AQA A-Level Chemistry: Redox and Electrochemistry — Complete Revision Guide (7405)

Redox underpins much of inorganic and biological chemistry, from the rusting of iron and the discharge of a lithium-ion cell to the electron transport chain that powers every breath you take. On AQA A-Level Chemistry (7405) the topic appears twice: once as an AS-level introduction to oxidation numbers and half-equations in §3.1.7, and again as a full A2 treatment of electrochemistry in §3.1.11. The whole topic reduces to two big ideas. The first is electron bookkeeping — track oxidation numbers, identify what is oxidised and what is reduced, and combine half-equations into balanced redox equations. The second is thermodynamics of electron transfer — standard electrode potentials measured against the standard hydrogen electrode predict the spontaneity of any redox reaction through E°cell, and ΔG° = −nFE° links electrochemistry directly to free energy.

This guide walks through both sections of the 7405 redox and electrochemistry content topic by topic. It covers oxidation numbers and balanced redox equations; half-equations and their combination; the standard hydrogen electrode and tabulated electrode potentials; electrochemical cells and cell diagram notation; EMF and the feasibility of redox reactions; commercial cells, including primary cells, rechargeable cells and hydrogen fuel cells; redox titrations using manganate(VII) and dichromate(VI); and electrolysis with the quantitative use of Faraday's laws.

The topic is also intensely synoptic. The transition-metal redox chemistry in inorganic chemistry — Fe²⁺/Fe³⁺, Mn-oxidation-state ladders, redox catalysis — is the same electrode-potential framework, only applied to d-block ions. The halogen displacement reactions are direct E° comparisons. The thermodynamic link ΔG° = −nFE° belongs to energetics, and the further bridge ΔG° = −RT ln K connects electrochemistry to the equilibrium constants you meet in kinetics and equilibrium. The alcohol oxidation reactions in organic foundations — primary alcohol to aldehyde to carboxylic acid using acidified dichromate(VI), secondary alcohol to ketone — and the halogenoalkane substitution mechanisms are redox processes whose feasibility you can now justify quantitatively. Time spent here pays back across more of the specification than almost any other single topic.

For each sub-topic you will find the core ideas, common pitfalls, a worked example and a link into the LearningBro Redox and Electrochemistry course.

What the AQA 7405 Specification Covers

AQA A-Level Chemistry (7405) is examined through Paper 1 (Inorganic and Physical, 2h, 105 marks), Paper 2 (Organic and Physical, 2h, 105 marks) and Paper 3 (a synoptic paper covering practical skills, structured questions and multiple choice, 2h, 90 marks). Redox is examined heavily on Paper 1 — virtually every Paper 1 contains a redox titration question and an electrode-potential question — and on Paper 3 through practical-based titration calculations.

Sub-topic	Spec area	Typical paper weight
Oxidation numbers and balanced redox equations	§3.1.7 (AS) and §3.1.11 (A2)	3-6 marks
Half-equations and their combination	§3.1.7	3-5 marks
Standard electrode potentials and the SHE	§3.1.11	4-6 marks
Electrochemical cells and cell diagrams	§3.1.11	3-5 marks
EMF and feasibility of redox reactions	§3.1.11	4-8 marks
Commercial cells and fuel cells	§3.1.11	4-6 marks
Redox titrations (MnO4⁻ and Cr2O7²⁻)	§3.1.7 / Practical 6	6-10 marks
Electrolysis and Faraday's laws	§3.1.11	4-6 marks

These weights are estimates modelled on recent AQA papers. What is reliable is that an extended-response electrode-potential or redox-titration question — typically eight to ten marks — appears on essentially every Paper 1.

Oxidation Numbers and Redox Equations (Extended)

The oxidation number of an atom in a species is the formal charge it would carry if every bond were ionic. The rules are short: elements in their standard state have ON = 0; in compounds, Group 1 = +1, Group 2 = +2, F = −1, O usually −2 (except in peroxides −1 and OF2 +2), H usually +1 (except in hydrides −1); the sum across a neutral compound is 0, and across an ion equals the ion charge.

Once you can assign oxidation numbers, balancing redox equations becomes mechanical. Identify the species whose ON changes; the one whose ON rises is oxidised, the one whose ON falls is reduced. The total electron loss must equal the total electron gain, which fixes the stoichiometric coefficients before you balance the spectator atoms.

