AQA A-Level Chemistry: Acids and bases (7405) – Revision Notes

These notes condense section 3.1.12 Acids and bases of the AQA AS and A-level Chemistry specification (7404/7405), version 1.2, July 2026, for exams from June 2016 onwards: 3.1.12.1 to 3.1.12.6 and Required practical 9. Everything here is A-level only. The specification places it on A-level Paper 1, and Paper 3 can test any content.

For full explanations and worked examples, read the acids and bases study guide. Then test yourself with the acids and bases practice questions. The course hub is AQA A-Level Chemistry, the printable checklist lists every specification point, and the free 10-minute diagnostics help you find weak spots.

Definitions to learn (3.1.12.1 to 3.1.12.4)

• Acid: a proton donor. Base: a proton acceptor.
• Acid–base equilibrium: an equilibrium involving the transfer of protons.
• Strong acid / strong base: fully dissociated in aqueous solution.
• Weak acid / weak base: dissociates only slightly in aqueous solution.
• pH: −log₁₀[H⁺].
• Kw: the ionic product of water, [H⁺][OH⁻]. It is derived from the equilibrium constant for the dissociation of water.
• Ka: the dissociation constant for a weak acid. pKa = −log₁₀ Ka.
• Buffer: a solution that keeps an approximately constant pH despite dilution or the addition of small amounts of acid or base.

Key equations

Quantity	Equation	Units
pH	pH = −log₁₀[H⁺]	none
[H⁺] from pH	[H⁺] = 10^(−pH)	mol dm⁻³
Ionic product of water	Kw = [H⁺][OH⁻]	mol² dm⁻⁶
Strong base	[H⁺] = Kw / [OH⁻]	mol dm⁻³
Pure water	[H⁺] = [OH⁻] = √Kw	mol dm⁻³
Weak acid HA	Ka = [H⁺][A⁻] / [HA]	mol dm⁻³
Weak acid alone	[H⁺] = √(Ka × [HA])	mol dm⁻³
pKa	pKa = −log₁₀ Ka; Ka = 10^(−pKa)	none
Acidic buffer	[H⁺] = Ka × [HA] / [A⁻]	mol dm⁻³
Half-neutralisation	pH = pKa	none

Use Kw = 1.00 × 10⁻¹⁴ mol² dm⁻⁶ at 298 K unless a question gives another value.

Proton transfer in one line

In NH₃ + H₂O ⇌ NH₄⁺ + OH⁻, water is the acid and NH₃ the base going forwards; NH₄⁺ is the acid and OH⁻ the base going backwards. Name the proton donor and acceptor in both directions.

Kw and temperature

• [H₂O] is effectively constant, so Kw = Kc × [H₂O].
• Kw increases with temperature, so the dissociation of water is endothermic.
• At higher temperature, pure water has pH below 7 but is still neutral: [H⁺] = [OH⁻].

Method in steps: pH of a strong acid or strong base

1. Strong monoprotic acid: [H⁺] = concentration of acid.
2. Strong base (NaOH, KOH): [OH⁻] = concentration of base.
3. For a base, [H⁺] = Kw / [OH⁻].
4. pH = −log₁₀[H⁺], to 2 decimal places.

Method in steps: pH of a weak acid

1. Write Ka = [H⁺][A⁻] / [HA].
2. Assume [H⁺] = [A⁻], so Ka = [H⁺]² / [HA].
3. Assume [HA] at equilibrium = starting concentration (only slight dissociation).
4. [H⁺] = √(Ka × [HA]); then pH.

Going the other way (Ka from pH): [H⁺] = 10^(−pH), then Ka = [H⁺]² / [HA]. Give Ka with units, mol dm⁻³.

Method in steps: pH after mixing a strong acid and a strong base

1. Moles of H⁺ and moles of OH⁻.
2. Subtract to find which is in excess and by how much.
3. Divide by the total volume in dm³.
4. Excess H⁺: pH directly. Excess OH⁻: use Kw first.

Method in steps: titration results

1. Ignore the rough titre. Average the concordant titres.
2. Moles of the known solution = concentration × titre (dm³).
3. Use the equation ratio (1 : 1 for monoprotic acid with NaOH).
4. Concentration = moles / volume of the unknown (dm³).

pH curves (3.1.12.5)

Titration	Start pH	Vertical section	Equivalence pH	Indicator
Strong acid + strong base	Low	Long	About 7	Methyl orange or phenolphthalein
Weak acid + strong base	Higher	Alkaline range	Above 7	Phenolphthalein
Strong acid + weak base	Low	Acidic range	Below 7	Methyl orange
Weak acid + weak base	Higher	None sharp	Unclear	None suitable; use a pH meter

Colour change ranges: methyl orange 3.1–4.4; phenolphthalein 8.3–10.0. An indicator works when its whole range sits inside the vertical section.

Sketching tips: mark the equivalence volume on the x-axis; show the flat buffer region for a weak acid with strong base; the curve levels off as excess base or acid is added.

Required practical 9

• Calibrate a pH meter with buffers of known pH.
• Pipette the acid into a beaker; add the base from a burette in small portions, swirling and recording pH each time.
• Use smaller additions near the end point.
• Plot pH against volume. Equivalence is the midpoint of the steep section.
• Weak acid run: pH at half the equivalence volume = pKa, so Ka = 10^(−pH).
• Repeat with a strong acid and a weak base.

