Reactions where electrons change hands. Batteries, corrosion, biology — all redox, all the time.
The required syllabus content for Unit 9, in order. Each card is one lesson-sized checkpoint.
Lesson 1 of Unit 9.
Lesson 2 of Unit 9.
Lesson 3 of Unit 9.
Lesson 5 of Unit 9.
In a primary cell: reactants are consumed – cannot reverse the reaction - Anode disintegrates due to oxidation
Fuel converted to water, any other gas, and heat
Lesson 9 of Unit 9.
Recap: Secondary cells have reversible electrochemical processes so are rechargeable
Each lesson card below mirrors the original teacher deck — syllabus refs, content, worked examples and practice questions in order.
A redox reaction is one in which electrons are transferred. OIL RIG: Oxidation Is Loss, Reduction Is Gain of electrons.
Three equivalent definitions:
| Oxidation | Reduction | |
|---|---|---|
| Electrons | Lost | Gained |
| Oxygen | Gained | Lost |
| Hydrogen | Lost | Gained |
| Oxidation state | Increases | Decreases |
The oxidising agent causes oxidation by accepting electrons — and is itself reduced. The reducing agent causes reduction by donating electrons — and is itself oxidised.
A half-equation shows just one half of a redox reaction (either the oxidation or the reduction), including the electrons explicitly.
Multiply each half-equation so the electrons cancel when summed. Then add and simplify (cancel any H⁺ or H₂O appearing on both sides).
A more reactive species displaces a less reactive one from its compound. Two kinds:
Reactivity series (top = most reactive): K, Na, Ca, Mg, Al, (C), Zn, Fe, Pb, (H), Cu, Ag, Au. A metal will displace any metal below it from a salt of the latter.
Example: Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s). Zn is above Cu in the series; Cu²⁺ takes Zn's electrons.
Reactivity: F > Cl > Br > I. A halogen displaces any halide ion below it.
Example: Cl₂(aq) + 2 KBr(aq) → 2 KCl(aq) + Br₂(aq). Orange Br₂ formation confirms the reaction.
A voltaic cell uses a spontaneous redox reaction to produce an electric current. Physically separating the two half-reactions forces electrons to travel through an external wire — that's what we tap for useful current.
Zn(s) | Zn²⁺(aq, 1 M) || Cu²⁺(aq, 1 M) | Cu(s)
Single bar | = phase boundary. Double bar || = salt bridge. Anode on left, cathode on right.
Primary cells are designed to be used once (the redox reaction proceeds in one direction). Secondary cells are rechargeable — applying an external current drives the redox reaction in reverse, regenerating the reactants.
Anode (discharge): Pb(s) + SO₄²⁻(aq) → PbSO₄(s) + 2 e⁻
Cathode (discharge): PbO₂(s) + 4 H⁺(aq) + SO₄²⁻(aq) + 2 e⁻ → PbSO₄(s) + 2 H₂O(l)
Overall (discharge): Pb(s) + PbO₂(s) + 2 H₂SO₄(aq) → 2 PbSO₄(s) + 2 H₂O(l)
Recharge reverses every step. E°cell ≈ 2 V per cell — six cells in series make a 12 V battery.
Lithium ions shuttle between a graphite anode and a metal-oxide cathode (e.g. LiCoO₂). Discharge: Li migrates from anode lattice to cathode. Charge: reversed. Very high energy density and long cycle life — universal in phones, laptops and EVs.
A fuel cell generates electricity by continuously supplying fuel (typically H₂) and oxidant (O₂) — unlike a battery, the reactants are not stored within the cell.
Anode: 2 H₂(g) + 4 OH⁻(aq) → 4 H₂O(l) + 4 e⁻
Cathode: O₂(g) + 2 H₂O(l) + 4 e⁻ → 4 OH⁻(aq)
Overall: 2 H₂(g) + O₂(g) → 2 H₂O(l). E°cell ≈ +1.23 V.
Only product is water — no CO₂, no NOx. Efficiency can exceed 60% (vs ~30% for a combustion engine doing the same chemistry).
In electrolysis, an external voltage drives a non-spontaneous redox reaction. Key uses: extracting reactive metals (Al, Na), purifying metals (Cu), electroplating, manufacturing chlorine and NaOH.