Worked example. In the reaction MnO4⁻ + Fe²⁺ + H⁺ → Mn²⁺ + Fe³⁺ + H2O, Mn goes from +7 to +2 (gain of 5e⁻ per Mn) and Fe goes from +2 to +3 (loss of 1e⁻ per Fe). To balance electrons, 1 Mn requires 5 Fe. Balancing oxygens with water and hydrogens with H⁺ gives MnO4⁻ + 5 Fe²⁺ + 8 H⁺ → Mn²⁺ + 5 Fe³⁺ + 4 H2O. Check: charges on left = (−1) + 5(+2) + 8(+1) = +17; charges on right = (+2) + 5(+3) + 0 = +17. Balanced.

A common pitfall is to forget that O is −1 in peroxides; another is to apply the +1 rule for H in metal hydrides. A third is to balance atoms before balancing electrons, which forces a redo. See the oxidation numbers lesson.

Redox Reactions and Half-Equations

A half-equation shows either the oxidation half or the reduction half of a redox process, with electrons written explicitly. The reduction half of permanganate in acidic solution is MnO4⁻ + 8 H⁺ + 5 e⁻ → Mn²⁺ + 4 H2O; the oxidation half for iron(II) is Fe²⁺ → Fe³⁺ + e⁻. To balance a half-equation: balance the redox atom first, balance oxygens with water, balance hydrogens with H⁺ (in acidic conditions) or OH⁻/H2O (in alkaline conditions), then balance charge with electrons.

To combine two half-equations into an overall redox equation, multiply each through so that the electrons cancel exactly. The MnO4⁻ half (5 e⁻ gained) combines with five copies of the Fe²⁺ half (1 e⁻ lost each), giving exactly the equation derived above.

Worked example. Combine Cr2O7²⁻ + 14 H⁺ + 6 e⁻ → 2 Cr³⁺ + 7 H2O with Fe²⁺ → Fe³⁺ + e⁻. Multiply the iron half by 6 to match 6 electrons. Overall: Cr2O7²⁻ + 6 Fe²⁺ + 14 H⁺ → 2 Cr³⁺ + 6 Fe³⁺ + 7 H2O. Charge check: left (−2) + 6(+2) + 14(+1) = +24; right 2(+3) + 6(+3) + 0 = +24. Correct.

A common pitfall is to leave electrons in the final combined equation. Another is to add hydrogens before water has balanced the oxygens. See the half-equations lesson.

Standard Electrode Potentials

A standard electrode potential E° is the potential of a half-cell measured under standard conditions (298 K, all aqueous species at 1.00 mol dm⁻³, all gases at 100 kPa, all solids pure) against the standard hydrogen electrode (SHE). The SHE is defined to have E° = 0.00 V exactly and consists of H⁺(aq, 1.00 mol dm⁻³) in contact with H2(g, 100 kPa) at an inert platinum electrode coated with platinum black.

Half-equations are tabulated as reductions, with the most negative E° at the top. A more positive E° means a stronger oxidising agent on the left of the reduction half; a more negative E° means a stronger reducing agent on the right. Worked snippet from the AQA data sheet:

Reduction half	E° / V
Li⁺ + e⁻ → Li	−3.04
Zn²⁺ + 2 e⁻ → Zn	−0.76
Fe²⁺ + 2 e⁻ → Fe	−0.44
2 H⁺ + 2 e⁻ → H2	0.00
Cu²⁺ + 2 e⁻ → Cu	+0.34
Fe³⁺ + e⁻ → Fe²⁺	+0.77
MnO4⁻ + 8 H⁺ + 5 e⁻ → Mn²⁺ + 4 H2O	+1.51
F2 + 2 e⁻ → 2 F⁻	+2.87

The SHE is the experimental anchor on which the whole table — and so the whole feasibility framework — sits. A common pitfall is to assume the SHE has a "true" zero potential rather than a defined one. Another is to expect E° values to scale with stoichiometry: doubling a half-equation does not double E°, because E° is an intensive property. See the standard electrode potentials lesson.

