What is the Brønsted–Lowry theory of acids and bases?

The Brønsted–Lowry theory defines acids as substances that donate protons (H⁺ ions) and bases as substances that accept protons. This theory is fundamental in understanding acid-base equilibria in AS-Level Physical Chemistry, particularly within the Cambridge International A Levels curriculum.

How does the Brønsted–Lowry theory differ from the Arrhenius theory?

While the Arrhenius theory defines acids and bases in terms of their ability to produce H⁺ and OH⁻ ions in water, the Brønsted–Lowry theory broadens this definition by focusing on proton transfer, making it applicable to a wider range of chemical reactions, including those not occurring in aqueous solutions.

Can you give an example of a Brønsted–Lowry acid-base reaction?

A classic example is the reaction between ammonia (NH₃) and water (H₂O). In this reaction, NH₃ acts as a Brønsted–Lowry base by accepting a proton from H₂O, which acts as a Brønsted–Lowry acid. This results in the formation of NH₄⁺ and OH⁻ ions.

What is a conjugate acid-base pair in the Brønsted–Lowry theory?

In the Brønsted–Lowry theory, a conjugate acid-base pair consists of two species that transform into each other by the gain or loss of a proton. For example, in the reaction of acetic acid (CH₃COOH) with water, CH₃COOH and CH₃COO⁻ form a conjugate acid-base pair.

Why is the Brønsted–Lowry theory important for understanding equilibria in chemistry?

The Brønsted–Lowry theory is crucial for understanding chemical equilibria because it explains how acids and bases interact in reversible reactions. This understanding helps predict the direction of equilibrium shifts and the relative strengths of acids and bases, which are key concepts in AS-Level Physical Chemistry under the Cambridge International A Levels curriculum.

Why is "Treating 'strong' and 'concentrated' as the same." a common mistake in Brønsted–Lowry theory of acids and bases?

Why it happens: Loose everyday language. How to avoid it: Strong/weak = degree of dissociation; concentrated/dilute = amount per volume. Source: 9701 Examiner Reports 2022-2024

Why is "Using 'base' and 'alkali' interchangeably." a common mistake in Brønsted–Lowry theory of acids and bases?

Why it happens: They overlap for soluble bases. How to avoid it: An alkali is a soluble base that releases OH⁻ in water; all alkalis are bases, not all bases are alkalis.

Why is "Thinking a weak acid must be dilute / less concentrated." a common mistake in Brønsted–Lowry theory of acids and bases?

Why it happens: Confusing strength with concentration again. How to avoid it: A weak acid can be very concentrated — it is just only partially dissociated.

Why is "Choosing an indicator outside the steep part of the curve." a common mistake in Brønsted–Lowry theory of acids and bases?

Why it happens: Picking a familiar indicator regardless of the equivalence pH. How to avoid it: Match the indicator's range to the equivalence pH (methyl orange acidic, phenolphthalein alkaline).

Why is "Choosing an indicator for a weak acid–weak base titration." a common mistake in Brønsted–Lowry theory of acids and bases?

Why it happens: Assuming every titration has a sharp end point. How to avoid it: There is no steep section, so use a pH meter, not an indicator.

Equilibria(AS-Level Physical Chemistry)Cambridge International A Level Chemistry 9701

Brønsted–Lowry theory of acids and bases

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Brønsted–Lowry Theory of Acids and Bases — Cambridge International AS & A Level Chemistry 9701 (2025-2027 syllabus)

The common acids and alkalis, the Brønsted–Lowry proton transfer theory, strong versus weak acids and bases, neutralisation, pH titration curves and how to choose the right indicator.

What you’ll learn

Mapped to the Cambridge International A Level 9701 syllabus (2025-2027).

7.2.1–7.2.2 — State the names and formulas of common acids (HCl, H₂SO₄, HNO₃, CH₃COOH) and alkalis (NaOH, KOH, NH₃).
7.2.3 — Describe the Brønsted–Lowry theory of acids and bases.
7.2.4 — Describe strong acids/bases as fully dissociated and weak acids/bases as partially dissociated in aqueous solution.
7.2.5 — Appreciate that water has pH 7, acids pH < 7, alkalis pH > 7.
7.2.6 — Explain qualitatively the differences in behaviour between strong and weak acids (reaction with a metal; pH; indicator; conductivity).
7.2.7–7.2.8 — Understand neutralisation (H⁺ + OH⁻ → H₂O) and that salts are formed.
7.2.9–7.2.10 — Sketch pH titration curves and select suitable indicators (pKa not used).