In a voltaic cell, the more reactive metal's terminal is negative. In an electrolytic cell, the battery's positive terminal is connected to the anode (oxidation site) and the negative terminal to the cathode (reduction site).
| Voltaic | Electrolytic | |
|---|---|---|
| Anode polarity | − | + |
| Cathode polarity | + | − |
| Anode = oxidation | ✓ | ✓ |
| Cathode = reduction | ✓ | ✓ |
Cations migrate to the cathode and are reduced. Anions migrate to the anode and are oxidised. Example: molten NaCl → Na(l) at cathode + Cl₂(g) at anode.
A fuel releases energy on combustion or other reaction. Two key metrics:
H₂ has very high specific energy (142 MJ kg⁻¹) but very low energy density at room T (low ρ). Petrol: 47 MJ kg⁻¹, 34 MJ dm⁻³ — much better for vehicles where volume is constrained.
Coal, oil, natural gas. Finite. Produce CO₂ (greenhouse gas), NOx (smog, acid rain), SO₂ (acid rain — coal especially), particulates (lung damage).
Biofuels (bioethanol, biodiesel): can be carbon-neutral if the next crop reabsorbs the CO₂. Drawbacks: land-use conflicts with food.
Hydrogen: carbon-free only if produced from renewable energy (electrolysis of water powered by solar/wind). Steam-reformed H₂ from CH₄ is currently the dominant source — not carbon-neutral.
HL: standard electrode potentials, electrochemical cells under non-standard conditions, electrolysis stoichiometry.
(HL ONLY) Reactivity 3.2.13—Standard cell potential, E⦵cell, can be calculated from standard electrode potentials. E⦵cell has a positive value for a spontaneous reaction.
Lesson 6 of Unit 9.
Lesson 10 of Unit 9.
Lesson 11 of Unit 9.
Any individual half-cell's potential cannot be measured in isolation — only differences between two half-cells. The Standard Hydrogen Electrode (SHE) is the universal reference, assigned E° = 0.00 V by definition.
The half-reaction is: 2 H⁺(aq) + 2 e⁻ ⇌ H₂(g), E° = 0.00 V.
Connect the half-cell to the SHE via a salt bridge and a high-resistance voltmeter. The voltmeter reading is the standard electrode potential of the half-cell of interest. The sign is positive if the half-cell is more strongly reducing (i.e. reduces H⁺), negative if more strongly oxidising.
E° values are tabulated in the IB data booklet — always quoted with the sign and against the SHE.
For any cell, the standard cell potential is:
E°cell = E°(cathode) − E°(anode)
Both values are looked up as reduction potentials in the data booklet. The half-reaction that occurs as reduction is the cathode; the one that runs in reverse (oxidation) is the anode.
If E°cell is positive, the reaction as written is spontaneous.
ΔG° = −n F E°cell
n = moles of electrons transferred per mole of reaction. F = Faraday constant (96 485 C mol⁻¹). A positive E° gives a negative ΔG° — same conclusion, different formalism.
In aqueous solution, water is always present — and it can be electrolysed too. At the cathode, the candidates are the metal cation OR water. At the anode, the candidates are the anion OR water.
The species with the more positive E° (more easily reduced) wins. Water: 2 H₂O + 2 e⁻ → H₂(g) + 2 OH⁻, E° = −0.83 V (at standard alkaline conditions; effectively ~0 V at pH 7).
So Na⁺ (E° = −2.71 V) cannot beat water → H₂ produced. Cu²⁺ (E° = +0.34 V) beats water → Cu deposited.
The species with the more negative E° (more easily oxidised) wins. Water: O₂(g) + 4 H⁺ + 4 e⁻ → 2 H₂O, E° = +1.23 V.
So OH⁻ in dilute solutions usually wins → O₂. But concentrated halide solutions give the halogen instead due to high concentration overriding E° predictions.
When the electrode itself can be oxidised, the rules change. The active anode reaction is the oxidation of the electrode metal rather than water/anion.
Impure Cu anode → Cu²⁺ + 2 e⁻ (anode dissolves). Cathode: Cu²⁺ + 2 e⁻ → pure Cu (deposits). Insoluble impurities (Ag, Au) fall to the bottom as "anode slime"; more reactive metals (Zn, Fe) dissolve but don't redeposit (their E° is too negative). Result: 99.99%+ pure copper.
The object to be plated is made the cathode. The plating metal is the anode. Electrolyte contains a salt of the plating metal. Example: silver-plating a fork — fork is cathode, Ag bar is anode, electrolyte is AgNO₃(aq). Ag dissolves from the anode, then redeposits on the fork.
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