Electrochemical Cells and Cell Diagrams

An electrochemical cell is built by connecting two half-cells through an external circuit and a salt bridge. Electrons flow externally from the more negative electrode (the anode, where oxidation occurs) to the more positive electrode (the cathode, where reduction occurs). The salt bridge — typically saturated KNO3 or KCl in agar gel — carries ions to maintain charge neutrality without allowing the two solutions to mix.

A cell diagram uses a compact notation. The convention is: anode species (most reduced first) | anode species (most oxidised) || cathode species (most oxidised first) | cathode species (most reduced). A single vertical line is a phase boundary; a double vertical line is the salt bridge.

Worked example. For a zinc-copper cell with Zn(s) | Zn²⁺(aq) on the left and Cu²⁺(aq) | Cu(s) on the right, the diagram is Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s). E°cell = E°(right) − E°(left) = (+0.34) − (−0.76) = +1.10 V.

Where a half-cell contains no solid metal — for example the Fe²⁺/Fe³⁺ couple — an inert platinum electrode is used, written explicitly: Pt | Fe²⁺(aq), Fe³⁺(aq) ||. The comma separates two aqueous species sharing the same phase. A common pitfall is to write the species in the wrong order around the salt bridge; another is to omit the inert electrode for a fully aqueous couple. See the electrochemical cells lesson.

EMF and Feasibility of Redox Reactions

The EMF (or cell potential) E°cell is the difference between the two standard electrode potentials, calculated as E°(reduction half) − E°(oxidation half), equivalently E°(right) − E°(left) when the cell diagram is written conventionally. If E°cell is positive, the reaction written left-to-right is thermodynamically feasible under standard conditions; if negative, it is not feasible as written and the reverse reaction is feasible.

The link to free energy is ΔG° = −nFE°cell, where n is the number of electrons transferred per formula unit of reaction and F is the Faraday constant (96 485 C mol⁻¹). A positive E°cell gives a negative ΔG°, the thermodynamic condition for spontaneity. A further synoptic link uses ΔG° = −RT ln K to derive the equilibrium constant of the redox reaction from the cell potential.

Worked example. For the cell Zn(s) | Zn²⁺ || Cu²⁺ | Cu(s) with E°cell = +1.10 V and n = 2, ΔG° = −(2)(96 485)(1.10) = −2.12 × 10⁵ J mol⁻¹ = −212 kJ mol⁻¹. The large negative value confirms strong feasibility — zinc spontaneously reduces copper(II).

Feasibility predictions are thermodynamic, not kinetic. A reaction may be feasible yet imperceptibly slow because of high activation energy: the classic case is the oxidation of water by oxygen in acidic solution, predicted feasible by E° but practically inert. Non-standard concentrations also shift the actual cell potential; rough rule, a tenfold deviation in concentration ratio shifts E by ~0.06/n V per the Nernst expression. A common pitfall is to call a reaction "impossible" rather than "not feasible under standard conditions". See the EMF and feasibility lesson.

Commercial Cells and Fuel Cells

The 7405 specification names three classes of commercial cell. Non-rechargeable (primary) cells such as the alkaline zinc-manganese dioxide cell deliver charge from a fixed mass of reactants and are then discarded; their EMF is typically 1.5 V. Rechargeable cells — the nickel-cadmium and especially the lithium-ion cell — are designed so that the cell reaction is reversed when current is passed in the opposite direction, regenerating the original reactants. The lithium-ion cell uses a lithium cobalt oxide cathode and a graphite anode; lithium ions shuttle between the two during charge and discharge. Cell EMF is around 3.6 V, energy density is high, and cycle life now reaches thousands of full charge-discharge cycles.

A hydrogen fuel cell is a non-rechargeable cell of a different kind: the reactants flow continuously through the cell and the products are removed. In the alkaline version, the anode half is H2 + 2 OH⁻ → 2 H2O + 2 e⁻ and the cathode half is ½ O2 + H2O + 2 e⁻ → 2 OH⁻; overall H2 + ½ O2 → H2O with E°cell = +1.23 V. The advantage over a hydrocarbon combustion engine is twofold: the only product is water, and the cell converts chemical energy directly to electrical energy without a heat-engine Carnot limit. Disadvantages are storage and transport of hydrogen (highly flammable, low volumetric energy density unless compressed or liquefied), the carbon footprint of the hydrogen production route (steam-methane reforming dominates today; only electrolytic "green" hydrogen is carbon-neutral, and only when the electricity is renewable), and the cost and durability of the platinum-loaded electrocatalysts.

