Acids and bases (A-level only)

Overview

The supplied curriculum label, 3.1.12 Acids and bases (A-level only), matches AQA A-level Chemistry 7405 rather than GCSE Chemistry. This guide therefore teaches the A-level content directly; GCSE students can use it as Higher-tier enrichment, but K_a, K_w, buffer calculations and full pH curves are A-level extensions. AQA places the topic in physical chemistry and assesses it in A-level Paper 1, so candidates need secure knowledge and reliable multi-step application. 1

At its heart, this topic is about moving protons and measuring the effect. You will identify Brønsted–Lowry acids and bases, calculate pH from concentrations, use the ionic product of water, calculate the pH of weak acids and buffers, and interpret titration curves. Questions may look numerical, practical or explanatory, but they all reward the same habits: identify the chemical system first, choose the correct relationship, show every substitution, and use precise equilibrium language. Candidates who can explain why an equilibrium shifts, rather than merely state that it does, gain the explanatory marks.

Key Concepts

1. Brønsted–Lowry acids, bases and conjugate pairs

A Brønsted–Lowry acid is a proton donor; a Brønsted–Lowry base is a proton acceptor. The proton is H^+. In an acid–base reaction, one species loses a proton and another gains it. The species formed when an acid loses H^+ is its conjugate base. The species formed when a base gains H^+ is its conjugate acid.

For ethanoic acid in water:

CH_3COOH(aq) + H_2O(l) \rightleftharpoons CH_3COO^-(aq) + H_3O^+(aq)

Ethanoic acid donates H^+, so it is the acid; ethanoate is its conjugate base. Water accepts H^+, so it is the base; H_3O^+ is its conjugate acid. The reversible arrow matters: weak acids establish an equilibrium rather than dissociating completely. Credit is given for pairing the acid with the species formed after it donates a proton, not merely for listing four labels. 1

Memory hook — BAAD: Bases Accept; Acids Donate. When you identify a species, track the proton: if it leaves, that species was the acid; if it arrives, that species was the base.

2. pH is logarithmic, not linear

The pH scale makes very small hydrogen ion concentrations manageable:

pH = -\log_{10}[H^+]

A pH change of one is a ten-fold change in [H^+], not a one-unit change. A solution at pH 3 has ten times the hydrogen ion concentration of a solution at pH 4. When pH is given, reverse the calculation using [H^+] = 10^{-pH}.

For a strong monoprotic acid, assume complete dissociation. Therefore 0.0500 mol,dm^{-3} nitric acid gives [H^+] = 0.0500,mol,dm^{-3}, so pH = -\log_{10}(0.0500) = 1.30. Do not make this shortcut for a weak acid: its initial acid concentration is not its hydrogen ion concentration.

Strength and concentration are different ideas. Strength is the extent of dissociation; concentration is the amount dissolved per dm^3. A concentrated weak acid may contain more H^+ than a very dilute strong acid. Candidates lose marks when they call a strong acid “strong because it is concentrated”.

3. Water and $K_w$: use the temperature given

Water is slightly dissociated:

H_2O(l) \rightleftharpoons H^+(aq) + OH^-(aq)

The ionic product of water is:

K_w = [H^+][OH^-]

At 298 K, K_w = 1.00 \times 10^{-14},mol^2,dm^{-6}. A strong base supplies OH^-, so first calculate [OH^-], then rearrange to [H^+] = \frac{K_w}{[OH^-]}, and finally calculate pH. For 0.150 mol,dm^{-3} Ba(OH)_2, two hydroxide ions are released per formula unit, giving [OH^-] = 0.300,mol,dm^{-3}. Missing this factor of two is a common lost mark.

K_w varies with temperature. Neutrality means [H^+] = [OH^-]; it does not always mean pH 7. At a higher temperature, pure water can have a pH below 7 and still be neutral because both ion concentrations have increased by the same amount. A precise explanation names both the changed K_w and the equal ion concentrations. 1

4. Weak acids, $K_a$ and $pK_a$

A weak acid only partially dissociates. For a general weak acid HA:

HA(aq) \rightleftharpoons H^+(aq) + A^-(aq)

K_a = \frac{[H^+][A^-]}{[HA]}

A larger K_a means the equilibrium lies further right and the acid dissociates more: the acid is stronger. pK_a = -\log_{10}K_a, so a stronger acid has a larger K_a but a smaller pK_a.

In most A-level weak-acid calculations, make two stated approximations. First, only a small amount of HA dissociates, so equilibrium [HA] \approx initial [HA]. Second, every dissociated HA produces one H^+ and one A^-, so [H^+] \approx [A^-]. The expression becomes:

K_a = \frac{[H^+]^2}{[HA]}, so [H^+] = \sqrt{K_a[HA]}

Use this only for a weak acid when the approximation is reasonable. The reliable workflow is: write the formula; substitute with brackets; calculate [H^+]; then apply the negative logarithm. Never take -\log before taking the square root.

