Acids and Bases: The Brønsted-Lowry Theory

This guide covers subtopic 7.2, Brønsted-Lowry theory of acids and bases, from Topic 7, Equilibria, of Cambridge International AS & A Level Chemistry 9701, 2025–2027 series. This is AS Level content.

A note on this resource

Subtopic 7.2 sits inside Topic 7, Equilibria, alongside 7.1 (Kc, Kp and Le Chatelier’s principle), but the two are conceptually distinct — one is about general equilibrium constants, the other is a specific theory of acid-base behaviour — so they’re covered as two separate resources rather than one combined page. See Chemical Equilibria: Kc, Kp and Le Chatelier’s Principle for 7.1.

Before studying this

At IGCSE or O Level (Topic 7, Acids, bases and salts), you met acids and bases qualitatively — naming common acids and alkalis, using indicators, and neutralisation reactions — without a formal theory of what an acid or base is. Marlbridge does not yet have a dedicated resource for that IGCSE/O Level topic; if it’s unfamiliar, the outcomes below are written to be self-contained.

AS Level introduces the Brønsted-Lowry theory as a precise definition of acids and bases, the strong/weak distinction (as opposed to concentrated/dilute, which is a separate idea entirely), and reading titration curves. Quantitative pH calculation, Ka, pKa and buffer solutions are A Level (A2) content, not AS — this resource stays qualitative throughout, exactly as the AS syllabus does.

Syllabus coverage

CAMBRIDGE INTERNATIONAL AS & A LEVEL CHEMISTRY 9701 — AS Level, Topic 7

7.2 Brønsted-Lowry theory of acids and bases — stating the names and formulas of common acids (hydrochloric, HCl; sulfuric, H₂SO₄; nitric, HNO₃; ethanoic, CH₃COOH) and common alkalis (sodium hydroxide, NaOH; potassium hydroxide, KOH; ammonia, NH₃); describing the Brønsted-Lowry theory of acids and bases; describing strong acids and bases as fully dissociated and weak acids and bases as partially dissociated in aqueous solution; appreciating that water has pH 7, acids below 7, and alkalis above 7; explaining qualitatively the differences in behaviour between strong and weak acids, including reaction with a reactive metal and differences in pH by pH meter, universal indicator or conductivity; understanding that neutralisation occurs when H⁺(aq) and OH⁻(aq) form H₂O(l); understanding that salts are formed in neutralisation reactions; sketching pH titration curves for combinations of strong and weak acids with strong and weak alkalis; selecting suitable indicators for acid-alkali titrations, given appropriate data (pKa values not used).

The Brønsted-Lowry theory

A Brønsted-Lowry acid is a proton (H⁺) donor. A Brønsted-Lowry base is a proton (H⁺) acceptor. This is a behavioural definition — it describes what a substance does in a reaction, not just what class of compound it belongs to, which is why ammonia (a gas, with no hydroxide ion in its own formula) still counts as a common alkali: dissolved in water, it accepts a proton from water to form NH₄⁺ and OH⁻.

Strong vs weak — not the same as concentrated vs dilute

This is the single most important distinction in this topic, and the one most often confused with something else:

• Strong and weak describe the extent of dissociation into ions in aqueous solution. A strong acid (like HCl) is essentially 100% dissociated; a weak acid (like CH₃COOH) exists mostly as undissociated molecules, with only a small proportion dissociated into ions at any moment.
• Concentrated and dilute describe how much acid is dissolved in a given volume of water — entirely independent of strength. You can have a concentrated weak acid or a dilute strong acid.

Observable differences between a strong and a weak acid of the same concentration:

• Reaction with a reactive metal: the strong acid reacts faster, because it has a higher concentration of free H⁺ ions immediately available.
• pH: the strong acid has a lower pH (more free H⁺), measurable with a pH meter or universal indicator.
• Conductivity: the strong acid conducts better, because full dissociation produces more ions in solution.

The pH scale and neutralisation

At this level, pH is understood descriptively rather than calculated: pure water has pH 7, acidic solutions have pH below 7, and alkaline solutions have pH above 7.

Neutralisation is the reaction of H⁺(aq) with OH⁻(aq) to form water:

H⁺(aq) + OH⁻(aq) → H₂O(l)

The positive ion left from the acid and the negative ion left from the base combine to form a salt — for example, hydrochloric acid and sodium hydroxide neutralise to form sodium chloride and water.

Titration curves

A pH titration curve plots pH against volume of acid or alkali added. Its shape depends on whether each reactant is strong or weak:

• Strong acid–strong base: a long, steep vertical section through the equivalence point, centred close to pH 7.
• Strong acid–weak base: a steep section, but the equivalence point sits below pH 7 (the salt formed is itself weakly acidic in solution).
• Weak acid–strong base: a steep section with the equivalence point above pH 7 (the salt formed is weakly alkaline in solution).
• Weak acid–weak base: no sharply defined vertical section — the pH change through the equivalence point is small and gradual, making the end point hard to detect with an indicator.

Choosing an indicator

An indicator is suitable for a titration if its colour-change range falls within the steep, near-vertical part of the curve — that’s the region where a single drop of titrant causes a large pH change, so the colour change happens sharply at the true equivalence point. Two indicators cover most AS-level cases: methyl orange (changes colour in acidic conditions) and phenolphthalein (changes colour in alkaline conditions). For a strong acid–strong base titration, either works, since the vertical section spans both ranges; for titrations with a weak reactant, the indicator must be chosen to match which side of pH 7 the equivalence point falls on. A weak acid–weak base titration has no sharp vertical section at all, so no simple indicator gives a reliable end point.

Common mistakes

• Equating “weak acid” with “dilute acid.” They describe different things — strength is about dissociation, concentration is about how much acid is dissolved. A weak acid can be concentrated; a strong acid can be dilute.
• Assuming a weak acid has a higher pH because it’s “less acidic” overall. A weak acid genuinely does have a higher pH than a strong acid of the same concentration — but that’s a consequence of partial dissociation producing fewer free H⁺ ions, not a statement about how much total acid is present.
• Picking phenolphthalein or methyl orange without checking which side of pH 7 the equivalence point falls on. The right indicator depends on whether the titration is acid–base, and which reactant (if either) is weak.
• Trying to calculate pH numerically at AS Level. Ka, pKa and pH calculations are A Level (A2) content — at AS, pH is described qualitatively only.

Quick revision checklist

• Brønsted-Lowry: acid = proton donor, base = proton acceptor
• Strong/weak = extent of dissociation; concentrated/dilute = amount dissolved — never conflate the two
• Observable differences between strong and weak acids: rate with a metal, pH, conductivity
• pH scale: water = 7, acids < 7, alkalis > 7 (descriptive, not calculated)
• Neutralisation: H⁺(aq) + OH⁻(aq) → H₂O(l), forming a salt
• Titration curve shapes for strong/strong, strong/weak, weak/strong and weak/weak combinations
• Choosing an indicator: colour-change range must sit on the vertical part of the curve

Related resources

• Chemical Equilibria: Kc, Kp and Le Chatelier’s Principle — the other AS Level equilibria topic
• Redox Processes: Oxidation Numbers and Electron Transfer — another electron/proton-transfer classification system
• Cambridge AS & A Level Chemistry hub

Written against Cambridge International AS & A Level Chemistry 9701, 2025–2027 series. Always check the current syllabus for your examination year.