A common pitfall is to write the fuel-cell half-equations in acidic form when the question specifies alkaline electrolyte, or vice versa. Another is to claim that a fuel cell is "100 percent efficient" — practical efficiencies are 40-60 percent. See the commercial cells and fuel cells lesson.

Redox Titrations

A redox titration uses one redox species of known concentration to determine the amount of a second redox species. The two standard AQA titrants are acidified potassium manganate(VII) (KMnO4) and acidified potassium dichromate(VI) (K2Cr2O7).

Manganate(VII) is its own indicator: the deep purple MnO4⁻ is reduced to almost colourless Mn²⁺, so the end point is the first persistent pink tinge after the analyte is exhausted. The half-equation MnO4⁻ + 8 H⁺ + 5 e⁻ → Mn²⁺ + 4 H2O fixes the 1:5 stoichiometry against any one-electron reducing agent such as Fe²⁺.

Worked example. 25.0 cm³ of a Fe²⁺ solution is titrated against 0.0200 mol dm⁻³ KMnO4 in excess dilute H2SO4; the end point is reached at 24.8 cm³. Moles MnO4⁻ = 0.0200 × 24.8 / 1000 = 4.96 × 10⁻⁴ mol. Moles Fe²⁺ = 5 × 4.96 × 10⁻⁴ = 2.48 × 10⁻³ mol. [Fe²⁺] = 2.48 × 10⁻³ / 0.0250 = 0.0992 mol dm⁻³.

Dichromate(VI) requires a separate indicator (often sodium diphenylaminesulfonate) because the orange-to-green colour change is not sharp at the end point. The half-equation Cr2O7²⁻ + 14 H⁺ + 6 e⁻ → 2 Cr³⁺ + 7 H2O fixes a 1:6 stoichiometry against Fe²⁺.

A common pitfall is to use HCl rather than H2SO4 as the acidifying acid for permanganate titrations: chloride is oxidised by MnO4⁻, consuming titrant and inflating the apparent analyte concentration. Another is to forget that the iron must be entirely as Fe²⁺ before titration; any pre-oxidation by air must be reversed by adding a stoichiometric reductant. See the redox titrations lesson.

Electrolysis and Faraday's Laws

Electrolysis is the use of an external EMF to force a non-spontaneous redox reaction. The cell potential applied must exceed the magnitude of the spontaneous E°cell of the reverse reaction (plus an overpotential to overcome kinetic barriers). The cation moves to the cathode and is reduced; the anion moves to the anode and is oxidised.

The quantitative laws are due to the work of Faraday in the 1830s. The first law says the mass of substance deposited or liberated at an electrode is proportional to the quantity of charge passed: Q = I × t (charge in coulombs, current in amperes, time in seconds). The second law says that to deposit one mole of a species requires nF coulombs, where n is the number of electrons in the half-equation and F = 96 485 C mol⁻¹.

Combining the two: moles deposited = Q / (nF) = (I × t) / (nF), and mass deposited = (I × t × M) / (nF) where M is the molar mass.

Worked example. A current of 1.50 A is passed for 1 hour through molten aluminium oxide. How much aluminium is deposited at the cathode? Half-equation Al³⁺ + 3 e⁻ → Al, so n = 3. Q = 1.50 × 3600 = 5400 C. Moles Al = 5400 / (3 × 96 485) = 0.01866 mol. Mass Al = 0.01866 × 27.0 = 0.504 g.

A common pitfall is to use n = 1 by default; n is set by the half-equation, not by the metal. Another is to forget to convert minutes or hours into seconds before computing Q. See the electrolysis and Faraday lesson.