The Brønsted–Lowry theory

Acid donates a proton (H⁺); base accepts one. It's all about proton transfer.

Brønsted–Lowry theory:

An acid is a proton (H⁺) donor.
A base is a proton (H⁺) acceptor.

A neutralisation is a proton transfer from acid to base.

Common acids (learn names + formulas): hydrochloric acid HCl, sulfuric acid H₂SO₄, nitric acid HNO₃, ethanoic acid CH₃COOH.

Common alkalis: sodium hydroxide NaOH, potassium hydroxide KOH, ammonia NH₃.

Worked. In $\text{HCl} + \text{NH}_3 \rightarrow \text{NH}_4^+ + \text{Cl}^-$ : HCl donates H⁺ (acid); NH₃ accepts H⁺ (base).

Acid = proton donor; base = proton acceptor.
Acids: HCl, H₂SO₄, HNO₃, CH₃COOH.
Alkalis: NaOH, KOH, NH₃.

Strong vs weak acids and bases

Strong = fully dissociated into ions; weak = only partially dissociated.

'Strong/weak' refers to the degree of dissociation, not the concentration.

Strong acid — fully dissociated in water: $\text{HCl} \rightarrow \text{H}^+ + \text{Cl}^-$ (essentially 100%).
Weak acid — partially dissociated, an equilibrium: $\text{CH}_3\text{COOH} \rightleftharpoons \text{H}^+ + \text{CH}_3\text{COO}^-$ (only a small % ionises).
Likewise strong base (NaOH, fully dissociated) vs weak base (NH₃, partially: $\text{NH}_3 + \text{H}_2\text{O} \rightleftharpoons \text{NH}_4^+ + \text{OH}^-$ ).

pH scale: pure water is pH 7 (neutral); acids are below 7; alkalis are above 7.

Comparing a strong vs a weak acid at the SAME concentration: the strong acid has more H⁺ ions, so it has:

a lower pH,
a faster reaction with a reactive metal (more vigorous effervescence),
higher electrical conductivity (more ions),
and can be distinguished by a pH meter / universal indicator (different pH).

Strong = fully dissociated; weak = partially (equilibrium).
Water pH 7; acid < 7; alkali > 7.
At same concentration, strong acid: lower pH, faster with metals, higher conductivity.

Neutralisation and salts

Acid + base → salt + water; the core ionic step is H⁺ + OH⁻ → H₂O.

In neutralisation, an acid reacts with a base to form a salt and water. The essential ionic reaction is: $\text{H}^+(aq) + \text{OH}^-(aq) \rightarrow \text{H}_2\text{O}(l)$

Salts are formed when the H⁺ of the acid is replaced by a metal (or ammonium) ion:

HCl + NaOH → NaCl + H₂O
H₂SO₄ + 2KOH → K₂SO₄ + 2H₂O
HNO₃ + NH₃ → NH₄NO₃

The salt's identity comes from the acid's anion (chloride, sulfate, nitrate) and the base's cation.

Acid + base → salt + water.
Ionic core: H⁺ + OH⁻ → H₂O.
Salt = base's cation + acid's anion.

pH titration curves and choosing an indicator

Pick an indicator whose colour-change range falls within the steep vertical part of the curve.

A pH titration curve plots pH against volume of one solution added. The shape depends on whether the acid and base are strong or weak.

A strong acid–strong base curve has a long, steep vertical section through pH 7 — any indicator changing colour in that range works.

The four combinations (where the steep 'vertical' section / equivalence pH lies):

Acid + base	Equivalence pH	Suitable indicator
Strong acid + strong base	~7 (long steep range)	Methyl orange or phenolphthalein
Strong acid + weak base	acidic (~5)	Methyl orange (3.1–4.4)
Weak acid + strong base	alkaline (~9)	Phenolphthalein (8.3–10.0)
Weak acid + weak base	no sharp vertical	No suitable indicator (use a pH meter)

Rule for choosing an indicator: its colour-change range must lie within the steep (vertical) part of the curve. If there is no steep section (weak–weak), no indicator gives a sharp end point.

