arrow), the second is weak (equilibrium arrow). At A-Level you usually treat H₂SO₄ as fully diprotic for stoichiometric calculations, but be alert that the rigorous pH of dilute H₂SO₄ is not simply −log(2c).
| Base / alkali | Formula | Solubility | Strength |
|---|---|---|---|
| Sodium hydroxide | NaOH | High | Strong alkali |
| Potassium hydroxide | KOH | High | Strong alkali |
| Calcium hydroxide | Ca(OH)₂ | Sparingly soluble | Strong (within solubility) |
| Ammonia (aq) | NH₃ | High | Weak alkali |
| Copper(II) oxide | CuO | Insoluble | Insoluble base |
| Magnesium oxide | MgO | Insoluble | Insoluble base (basic oxide) |
| Calcium carbonate | CaCO₃ | Insoluble | Insoluble base |
| Sodium carbonate | Na₂CO₃ | Soluble | Weak alkali (hydrolysis) |
The distinction between a base and an alkali is one of solubility — every alkali is a base, but not every base is an alkali. CuO and MgO are bases but not alkalis because they are insoluble in water; they react with acid in their solid form. Aqueous ammonia is an alkali, but it is a weak one because most NH₃ molecules in solution remain unprotonated:
NH3(aq)+H2O(l)⇌NH4+(aq)+OH−(aq)
This is the single most-confused pair of distinctions in A-Level acid-base chemistry. Hold them apart deliberately:
| Term | What it describes | Example |
|---|---|---|
| Strong | Acid (or base) is essentially fully dissociated in water | 0.1 mol dm⁻³ HCl is fully dissociated |
| Weak | Acid (or base) is only partially dissociated; an equilibrium exists | 0.1 mol dm⁻³ CH₃COOH is ~1 % dissociated |
| Concentrated | Many moles of acid per dm³ of solution | 17 mol dm⁻³ glacial ethanoic acid |
| Dilute | Few moles of acid per dm³ of solution | 0.001 mol dm⁻³ HCl |
You can have:
Conductivity is an excellent diagnostic: glacial ethanoic acid conducts electricity poorly compared with dilute HCl, despite holding thousands of times more acid molecules per dm³ — because only the dissociated fraction carries charge. Strength (degree of dissociation) is a property of the acid; concentration is a property of the solution.
flowchart TD
A[Acid HA in water] --> B{Fully dissociated?}
B -- Yes --> C[Strong acid HCl HNO3 H2SO4]
B -- No --> D[Weak acid CH3COOH H3PO4 HF]
A --> E{Mol per dm3 of solution?}
E -- High --> F[Concentrated solution]
E -- Low --> G[Dilute solution]
C --> H[Strong dilute or strong concentrated both possible]
D --> I[Weak dilute or weak concentrated both possible]
The decision tree above makes the orthogonality explicit: strength and concentration are independent axes.
Every Brønsted–Lowry acid-base reaction involves the transfer of a single proton, which automatically generates two conjugate acid-base pairs. Consider ethanoic acid in water:
CH3COOH(aq)+H2O(l)⇌CH3COO−(aq)+H3O+(aq)
Pair 1: CH₃COOH / CH₃COO⁻. Pair 2: H₂O / H₃O⁺.
Reading the equation right-to-left, H₃O⁺ donates a proton back to CH₃COO⁻ — so H₃O⁺ is acting as the acid and CH₃COO⁻ as the base in the reverse reaction. This is why we use the equilibrium arrow: both directions are occurring, and the position of equilibrium depends on the relative strengths of the two acids on either side.
Rule of thumb: an equilibrium always lies on the side of the weaker acid (and weaker base). For CH₃COOH + H₂O, the equilibrium lies far to the left because CH₃COOH is a weaker acid than H₃O⁺.
Water can act as either acid or base — it is amphiprotic (specifically) and amphoteric (the more general term):
Aluminium oxide and zinc oxide are amphoteric oxides — they react with both acids and alkalis:
Al2O3+6HCl→2AlCl3+3H2O Al2O3+2NaOH+3H2O→2NaAl(OH)4
Amphoteric oxide behaviour is a useful diagnostic in qualitative inorganic chemistry (a topic you will return to in the periodicity-groups course).
