OxfordAQA A-Level Chemistry: Acids and bases (9620) – Revision Notes

These notes condense section 3.1.10, Acids and bases, of the OxfordAQA International AS and A-level Chemistry (9620) specification, Version 5.3, for International AS and A-level exams May/June 2020 onwards: 3.1.10.1 to 3.1.10.6 and required practical 7. All of it is International A2 material (International A-level only), assessed in Unit 3, with Unit 5 able to draw on any topic.

The study guide explains each idea in full with worked examples; the practice questions let you check yourself. Kw and Ka rest on equilibria. Use the course hub, the specification checklist and our diagnostics to track gaps.

Definitions (3.1.10.1 to 3.1.10.4)

• Acid: proton donor. Base: proton acceptor.
• Acid–base equilibrium: one in which protons (H⁺) are transferred. Check both directions: the species that loses H⁺ in the reverse reaction is also an acid.
• Strong acid or base: completely dissociated in water.
• Weak acid or base: dissociates only slightly in water.
• pH = −log₁₀[H⁺].
• Kw = [H⁺][OH⁻], the ionic product of water; it comes from Kc for water ionising.
• Ka: dissociation constant for a weak acid; pKa = −log₁₀ Ka.
• Neutral: [H⁺] = [OH⁻]. This is pH 7 only at 25 °C.

Equations at a glance

Situation	Relationship	Units of the constant or answer
Any solution	pH = −log₁₀[H⁺], so [H⁺] = 10^(−pH)	pH has none; [H⁺] in mol dm⁻³
Water equilibrium	Kw = [H⁺][OH⁻]	Kw in mol² dm⁻⁶
Pure water	[H⁺] = √Kw = [OH⁻]	mol dm⁻³
Fully dissociated base	[H⁺] = Kw ÷ [OH⁻]	mol dm⁻³
Weak acid equilibrium	Ka = [H⁺][A⁻] ÷ [HA]	Ka in mol dm⁻³
Weak acid on its own	[H⁺] = √(Ka[HA])	mol dm⁻³
Strength scale	pKa = −log₁₀ Ka and Ka = 10^(−pKa)	pKa has none
Buffer (acidic)	[H⁺] = Ka[HA] ÷ [A⁻]	mol dm⁻³
Half-way to equivalence	pH = pKa	none

Kw = 1.00 × 10⁻¹⁴ mol² dm⁻⁶ at 25 °C. Use the value a question gives at any other temperature.

Why Kw = [H⁺][OH⁻]

H₂O ⇌ H⁺ + OH⁻        Kc = [H⁺][OH⁻] / [H₂O]

[H₂O] is so large compared with the ions that it is effectively constant, so Kc × [H₂O] = Kw. The dissociation is endothermic, so raising the temperature moves the equilibrium right and Kw increases. Pure water then has pH below 7 but stays neutral.

Method boxes

pH of a strong acid (monoprotic)

1. [H⁺] = concentration of the acid.
2. pH = −log₁₀[H⁺].

pH of a strong base

1. [OH⁻] = concentration × number of OH⁻ per formula unit.
2. [H⁺] = Kw / [OH⁻].
3. pH = −log₁₀[H⁺].

pH of a weak acid from Ka

1. Assume [H⁺] = [A⁻] and [HA] ≈ starting concentration.
2. [H⁺] = √(Ka × [HA]).
3. pH = −log₁₀[H⁺].

Ka from the pH of a weak acid

1. [H⁺] = 10^(−pH).
2. Ka = [H⁺]² / [HA].
3. pKa = −log₁₀ Ka if asked.

Strong acid mixed with strong base

1. Work out n(H⁺) and n(OH⁻) from concentration × volume in dm³.
2. Take the smaller from the larger; whichever is left over decides the pH.
3. Convert that leftover to a concentration using the combined volume.
4. Leftover H⁺ gives pH at once; leftover OH⁻ needs Kw to reach [H⁺].

Acidic buffer pH

1. Get the final amounts of HA and A⁻, allowing for any base or acid that has reacted.
2. [H⁺] = Ka × n(HA) / n(A⁻).
3. pH = −log₁₀[H⁺].

Small worked reminders

• 0.0055 mol dm⁻³ HNO₃: pH = −log₁₀(0.0055) = 2.26.
• A weak acid with Ka = 2.0 × 10⁻⁶ mol dm⁻³ at 0.500 mol dm⁻³: [H⁺] = √(1.0 × 10⁻⁶) = 1.0 × 10⁻³, so pH = 3.00.
• Buffer with n(HA) = n(A⁻): [H⁺] = Ka, so pH = pKa.

Titrations, pH curves and indicators (3.1.10.5)

Titration results: use only concordant titres (usually within 0.10 cm³) for the mean, and ignore the rough titre.

Acid + base	Start pH (acid in flask)	Vertical section	Equivalence pH	Indicator
Strong + strong	about 1	about 3 to 11	7	methyl orange or phenolphthalein
Weak + strong	higher	about 7 to 11	above 7	phenolphthalein
Strong + weak	about 1	about 3 to 7	below 7	methyl orange
Weak + weak	higher	none	hard to find	none: use a pH meter

Indicator ranges to know: methyl orange 3.1–4.4 (red to yellow); phenolphthalein 8.3–10.0 (colourless to pink).

Rule: an indicator is suitable if its entire colour-change range sits inside the vertical section.

