OCR A Level Chemistry A: Acids, bases and buffers (H432) – Revision Notes

Full explanations and longer worked examples are in the Acids, bases and buffers study guide.

These notes condense section 5.1.3, Acids, bases and buffers, of the H432 specification for OCR A Level Chemistry A (Version 3.1, May 2026; first assessment 2017). Every lettered outcome, (a) through (o), is included. H432 is the full A level, and this content is part of its Module 5 on physical chemistry and transition elements. After revising, attempt the practice questions, mark off each statement on the printable OCR chemistry checklist, and use the course hub to move to the next topic. A free diagnostic will tell you where to spend your time.

Definitions to learn word for word

• Brønsted–Lowry acid: a species that donates a proton (H⁺).
• Brønsted–Lowry base: a species that accepts a proton.
• Conjugate acid–base pair: two species that differ by one H⁺, e.g. HNO₂/NO₂⁻.
• Monobasic / dibasic / tribasic: an acid that can donate 1 / 2 / 3 protons per molecule (HCl / H₂SO₄ / H₃PO₄).
• Ka (acid dissociation constant): the equilibrium constant for HA ⇌ H⁺ + A⁻; it measures the extent of dissociation.
• pH: –log[H⁺].
• Kw (ionic product of water): [H⁺][OH⁻].
• Buffer solution: a system that minimises pH changes on addition of small amounts of an acid or a base.
• End point: the point in a titration where the indicator changes colour, when [HIn] = [In⁻].

Formulas (5.1.3(c)–(g), (l))

Quantity	Formula	Units
Ka	[H⁺][A⁻] / [HA]	mol dm⁻³
pKa	–log Ka	none
Ka from pKa	10⁻ᵖᴷᵃ	mol dm⁻³
pH	–log[H⁺]	none
[H⁺] from pH	10⁻ᵖᴴ	mol dm⁻³
Kw	[H⁺][OH⁻] = 1.00 × 10⁻¹⁴ at 298 K	mol² dm⁻⁶
Strong monobasic acid	[H⁺] = [acid]
Strong base	[H⁺] = Kw / [OH⁻]
Weak acid	[H⁺] = √(Ka × [HA])
Buffer	[H⁺] = Ka × [HA] / [A⁻]

Ka is applied in the same way as Kc in How far?. A bigger Ka, or a smaller pKa, signals a stronger acid. Each step of 1 in pH is a factor of 10 in [H⁺].

Ionic equations: the role of H⁺ (5.1.3(b))

Show only H⁺ and the species it reacts with:

• metal: Mg(s) + 2H⁺(aq) → Mg²⁺(aq) + H₂(g)
• carbonate: CO₃²⁻(aq) + 2H⁺(aq) → CO₂(g) + H₂O(l)
• metal oxide: ZnO(s) + 2H⁺(aq) → Zn²⁺(aq) + H₂O(l)
• alkali: H⁺(aq) + OH⁻(aq) → H₂O(l)

Insoluble solids keep their full formulae.

Method in steps

Strong base pH (5.1.3(f)(ii))

1. Find [OH⁻]. Multiply by 2 for a base like Sr(OH)₂.
2. [H⁺] = 1.00 × 10⁻¹⁴ / [OH⁻].
3. pH = –log[H⁺].

Working backwards: a strong base solution of pH 11.65 has [H⁺] = 10⁻¹¹·⁶⁵ = 2.24 × 10⁻¹² mol dm⁻³, so [OH⁻] = (1.00 × 10⁻¹⁴) / (2.24 × 10⁻¹²) = 4.47 × 10⁻³ mol dm⁻³.

Weak acid (5.1.3(g))

1. Write Ka = [H⁺]² / [HA].
2. State the approximations: [H⁺] = [A⁻] (water’s H⁺ negligible) and [HA]equilibrium ≈ [HA]undissociated.
3. For pH: [H⁺] = √(Ka × [HA]), then –log.
4. For Ka: [H⁺] = 10⁻ᵖᴴ, then Ka = [H⁺]² / [HA].