Curve shape depends on strong/weak combination.
Indicator's range must fall in the vertical section.
Methyl orange for acidic equivalence; phenolphthalein for alkaline.
Weak acid + weak base: no suitable indicator.

How it’s examined

Brønsted–Lowry definitions, strong-vs-weak comparisons, neutralisation equations and indicator selection are reliable Paper 2 marks; titration curves and indicators link to the Paper 3/5 practicals. Examiner reports flag: confusing strong/weak with concentrated/dilute, and choosing an indicator that changes colour outside the steep part of the curve.

Sources: Cambridge International AS & A Level Chemistry 9701 syllabus (2025-2027), section 7.2; 9701 Examiner Reports 2022-2024; 9701/22 May/Jun 2024 question paper and mark scheme. Last reviewed 2026-05-20.

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Step-by-step worked examples — Brønsted–Lowry theory of acids and bases

Step-by-step solutions to past-paper-style questions on brønsted–lowry theory of acids and bases, written exactly the way a tutor would explain them at the board.

1Identifying acid and base (2 marks)
Getting started• Bronsted-Lowry
Question
In $\text{HNO}_3 + \text{H}_2\text{O} \rightarrow \text{H}_3\text{O}^+ + \text{NO}_3^-$ , identify the Brønsted–Lowry acid and base.
Step-by-step solution
1. Step 1
  HNO₃ donates H⁺ → Brønsted–Lowry acid.
2. Step 2
  H₂O accepts H⁺ → Brønsted–Lowry base.
Answer
HNO₃ is the acid (proton donor); H₂O is the base (proton acceptor).
Examiner tip
Define by what they DO with the proton: donor = acid, acceptor = base.
2Identifying conjugate pairs (3 marks)
Getting started• conjugate pairs
Question
In $\text{NH}_3 + \text{H}_2\text{O} \rightleftharpoons \text{NH}_4^+ + \text{OH}^-$ , identify the two conjugate acid–base pairs.
Step-by-step solution
1. Step 1
  NH₃ (base) gains H⁺ → NH₄⁺ (its conjugate acid).
2. Step 2
  H₂O (acid) loses H⁺ → OH⁻ (its conjugate base).
Answer
Pairs: NH₃ / NH₄⁺ and H₂O / OH⁻ (each pair differs by one H⁺).
Examiner tip
Conjugate pairs always differ by exactly one proton — H₂O here is acting as the acid.
3Strong vs weak acid (4 marks)
Building confidence• strong weak
Question
Equal concentrations of hydrochloric acid and ethanoic acid are compared. State and explain two differences.
Step-by-step solution
1. Step 1
  HCl is fully dissociated; CH₃COOH only partially, so HCl has a higher [H⁺].
2. Step 2
  Lower pH for HCl (more H⁺).
3. Step 3
  Faster reaction with a reactive metal / higher conductivity for HCl (more H⁺/ions).
Answer
HCl: lower pH and faster reaction with metals / higher conductivity, because it is fully dissociated (more H⁺).
Examiner tip
Must specify 'same concentration' and link the differences to degree of dissociation / [H⁺].
4Neutralisation product (2 marks)
Building confidence• neutralisation
Question
Write the balanced equation for the reaction of sulfuric acid with potassium hydroxide, naming the salt.
Step-by-step solution
1. Step 1
  Acid + base → salt + water; H₂SO₄ is diprotic.
  $\text{H}_2\text{SO}_4 + 2\text{KOH} \rightarrow \text{K}_2\text{SO}_4 + 2\text{H}_2\text{O}$
Answer
Salt is potassium sulfate, K₂SO₄.
Examiner tip
Diprotic acid needs 2 KOH; the underlying ionic step is H⁺ + OH⁻ → H₂O.
5Choosing an indicator (3 marks)
Stretch• titration, indicator
Question
Ethanoic acid (weak) is titrated with sodium hydroxide (strong). Which indicator is suitable and why?
Step-by-step solution
1. Step 1
  Weak acid + strong base → equivalence point is alkaline (~pH 9).
2. Step 2
  The steep section is in the alkaline region, so the indicator must change colour there.
3. Step 3
  Phenolphthalein (8.3–10.0) changes in that range; methyl orange would change too early.
Answer
Phenolphthalein — its range (8.3–10) lies within the alkaline steep section.
Examiner tip
Match the indicator's range to the vertical (steep) part of the curve, set by the acid/base strengths.
6Distinguishing strong from weak acid (4 marks)
Stretch• strong weak, practical
Question
You have equal-concentration solutions of a strong acid (HA) and a weak acid (HB). Describe two methods to tell which is which, with the expected results. (4)
Step-by-step solution
1. Step 1
  pH meter / measure pH: the strong acid HA has the LOWER pH (higher [H⁺] from full dissociation).
2. Step 2
  Electrical conductivity: HA conducts better (more ions in solution).
3. Step 3
  Alternatively, rate with Mg / carbonate: HA fizzes faster initially (higher [H⁺]).
4. Step 4
  Note: the same concentration is essential — strength is about dissociation, not amount.
Answer
Lower pH, higher conductivity and a faster initial reaction all point to the strong acid (more dissociated).
Examiner tip
Total volume of gas / amount of base needed in a titration is the SAME for both — only the rate, pH and conductivity differ.