Features to explain on a sketch

• Weak acid starts higher: only slightly dissociated.
• Buffer region early in a weak acid–strong base titration: HA and A⁻ both present.
• pH = pKa at half the equivalence volume, because [HA] = [A⁻].
• Equivalence above 7 with a weak acid and strong base: A⁻ removes H⁺ from water, leaving OH⁻.
• Equivalence below 7 with a strong acid and weak base: the cation of the salt donates H⁺ to water.

Required practical 7

Aim: record pH while a strong base is run into a weak acid, then while a strong acid is run into a weak base.

• Calibrate the pH meter with buffers of known pH first.
• Measure the solution in the beaker with a pipette; add the other from a burette.
• Stir, wait for a steady reading, then record pH and total volume.
• Use smaller additions near the end point to define the steep section.
• Plot pH (y-axis) against volume added (x-axis). Equivalence volume is at the middle of the steep section.
• For the weak acid run, read pH at half the equivalence volume to get pKa.

Buffers (3.1.10.6)

Buffer: a solution whose pH stays nearly constant when it is diluted or when a little acid or base is added.

Type	Contains	Example
Acidic	weak acid + its salt	HA with NaA, or HA partly neutralised by NaOH
Basic	weak base + its salt	methylamine with methylammonium chloride

Acidic buffer, qualitative

• Add H⁺: H⁺ + A⁻ → HA. Equilibrium HA ⇌ H⁺ + A⁻ shifts left.
• Add OH⁻: OH⁻ + H⁺ → H₂O; HA dissociates to replace H⁺. Equilibrium shifts right.
• Large reservoirs of HA and A⁻ mean the ratio [HA]/[A⁻] hardly changes.

Basic buffer, qualitative

• Add H⁺: base + H⁺ → conjugate acid (CH₃NH₂ + H⁺ → CH₃NH₃⁺).
• Add OH⁻: conjugate acid + OH⁻ → base + H₂O.

Applications: calibrating pH meters; controlling the pH of blood; controlling acidity in foods and drinks; any process where enzymes or reactions need a steady pH.

Must-know distinctions

• Strong vs concentrated: “strong” describes the extent of dissociation (complete); “concentrated” describes how many moles are in each dm³.
• Weak acid vs buffer: for a weak acid alone [H⁺] = [A⁻]; in a buffer [A⁻] is much larger, so [H⁺] = Ka[HA]/[A⁻].
• End point vs equivalence point: the end point is where the indicator changes; the equivalence point is where reacting amounts match.
• Ka vs pKa: larger Ka means stronger; larger pKa means weaker.

Quick self-test

1. In HCO₃⁻ + OH⁻ → CO₃²⁻ + H₂O, which species is the Brønsted–Lowry acid?
2. Calculate the pH of 0.0072 mol dm⁻³ HI.
3. Calculate [H⁺] in a solution of pH 10.46.
4. Calculate the pH of 0.0028 mol dm⁻³ NaOH at 25 °C.
5. Ka = 4.9 × 10⁻¹⁰ mol dm⁻³. Calculate pKa.
6. pKa = 2.95. Calculate Ka.
7. A weak acid has Ka = 7.2 × 10⁻⁶ mol dm⁻³. Calculate the pH of a 0.0350 mol dm⁻³ solution.
8. Which of methyl orange and phenolphthalein suits a strong acid–weak base titration?
9. In a pH curve, the pH at half the equivalence volume is 5.13. State pKa.
10. A buffer contains 0.024 mol HA and 0.016 mol A⁻ (Ka = 3.0 × 10⁻⁵ mol dm⁻³). Calculate its pH.
11. At 0 °C, Kw is about 1.12 × 10⁻¹⁵ mol² dm⁻⁶. Calculate the pH of pure water and say whether it is neutral.

Answers

1. HCO₃⁻ (it donates a proton to OH⁻).
2. pH = −log₁₀(0.0072) = 2.14.
3. [H⁺] = 10^(−10.46) = 3.47 × 10⁻¹¹ mol dm⁻³.
4. [H⁺] = 1.00 × 10⁻¹⁴ / 0.0028 = 3.57 × 10⁻¹², so pH = 11.45.
5. pKa = 9.31.
6. Ka = 1.12 × 10⁻³ mol dm⁻³.
7. [H⁺] = √(7.2 × 10⁻⁶ × 0.0350) = 5.02 × 10⁻⁴, so pH = 3.30.
8. Methyl orange: its range lies in the acidic vertical section.
9. pKa = 5.13.
10. [H⁺] = 3.0 × 10⁻⁵ × 0.024 / 0.016 = 4.5 × 10⁻⁵, so pH = 4.35.
11. [H⁺] = √(1.12 × 10⁻¹⁵) = 3.35 × 10⁻⁸, so pH = 7.48; still neutral because [H⁺] = [OH⁻].

Where marks are usually lost

• Leaving H₂O in the Kw expression, or giving Kw units of mol dm⁻³ instead of mol² dm⁻⁶.
• Forgetting that a hydroxide with two OH⁻ per formula unit doubles [OH⁻].
• Using √(Ka × [HA]) in a buffer calculation.
• Using only one solution’s volume, rather than the combined volume, for the leftover reagent.
• Describing an indicator as suitable when only part of its range lies in the vertical section.
• Stating the pH of neutral water is always 7.
• Writing the buffer explanation without naming which equilibrium shifts and in which direction.
• Missing the half-equivalence trick (pH = pKa) when a question gives a pH curve.

Official syllabus

OxfordAQA International AS and A-level Chemistry (9620) specification, Version 5.3, for International AS and A-level exams May/June 2020 onwards, section 3.1.10 Acids and bases (International A2), with required practical 7. Published by OxfordAQA.