Memory hook — “To find the H, root the K times A.” It prompts [H^+] = \sqrt{K_a \times [HA]} for the simplified weak-acid calculation.

5. Titration curves, equivalence points and indicators

A pH curve plots pH on the y-axis against volume of titrant added on the x-axis. It shows the gradual and rapid pH changes during a titration. The equivalence point is the stoichiometric point where the reacting acid and base have exactly neutralised each other according to the balanced equation. It is not automatically pH 7.

Select an indicator whose colour-change range lies inside the vertical section. The endpoint is the observed colour change; the equivalence point is the exact stoichiometric neutralisation point. Examiners can test this distinction in a practical or graph question. At the half-neutralisation point of a weak acid–strong base titration, half the acid has become conjugate base, so [HA] = [A^-]. Therefore [H^+] = K_a and pH = pK_a. 1

Memory hook — “Half-way there, pH is pK_a.” Find the equivalence volume, halve it, then read the pH at that half-volume.

6. Buffers: two reservoirs working together

A buffer maintains an approximately constant pH after dilution or addition of a small amount of acid or base. An acidic buffer contains a weak acid and the salt of that acid, for example CH_3COOH and sodium ethanoate. The salt supplies a high concentration of CH_3COO^-; the weak acid supplies CH_3COOH.

If a small amount of strong acid is added, the added H^+ reacts with ethanoate ions:

CH_3COO^-(aq) + H^+(aq) \rightarrow CH_3COOH(aq)

This removes most added H^+; the equilibrium shifts left. If a small amount of alkali is added, hydroxide ions react with the weak acid:

CH_3COOH(aq) + OH^-(aq) \rightarrow CH_3COO^-(aq) + H_2O(l)

This removes most added OH^-; the weak acid dissociation equilibrium shifts right to replace some lost H^+. For full explanatory credit, name the species that reacts, write or describe the reaction, and state the effect on [H^+] or pH. Avoid the vague phrase “the buffer neutralises it” without identifying how.

For an acidic buffer, rearrange the K_a expression:

[H^+] = K_a \times \frac{[HA]}{[A^-]}

If both components are mixed into one final volume, moles can be used in the ratio because the common final volume cancels. However, candidates must first ensure they have used the amounts after any neutralisation reaction, not simply the original label concentrations.

Mathematical and Scientific Relationships

Unit conversion check: 1000,cm^3 = 1,dm^3; convert a volume in cm^3 to dm^3 before using n = cV. Concentration is normally mol,dm^{-3}; K_w has units mol^2,dm^{-6} in the concentration treatment used here. pH and pK_a have no units. In a pH answer, two decimal places is normally appropriate because it reflects the logarithmic calculation.

Required Practical 9: pH curves from titrations

AQA Required Practical 9 investigates how pH changes when a weak acid reacts with a strong base and when a strong acid reacts with a weak base. It connects apparatus technique, graph plotting, indicator choice and equilibrium reasoning. 1

Common errors include failing to calibrate the pH meter, reading pH before mixing is complete, using large titrant additions near equivalence, leaving air bubbles in the burette tip, and plotting the axes the wrong way round. In written papers, examiners may ask why smaller additions are needed near the endpoint, how to improve repeatability, why a pH meter is better than a colour change for continuous data, or which indicator is justified by the curve. A strong answer links each improvement to the error it reduces.

Exam-Focused Checklist

Before submitting a calculation, candidates should check four things. First, have you identified the system as strong acid, strong base, weak acid or buffer? Second, have you used the correct concentrations after dilution, stoichiometry or the number of hydroxide ions released? Third, have you shown the formula, substitution and intermediate value so method marks can be awarded? Finally, does the answer make chemical sense: pH below 7 for acids, above 7 for bases at 298 K, with a realistic magnitude?

For a graph question, describe the data before explaining it. For example: “The pH rises gradually from about 3, then rises steeply at 25.0 cm^3.” Then explain: “The initial gradual region is a buffer because both weak acid and ethanoate are present; the steep change occurs as the weak acid is neutralised.” This separates observation from chemical reasoning and makes every marking point easy to credit.

GCSE Bridge and Tier Note

There is no GCSE Foundation/Higher tier split for this topic because AQA 3.1.12 is A-level only. GCSE knowledge of neutralisation, indicators, the pH scale and simple acid reactions is useful prior knowledge. The A-level step-up is the use of Brønsted–Lowry definitions, logarithms, equilibrium constants, temperature-dependent K_w, buffers and fully interpreted pH curves. If you are revising GCSE only, prioritise the bridge concepts and use the rest as extension work.

References