Key Definitions and Keywords — Brønsted–Lowry theory of acids and bases

Definitions to memorise and the exact keywords mark schemes credit for brønsted–lowry theory of acids and bases answers — sharpened from recent examiner reports for the 2026 Cambridge International A Level 9701 sitting.

Brønsted–Lowry acid / base
Examiner keyword
An acid is a proton (H⁺) donor; a base is a proton (H⁺) acceptor.
Conjugate acid–base pair
Examiner keyword
Two species that differ by one proton: removing H⁺ from an acid gives its conjugate base; adding H⁺ to a base gives its conjugate acid.
Strong / weak acid
Examiner keyword
A strong acid is fully dissociated in aqueous solution; a weak acid is only partially dissociated (a dissociation equilibrium).
Amphoteric / amphiprotic
Examiner keyword
A species that can act as either an acid or a base (e.g. water can donate or accept a proton).
Neutralisation
Examiner keyword
The reaction of an acid with a base to form a salt and water; the ionic equation is H⁺(aq) + OH⁻(aq) → H₂O(l).

Common Mistakes and Misconceptions — Brønsted–Lowry theory of acids and bases

The traps other students keep falling into on brønsted–lowry theory of acids and bases questions — taken from recent Cambridge International A Level 9701 examiner reports and mark schemes — and how to avoid them.

✕Treating 'strong' and 'concentrated' as the same.
9701 Examiner Reports 2022-2024
Why it happens
Loose everyday language.
How to avoid it
Strong/weak = degree of dissociation; concentrated/dilute = amount per volume.
✕Using 'base' and 'alkali' interchangeably.
Why it happens
They overlap for soluble bases.
How to avoid it
An alkali is a soluble base that releases OH⁻ in water; all alkalis are bases, not all bases are alkalis.
✕Thinking a weak acid must be dilute / less concentrated.
Why it happens
Confusing strength with concentration again.
How to avoid it
A weak acid can be very concentrated — it is just only partially dissociated.
✕Choosing an indicator outside the steep part of the curve.
Why it happens
Picking a familiar indicator regardless of the equivalence pH.
How to avoid it
Match the indicator's range to the equivalence pH (methyl orange acidic, phenolphthalein alkaline).
✕Choosing an indicator for a weak acid–weak base titration.
Why it happens
Assuming every titration has a sharp end point.
How to avoid it
There is no steep section, so use a pH meter, not an indicator.

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Brønsted–Lowry theory of acids and bases

Short Study Notes in the form of Slides

Short Study Notes — Brønsted–Lowry theory of acids and bases

Detailed Notes

What you’ll learn

How it’s examined

1Identifying acid and base (2 marks)

2Identifying conjugate pairs (3 marks)

3Strong vs weak acid (4 marks)

4Neutralisation product (2 marks)

5Choosing an indicator (3 marks)

6Distinguishing strong from weak acid (4 marks)

Brønsted–Lowry acid / base

Conjugate acid–base pair

Strong / weak acid

Amphoteric / amphiprotic

Neutralisation

✕Treating 'strong' and 'concentrated' as the same.

✕Using 'base' and 'alkali' interchangeably.

✕Thinking a weak acid must be dilute / less concentrated.

✕Choosing an indicator outside the steep part of the curve.

✕Choosing an indicator for a weak acid–weak base titration.

Practice questions

Past paper style quiz

Assess your understanding

Video lesson

Brønsted–Lowry theory of acids and bases — frequently asked questions