Neutralisation is the reaction of an acid with a base to form a salt and water (with possible additional products). OCR examines four canonical reaction types — you must be able to write balanced equations with state symbols for each.
Mg(s)+2HCl(aq)→MgCl2(aq)+H2(g) Zn(s)+H2SO4(aq)→ZnSO4(aq)+H2(g)
Only metals above hydrogen in the reactivity series react significantly. The chemistry is a redox reaction (the metal is oxidised, H⁺ is reduced) — a thematic link to the redox lesson later in this course. Lead reacts very slowly with dilute HCl and H₂SO₄ because PbCl₂ and PbSO₄ are insoluble and coat the metal surface, preventing further attack.
CuO(s)+2HCl(aq)→CuCl2(aq)+H2O(l) MgO(s)+H2SO4(aq)→MgSO4(aq)+H2O(l)
CuO is a black powder; on warming with HCl it dissolves to give a blue-green CuCl₂ solution — a classic A-Level practical illustration.
NaOH(aq)+HCl(aq)→NaCl(aq)+H2O(l) 2KOH(aq)+H2SO4(aq)→K2SO4(aq)+2H2O(l)
The ionic equation stripping out spectator Na⁺ and Cl⁻ is the universal one for any strong acid + strong alkali pair:
H+(aq)+OH−(aq)→H2O(l),ΔH≈−57 kJ mol−1
The enthalpy of neutralisation for any strong acid + strong alkali is the same value because the underlying reaction is identical at the ionic level — a beautiful illustration of why ionic equations matter.
CaCO3(s)+2HCl(aq)→CaCl2(aq)+H2O(l)+CO2(g) Na2CO3(aq)+H2SO4(aq)→Na2SO4(aq)+H2O(l)+CO2(g)
Effervescence and a positive limewater test (CO₂ turns limewater milky) confirm the carbonate identification — a key diagnostic in qualitative analysis.
A useful unifying mnemonic — the SVG below summarises which class of neutralisation produces which by-product alongside the salt and water:
When an acid neutralises a base, the salt name is constructed as metal name + anion name, where the anion comes from the acid:
| Acid | Anion | Salt name suffix | Example salt |
|---|---|---|---|
| HCl | Cl⁻ | chloride | NaCl — sodium chloride |
| HNO₃ | NO₃⁻ | nitrate | KNO₃ — potassium nitrate |
| H₂SO₄ | SO₄²⁻ | sulfate | CuSO₄ — copper(II) sulfate |
| H₃PO₄ | PO₄³⁻ | phosphate | Na₃PO₄ — trisodium phosphate |
| CH₃COOH | CH₃COO⁻ | ethanoate | Na(CH₃COO) — sodium ethanoate |
| H₂CO₃ | CO₃²⁻ / HCO₃⁻ | carbonate / hydrogencarbonate | Na₂CO₃, NaHCO₃ |
The cation in solution and the spectator anion in the ionic equation are formal — they are present but do not participate in the bond-breaking / bond-making chemistry. Stripping out the spectator ions reveals the true reaction:
Q: Write a balanced equation, with state symbols, for the reaction of magnesium with dilute sulfuric acid.
Mg(s)+H2SO4(aq)→MgSO4(aq)+H2(g)
Comment: one mole of Mg provides two electrons (Mg → Mg²⁺ + 2e⁻) and one mole of H₂SO₄ provides two H⁺, so the stoichiometry is 1 : 1, not 1 : 2 as it would be for monoprotic HCl.
Q: Write the symbol and ionic equations for the reaction of sodium carbonate with dilute nitric acid.
Symbol equation: Na2CO3(aq)+2HNO3(aq)→2NaNO3(aq)+H2O(l)+CO2(g)
Ionic equation (Na⁺ and NO₃⁻ are spectators): CO32−(aq)+2H+(aq)→H2O(l)+CO2(g)
Q: In the reaction NH3+H2O⇌NH4++OH−, identify the two conjugate acid–base pairs.
NH₃ accepts a proton to become NH₄⁺ — so NH₃ is the base and NH₄⁺ its conjugate acid. H₂O donates a proton to become OH⁻ — so H₂O is the acid and OH⁻ its conjugate base.
Pair 1: NH₃ / NH₄⁺ (base / conjugate acid). Pair 2: H₂O / OH⁻ (acid / conjugate base).