Buffer made by partial neutralisation (5.1.3(j)(ii), (l))

1. Moles of weak acid and moles of OH⁻.
2. Moles of A⁻ formed = moles of OH⁻ added.
3. Moles of HA left = starting HA – OH⁻.
4. [H⁺] = Ka × n(HA) / n(A⁻) (the volume cancels).
5. pH = –log[H⁺].

Mass of salt for a target pH (5.1.3(l))

1. [H⁺] = 10⁻ᵖᴴ.
2. [A⁻] = Ka × [HA] / [H⁺].
3. Moles = [A⁻] × volume in dm³; mass = moles × M.

Small worked reminders

Half-neutralised acid. 30.0 cm³ of 0.180 mol dm⁻³ HZ (Ka = 7.2 × 10⁻⁵ mol dm⁻³) is mixed with 27.0 cm³ of 0.100 mol dm⁻³ NaOH.

• n(HZ) = 0.0300 × 0.180 = 0.00540 mol; n(OH⁻) = 0.0270 × 0.100 = 0.00270 mol
• n(Z⁻) = 0.00270 mol; n(HZ) left = 0.00270 mol
• Equal amounts, so pH = pKa = –log(7.2 × 10⁻⁵) = 4.14

Dilution check. Make a strong monobasic acid 100 times more dilute and its pH goes up by 2, since [H⁺] drops by a factor of 100. Do the same to a weak acid and the pH goes up by only about 1: [H⁺] follows √[HA], so it drops by a factor of about 10. A larger fraction of the weak acid dissociates as it becomes more dilute.

Units of Ka. In [H⁺][A⁻] / [HA], two concentration units on top and one underneath leave mol dm⁻³. In Kw nothing is divided, so the units are mol² dm⁻⁶.

Approximation limits (5.1.3(h))

The assumption [HA]equilibrium ≈ [HA]undissociated breaks down for “stronger” weak acids (large Ka), particularly in dilute solution. A significant fraction dissociates, so the real [HA] is lower than the starting value. The approximation then gives an [H⁺] that is too high and a pH that is too low. A useful test: if the approximate [H⁺] is more than a few per cent of the starting concentration, treat the answer with suspicion. You will not be asked to solve the quadratic.

Buffers in one box (5.1.3(i)–(m))

• Two ways to make an acid buffer: weak acid + its salt (CH₃COOH/CH₃COONa), or excess weak acid + strong alkali (excess CH₃COOH/NaOH).
• The conjugate pair: in HA ⇌ H⁺ + A⁻, the buffer keeps big reserves of HA and of A⁻.
• Add acid: H⁺ + A⁻ → HA; equilibrium shifts left.
• Add alkali: OH⁻ + H⁺ → H₂O; HA dissociates, equilibrium shifts right.
• Equal [HA] and [A⁻]: pH = pKa.
• Blood: H₂CO₃ ⇌ H⁺ + HCO₃⁻ in blood plasma holds pH between 7.35 and 7.45. Extra H⁺ reacts with HCO₃⁻ to form H₂CO₃, which leaves as CO₂ and H₂O. Extra OH⁻ removes H⁺, and H₂CO₃ dissociates to replace it.

Titration curves and indicators (5.1.3(n))

Titration	Steep section (approx.)	Equivalence pH	Indicator
strong acid / strong base	3–11	7	methyl orange or phenolphthalein
weak acid / strong base	7–11	above 7	phenolphthalein
strong acid / weak base	3–7	below 7	methyl orange
weak acid / weak base	none sharp	varies	none suitable

Commonly quoted ranges: methyl orange about 3.1–4.4 (red in acid, yellow in alkali); phenolphthalein about 8.3–10.0 (colourless in acid, pink in alkali). In a question, always use the ranges it supplies.

• Rule: the indicator’s whole range must sit inside the steep section.
• Indicator as a weak acid: HIn ⇌ H⁺ + In⁻. Acid shifts it left (colour of HIn); alkali removes H⁺ and shifts it right (colour of In⁻).
• Reading pKa from a curve: when a weak acid is titrated with a strong base, the pH halfway to equivalence equals pKa.
• Sketching: mark the start pH, the equivalence volume and the steep section clearly.

pH meter (5.1.3(o), PAG11)

Calibrate with buffers of known pH; rinse the electrode with deionised water between readings; let the reading settle; keep the temperature constant; store the electrode wet.