Buffers (3.1.12.6)

Type	Made from	Example	Added H⁺	Added OH⁻
Acidic	Weak acid + its salt	CH₃COOH / CH₃COONa	H⁺ + A⁻ → HA	OH⁻ + HA → A⁻ + H₂O
Basic	Weak base + its salt	NH₃ / NH₄Cl	NH₃ + H⁺ → NH₄⁺	NH₄⁺ + OH⁻ → NH₃ + H₂O

Why the pH barely changes: the buffer holds large amounts of the acid and its salt (or base and its salt). Small additions change the ratio [HA]/[A⁻] only slightly. Dilution leaves the ratio unchanged.

An acidic buffer can be made by partially neutralising a weak acid with a strong base: HA left = starting HA − OH⁻ added; A⁻ formed = OH⁻ added.

Applications: blood (carbonic acid–hydrogencarbonate keeps pH 7.35–7.45), fermentation, fabric dyeing, calibrating pH meters.

Method in steps: pH of an acidic buffer

1. Find moles of HA and A⁻ (from the mixture, or after a reaction with added acid or base).
2. Added H⁺: HA goes up, A⁻ goes down by the same amount. Added OH⁻: the reverse.
3. [H⁺] = Ka × moles HA / moles A⁻ (volume cancels).
4. pH = −log₁₀[H⁺].

Must-know distinctions

• Strong vs concentrated: strong means fully dissociated; concentrated means a lot of acid per dm³.
• Weak acid alone vs buffer: alone, [H⁺] = [A⁻]; in a buffer, [A⁻] comes mainly from the salt.
• Neutral vs pH 7: neutral means [H⁺] = [OH⁻]; pH 7 is neutral only at 298 K.
• Equivalence point vs end point: equivalence is where the moles match the equation; the end point is where the indicator changes colour.
• Ka vs pKa: larger Ka means stronger acid; larger pKa means weaker acid.

Quick self-test

1. Define a Brønsted–Lowry base.
2. Calculate the pH of 0.0040 mol dm⁻³ HNO₃.
3. Calculate [H⁺] in a solution of pH 4.85.
4. Calculate the pH of 0.0125 mol dm⁻³ NaOH at 298 K.
5. Ka = 3.2 × 10⁻⁶ mol dm⁻³. Calculate pKa.
6. pKa = 3.40. Calculate Ka.
7. Calculate the pH of 0.120 mol dm⁻³ HA, Ka = 2.0 × 10⁻⁵ mol dm⁻³.
8. State whether the equivalence pH of a strong acid–weak base titration is above or below 7, and name a suitable indicator.
9. A buffer holds 0.300 mol HA and 0.100 mol NaA (Ka = 1.0 × 10⁻⁵ mol dm⁻³). Calculate its pH.
10. Explain why diluting a buffer hardly changes its pH.
11. At temperature T, Kw = 1.2 × 10⁻¹³ mol² dm⁻⁶. Calculate the pH of pure water at T.
12. In RP9, the pH at half the equivalence volume is 3.70. Calculate Ka.

Answers

1. A proton acceptor.
2. pH = −log₁₀(0.0040) = 2.40
3. 10^(−4.85) = 1.41 × 10⁻⁵ mol dm⁻³
4. [H⁺] = 1.00 × 10⁻¹⁴ / 0.0125 = 8.0 × 10⁻¹³; pH = 12.10
5. pKa = 5.49
6. Ka = 10^(−3.40) = 3.98 × 10⁻⁴ mol dm⁻³
7. [H⁺] = √(2.0 × 10⁻⁵ × 0.120) = 1.55 × 10⁻³; pH = 2.81
8. Below 7; methyl orange.
9. [H⁺] = 1.0 × 10⁻⁵ × 0.300/0.100 = 3.0 × 10⁻⁵; pH = 4.52
10. [HA] and [A⁻] fall by the same factor, so their ratio, and [H⁺] = Ka[HA]/[A⁻], stay the same.
11. [H⁺] = √(1.2 × 10⁻¹³) = 3.46 × 10⁻⁷; pH = 6.46 (still neutral).
12. pH = pKa, so Ka = 10^(−3.70) = 2.00 × 10⁻⁴ mol dm⁻³

Where marks are usually lost

• Using √(Ka[HA]) for a buffer instead of Ka × [HA]/[A⁻].
• Forgetting to use the total volume after mixing two solutions.
• Leaving out units for Ka (mol dm⁻³) or Kw (mol² dm⁻⁶).
• Calling hot pure water acidic because its pH is below 7.
• Writing the Ka expression with [H₂O] in it, or with [H⁺]² before stating the assumption.
• Choosing an indicator whose range only partly overlaps the vertical section.
• Sketching a weak acid–strong base curve with an equivalence pH of 7.
• Not adjusting moles of HA and A⁻ after acid or base is added to a buffer.
• Rounding [H⁺] heavily before taking the log.

Next steps

• Acids and bases study guide
• Acids and bases practice questions
• Equilibria and Kc study guide
• Mole revision notes
• All free 10-minute diagnostics

Official syllabus

AQA, AS and A-level Chemistry specification (7404/7405), version 1.2, July 2026, for AS and A-level exams June 2016 onwards: section 3.1.12 Acids and bases (3.1.12.1 to 3.1.12.6, A-level only), with Required practical 9.