Q: Estimate the pH of 0.05 mol dm⁻³ HCl(aq).
HCl is a strong monoprotic acid, so [H+]=0.05 mol dm⁻³.
pH=−log10(0.05)≈1.30
Quantitative weak-acid pH calculations follow the same procedure but require Ka, which is the content of Module 5.1.3.
Q: A student labels two beakers: (A) 0.01 mol dm⁻³ HCl(aq) and (B) 4 mol dm⁻³ CH₃COOH(aq). Which beaker has the higher H⁺?
A: [H+]=0.01 mol dm⁻³ (fully dissociated).
B: ethanoic acid is weak; using Ka≈1.7×10−5, [H+]≈Ka⋅c=1.7×10−5×4≈8.2×10−3 mol dm⁻³.
So A has the higher [H⁺] despite being far more dilute. This is the strong-vs-concentrated lesson made quantitative — and a classic OCR exam trap.
Synoptic Links — Connects to:
ocr-alevel-chemistry-acids-redox-bonding / acid-base-titrations(the volumetric framework that uses the stoichiometry established here).ocr-alevel-chemistry-acids-bases-buffers(the Year 13 quantitative treatment — Ka, Kw, Henderson–Hasselbalch, buffer derivations all sit on this Brønsted–Lowry foundation).ocr-alevel-chemistry-periodicity-groups(acid-base behaviour of Period 3 oxides; amphoteric Al₂O₃; group 2 hydroxide solubility trend).ocr-alevel-chemistry-enthalpy-rates-equilibrium(the universal ΔH≈−57 kJ mol−1 for strong acid + strong alkali is examined in calorimetry questions).
Practical Activity Group anchor: PAG 2 — Acid–base titration. While the quantitative titration practical is anchored in the next lesson, the conceptual framework for acid identity (mono- vs diprotic), choice of indicator (function of strong/weak character), and salt naming all originate here. The exam-style write-up of any PAG 2 titration must use the Brønsted–Lowry vocabulary established in this lesson.
Question (6 marks): A student is given two unlabelled colourless solutions: 0.1 mol dm⁻³ hydrochloric acid and 0.1 mol dm⁻³ ethanoic acid. Both have the same concentration, but they behave differently in the laboratory. Explain, using the Brønsted–Lowry framework, why the two acids have different pH values and different electrical conductivities, and outline an experiment the student could use to distinguish them.
| Mark | AO | Awarded for |
|---|---|---|
| 1 | AO1 | Definition of acid as proton donor (Brønsted–Lowry) |
| 2 | AO1 | Statement that HCl is a strong acid (fully dissociated) and CH₃COOH is a weak acid (partially dissociated) |
| 3 | AO2 | Linking degree of dissociation to [H⁺] and therefore to pH |
| 4 | AO2 | Linking [ions] to electrical conductivity |
| 5 | AO2 | Sensible experiment — pH meter, conductivity meter, or rate of reaction with Mg ribbon |
| 6 | AO3 | Predicted comparison — HCl has lower pH and higher conductivity than CH₃COOH at the same concentration |
AO split: AO1 = 2, AO2 = 3, AO3 = 1.
Both HCl and ethanoic acid are Brønsted–Lowry acids because they donate protons to water. However, HCl is a strong acid, which means it fully dissociates in water. Ethanoic acid is a weak acid, which means only a small fraction of it dissociates. So 0.1 mol dm⁻³ HCl gives nearly 0.1 mol dm⁻³ H⁺, but 0.1 mol dm⁻³ ethanoic acid gives only a small concentration of H⁺. Because pH depends on [H⁺], HCl has a lower pH than ethanoic acid. The student could measure the pH of both solutions with a pH meter — HCl would give about pH 1 and ethanoic acid about pH 3. Alternatively a conductivity meter would show that HCl conducts more electricity because it has more ions in solution. So the strong acid has both lower pH and higher conductivity than the weak acid at the same concentration.