Industrial Applications

Three large-scale electrochemical processes anchor the topic in real chemistry. The Hall-Héroult process, devised independently by Hall and Héroult in 1886, extracts aluminium by electrolysis of alumina dissolved in molten cryolite at around 950 °C. Steel-lined carbon-anode cells operate at currents of order 300 kA per cell, and the world's smelters consume something like 2 percent of global electricity to produce ~70 megatonnes of aluminium a year. The chlor-alkali process electrolyses concentrated brine in membrane cells to produce chlorine (at the anode), hydrogen (at the cathode) and sodium hydroxide (in the catholyte) — three commodity chemicals from one electrolytic unit. Hydrogen fuel cells power buses, forklifts and an increasing number of long-haul lorries; the chemistry is identical to the textbook half-cells above, only scaled up and engineered for durability.

A connection that often surprises students: a closely related fixed-potential framework, the Pourbaix diagram, plots stable redox species against pH and electrode potential, and is the working language of corrosion engineering and the metallurgy of d-block extraction. The same E°cell calculations you do for a Paper 1 question are, on a different page, the design tool that decides whether a steel pipeline survives in seawater. And on the most exotic edge of the same physics, the BCS theory of superconductivity by Bardeen, Cooper and Schrieffer in 1957 describes how electrons in a metal can flow with literally zero resistance — a phenomenon that is now used to build the magnetic-resonance imaging scanners that themselves rely on tightly controlled redox chemistry to operate.

Common Mark-Loss Patterns

• Forgetting that O in peroxides has ON = −1 (and H in hydrides has ON = −1).
• Balancing atoms before balancing electrons in a redox equation.
• Leaving electrons in the final combined redox equation.
• Doubling E° when a half-equation is doubled (E° is intensive).
• Writing E°cell as E°(left) − E°(right) instead of E°(right) − E°(left).
• Omitting the inert platinum electrode for an aqueous-only couple.
• Calling a slow but thermodynamically feasible reaction "impossible".
• Using HCl with acidified permanganate (chloride is oxidised).
• Using n = 1 by default in Faraday calculations.
• Forgetting to convert minutes/hours to seconds in Q = It.
• Confusing alkaline and acidic versions of the fuel-cell half-equations.
• Quoting "100 percent efficient" for a fuel cell.

How to Revise This Topic

• Drill oxidation-number assignment until it is automatic — twenty species a day for a fortnight.
• Practise balancing half-equations and combining them; aim for ten balanced overall equations a session.
• Memorise the SHE definition and the top, middle and bottom of the AQA data sheet so you can spot the strong oxidising and reducing agents instantly.
• Build a cell-diagram flashcard set for every named couple in the spec.
• Drill E°cell and ΔG° = −nFE° calculations in mixed batches with feasibility questions.
• Practise redox-titration arithmetic with both MnO4⁻ (1:5 with Fe²⁺) and Cr2O7²⁻ (1:6 with Fe²⁺).
• Drill Faraday's laws by computing the mass of metal deposited in mixed-n problems (n = 1 for Ag, n = 2 for Cu, n = 3 for Al).
• Use the LearningBro practice quizzes to test under timed conditions, and the AI tutor for feedback on extended-response electrode-potential questions.

Linking to Other Topics

Redox and electrochemistry sits at the junction of physical chemistry, inorganic chemistry and organic chemistry. The transition-metal chemistry in inorganic chemistry is essentially applied electrode-potential theory: the colours of Fe²⁺/Fe³⁺ solutions, the disproportionation of Cu⁺, the redox titrations of the d-block all sit on the same E° table. The thermodynamic link ΔG° = −nFE° belongs to energetics and lets you derive Gibbs energies from cell potentials without calorimetry. The further bridge ΔG° = −RT ln K — together with the kinetics and equilibrium framework — converts cell potentials into equilibrium constants for redox reactions. The alcohol oxidation reactions and halogenoalkane substitutions in organic foundations are redox processes whose feasibility is now quantitatively justifiable. Time spent here returns in every other physical-chemistry section.

Final Word

Redox and electrochemistry is calculation-heavy, but the calculations come from a very short list: oxidation-number tracking, half-equation balancing, E°cell = E°(right) − E°(left), ΔG° = −nFE°, and Q = It with moles = Q/(nF). Master those, drill the manganate(VII) and dichromate(VI) titration stoichiometries, and learn the alkaline and acidic forms of the fuel-cell half-equations cold. The full LearningBro Redox and Electrochemistry course walks through every sub-topic with worked examples and AI tutor feedback. Get this section fluent and a substantial fraction of the physical chemistry on every paper becomes routine.