Must-know distinctions

• Strength and concentration: a strong acid dissociates completely; a concentrated acid simply has a large amount of acid in each dm³. So you can meet a weak acid at high concentration, or a strong one at low concentration.
• Equivalence point and end point: at equivalence, exactly enough of one solution has been run in to react with the other; the end point is simply the volume at which you see the indicator turn.
• Ka and pKa: the two move in opposite directions, so a weak acid has a small Ka but a large pKa.
• Neutral and pH 7: a neutral solution is one where [H⁺] and [OH⁻] are equal; that coincides with pH 7 only at 298 K.
• Buffer and neutral solution: a buffer resists pH change at whatever pH it is set; it need not be at pH 7.

Quick self-test

1. Calculate the pH of 0.00320 mol dm⁻³ HCl.
2. A solution has pH 4.15. Calculate [H⁺].
3. Calculate the pH of 0.0650 mol dm⁻³ NaOH at 298 K.
4. An acid has Ka = 2.9 × 10⁻⁴ mol dm⁻³. Calculate pKa.
5. An acid has pKa = 3.14. Calculate Ka.
6. A weak acid with Ka = 2.7 × 10⁻⁵ mol dm⁻³ is at 0.0800 mol dm⁻³. Calculate its pH.
7. A buffer has [HA] = [A⁻] and the acid’s pKa is 4.38. State the pH.
8. Give the conjugate base of H₂PO₄⁻ and the conjugate acid of HSO₄⁻.
9. Which indicator suits a strong acid–weak base titration, and why?
10. Explain why H₃PO₄ is tribasic but CH₃COOH is monobasic.
11. Why must a pH meter be calibrated before use?

Answers

1. pH = –log(0.00320) = 2.49
2. [H⁺] = 10⁻⁴·¹⁵ = 7.08 × 10⁻⁵ mol dm⁻³
3. [H⁺] = 1.00 × 10⁻¹⁴ / 0.0650 = 1.54 × 10⁻¹³ mol dm⁻³; pH = 12.81
4. pKa = –log(2.9 × 10⁻⁴) = 3.54
5. Ka = 10⁻³·¹⁴ = 7.24 × 10⁻⁴ mol dm⁻³
6. [H⁺] = √(2.7 × 10⁻⁵ × 0.0800) = 1.47 × 10⁻³ mol dm⁻³; pH = 2.83
7. pH = pKa = 4.38
8. HPO₄²⁻ and H₂SO₄
9. Methyl orange: the steep section runs from roughly pH 3 to 7, and methyl orange changes colour inside it; phenolphthalein changes above it.
10. H₃PO₄ can donate three protons per molecule; CH₃COOH can donate only the one on its COOH group.
11. Electrode readings drift, so calibrating against buffers of known pH makes the readings accurate.

Where marks are usually lost

• Writing pH = –log[OH⁻] for a strong base instead of using Kw.
• Forgetting to double [OH⁻] for a group 2 hydroxide such as Sr(OH)₂.
• Leaving out “proton” (or H⁺) in a Brønsted–Lowry definition, or swapping donor and acceptor.
• Giving only one approximation when asked for both weak acid approximations.
• In a partial-neutralisation buffer, using the starting moles of acid rather than the moles left.
• Saying a buffer “keeps the pH constant” instead of “minimises pH changes”.
• Explaining buffer action without saying which way the equilibrium shifts and which species removes the added H⁺ or OH⁻.
• Choosing an indicator for a weak acid–weak base titration: none is suitable.
• Including spectator ions in an ionic equation, or splitting an insoluble solid into ions.
• Giving units for pH or pKa: both have none.

Official syllabus

Source document: OCR A Level Chemistry A (code H432), specification Version 3.1 dated May 2026; first assessment 2017. These notes follow section 5.1.3 Acids, bases and buffers, in Module 5: Physical chemistry and transition elements.