Examiner commentary: M1 (AO1) for the Brønsted–Lowry definition; M1 (AO1) for naming strong vs weak; M1 (AO2) for the dissociation → [H⁺] → pH chain; M1 (AO2) for the ion-count → conductivity link; M1 (AO2) for proposing a pH meter; M1 (AO3) for the predicted comparison. Around 6/6 on a generous mark scheme but the answer feels mechanical. The candidate has not invoked Ka, has not written the dissociation equations explicitly, and has not justified why dissociation is partial in the weak acid case.
Under the Brønsted–Lowry definition both acids donate a proton to water — HCl → H⁺ + Cl⁻, CH₃COOH ⇌ CH₃COO⁻ + H⁺. The arrow choice carries all the chemistry: the single arrow for HCl signals essentially complete dissociation (a "strong" acid), the equilibrium arrow for ethanoic acid signals partial dissociation governed by Ka≈1.7×10−5 mol dm⁻³. For a 0.1 mol dm⁻³ ethanoic acid solution, [H+]≈Ka⋅c≈1.3×10−3 mol dm⁻³, giving pH ≈ 2.9. For 0.1 mol dm⁻³ HCl, [H+]≈0.1 mol dm⁻³, giving pH = 1. The pH difference of approximately two units corresponds to roughly a hundred-fold difference in [H⁺] — a striking demonstration that strength matters independently of concentration.
Conductivity is proportional to total ion concentration (modulated by ion mobility). In 0.1 mol dm⁻³ HCl the dissociated ion count is roughly 0.2 mol dm⁻³ (H⁺ and Cl⁻ together); in 0.1 mol dm⁻³ ethanoic acid it is roughly 2.6 × 10⁻³ mol dm⁻³. A conductivity meter will therefore read about two orders of magnitude higher for HCl.
To distinguish them I would measure the pH of each using a calibrated pH meter — HCl pH ≈ 1, ethanoic acid pH ≈ 2.9. As a confirmatory experiment I would observe the rate of effervescence when equal-sized magnesium ribbons are added: HCl produces vigorous H₂ effervescence almost immediately, while ethanoic acid produces only slow gas evolution because the much lower [H⁺] limits the rate of the redox reaction Mg+2H+→Mg2++H2.
Examiner commentary: Full 6/6 with comfort. M1 (Brønsted–Lowry), M1 (strong/weak identification with equations), M1 (AO2 dissociation → pH using Ka), M1 (AO2 ion count → conductivity), M1 (AO2 named experiment), M1 (AO3 quantitative comparison). The discriminators that lift this to A*: explicit use of Ka, the Kac approximation, the two-orders-of-magnitude framing, and the Mg ribbon backup experiment. Synoptic awareness of the redox mechanism is signal of top-band thinking.
Pedagogical observations — not fabricated statistics:
Acid-base chemistry beyond OCR H432 opens onto two large university programmes: inorganic chemistry (where Lewis acid-base concepts dominate the description of transition-metal complex formation and Frustrated Lewis Pair catalysis — work pioneered by Stephan and Erker since 2006) and physical chemistry (where the proton activity, not just [H⁺], drives reaction thermodynamics — the Debye-Hückel theory of activity coefficients was developed in 1923, the same year as the Brønsted–Lowry framework). Recommended reading: Atkins, Physical Chemistry, chapters on chemical equilibria and electrochemistry; Housecroft & Sharpe, Inorganic Chemistry, on acid-base theory and superacid behaviour (the Magic Acid HF⋅SbF5 system, H0≈−25, can protonate methane). Oxbridge interview-style prompt: "Carborane acids (H(CHB11Cl11)) are even 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 definition — acids donate protons, bases accept them — frames all of A-Level acid-base chemistry. Strong acids dissociate fully; weak acids partially; concentration and strength are independent. Four neutralisation reaction classes (acid + metal, acid + metal oxide, acid + metal hydroxide, acid + metal carbonate) cover OCR's H432 Module 2.1.4 examination scope. Conjugate acid-base pairs differ by a single proton; water is amphoteric, acting as base toward HCl and as acid toward NH₃. The universal ionic equation H⁺(aq) + OH⁻(aq) → H₂O(l) explains why neutralisation enthalpy is the same for any strong acid + strong alkali pair. Master these qualitative ideas now — the next lesson layers titration arithmetic onto exactly this foundation.
Reference: OCR A-Level Chemistry A (H432) Module 2.1.4 (a)–(g) (refer to the official OCR H432 specification document for exact wording).