IB DP Chemistry Reactivity 3.2: Electron transfer: redox, cells and electrolysis – Revision Notes

These notes condense Reactivity 3.2, Electron transfer reactions, with the February 2023 IB Diploma Programme Chemistry guide (first assessment 2025) as the checklist. They cover understandings Reactivity 3.2.1 to 3.2.11, which are common to SL and HL; the additional higher level material from 3.2.12 onwards is in the HL unit on standard electrode potentials. The study guide has the longer explanations and worked examples.

When you have read these notes, try the practice questions. Follow your progress with the printable checklist and the IB DP chemistry hub, and try our free diagnostics to find weak spots.

Redox definitions (3.2.1)

Process	Electrons	Oxidation state	Oxygen	Hydrogen
Oxidation	lost	rises	gained	lost
Reduction	gained	falls	lost	gained

• Oxidizing agent: accepts electrons, so it is reduced.
• Reducing agent: donates electrons, so it is oxidized.
• Oxidation states are written sign first: +2, −1. “Oxidation number” and “oxidation state” are used interchangeably.

Oxidation state rules

1. Atoms in an element on its own (Na, O₂, S₈) are 0.
2. A monatomic ion takes the value of its charge.
3. Oxygen −2, except −1 in peroxides.
4. Hydrogen +1, except −1 in metal hydrides.
5. All values total 0 for a neutral compound; for an ion they total its overall charge.

Small reminder: S in HSO₃⁻ → +1 + x + 3(−2) = −1, so x = +4.

Method in steps: finding the agents

1. Write the oxidation state of every atom that could change, on both sides.
2. The atom whose oxidation state rises belongs to the species that is oxidized; that species is the reducing agent.
3. The atom whose oxidation state falls belongs to the species that is reduced; that species is the oxidizing agent.
4. Name whole species (Cu²⁺, not “copper”), and check that the total rise equals the total fall.

Small reminder: in Cu²⁺ + Fe → Cu + Fe²⁺, iron rises 0 → +2 (oxidized, so Fe is the reducing agent) and copper falls +2 → 0 (reduced, so Cu²⁺ is the oxidizing agent).

Variable oxidation states and naming

• Transition elements: iron +2 and +3; copper +1 and +2; chromium +3 and +6; manganese +2, +4, +6, +7.
• Main group non-metals: sulfur −2 (H₂S), +4 (SO₂), +6 (SO₄²⁻); nitrogen −3 (NH₃) up to +5 (NO₃⁻); chlorine −1 up to +7.
• Stock names use Roman numerals: CuCl is copper(I) chloride, Cr₂O₃ is chromium(III) oxide, NO₃⁻ can be nitrate(V). Common names such as nitrate and sulfite are still accepted.
• Limitation: oxidation states assume ionic bonding, so they are a counting device, not real charges.

Half-equations (3.2.2)

Method in steps (acidic solution)

1. Balance the element changing oxidation state.
2. Make oxygen atoms equal using water.
3. Make hydrogen atoms equal using H⁺.
4. Make charges equal using electrons, placed on the more positive side.
5. To combine: scale to equal electrons, add, cancel species on both sides.

Small reminder: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O. Charge check: −1 + 8 − 5 = +2 on the left; +2 on the right.

Neutral solution: simple metal/ion pairs need no H⁺ or H₂O, for example 2Al + 3Cu²⁺ → 2Al³⁺ + 3Cu.

Self-indicating titrations: a coloured reagent that turns colourless when reduced (purple manganate(VII)) shows its own end point.

Group trends and displacement (3.2.3)

Group	Process	Trend down the group	Reason
1 and 2 metals	oxidation (lose e⁻)	easier	outer electrons further out, more shielded
17 halogens	reduction (gain e⁻)	harder	incoming electron further out, more shielded

• Halogen displacement: Cl₂ oxidizes Br⁻ and I⁻; Br₂ oxidizes I⁻ only; I₂ oxidizes neither.
• Colours in water: bromine yellow-orange, iodine brown.
• Metal/metal ion data are supplied in questions. The metal that reduces the most other ions is oxidized most easily.

Metals with dilute acids (3.2.4)

• General: metal + acid → salt + hydrogen.
• Ionic: M + 2H⁺ → M²⁺ + H₂ (for a 2+ ion); 2M + 6H⁺ → 2M³⁺ + 3H₂ (for a 3+ ion).
• Examples: Zn + 2HCl → ZnCl₂ + H₂; Mg + H₂SO₄ → MgSO₄ + H₂.
• The metal is the reducing agent; H⁺ is the oxidizing agent.

Electrodes and cells (3.2.5 to 3.2.8)

Electrode signs

	Voltaic (primary, or secondary discharging)	Electrolytic (or secondary charging)
Anode	oxidation, negative	oxidation, positive
Cathode	reduction, positive	reduction, negative
Energy change	chemical → electrical	electrical → chemical
Reaction	spontaneous	non-spontaneous

Memory aid: an ox, red cat (anode oxidation, reduction cathode) holds in every cell.

Primary (voltaic) cell parts

Two metal/metal ion half-cells, an anode and a cathode, an external circuit and a salt bridge.

• Electrons: anode → cathode, through the wire only.
• Salt bridge: anions toward the anode half-cell, cations toward the cathode half-cell. It keeps each solution electrically neutral and completes the circuit.
• Choose bridge ions that will not precipitate with ions in either half-cell.

Secondary cells

• To get a charging half-equation, reverse the discharge half-equation; to get discharge from given charging reactions, reverse those. The terminal signs stay the same; the names anode and cathode swap.
• Examples: lead–acid (Pb and PbO₂ in sulfuric acid), nickel–cadmium, lithium-ion.

	Primary	Secondary	Fuel cell
Plus	cheap, light, long shelf life	rechargeable, less waste	continuous while fuelled; efficient; H₂ gives water only
Minus	single use, disposal	higher initial cost; capacity fades; some toxic metals	costly catalysts; H₂ storage; H₂ source

Electrolytic cell

• Parts: DC power supply, anode, cathode, electrolyte.
• Conduction: electrons in the wires; mobile ions in the electrolyte; electron transfer at electrode surfaces.
• Molten binary salt: metal at the cathode, non-metal at the anode. Example, molten NaBr: Na⁺ + e⁻ → Na; 2Br⁻ → Br₂ + 2e⁻.

Organic redox (3.2.9 to 3.2.11)

Starting compound	Oxidation [O]	Reduction [H]
Primary alcohol	aldehyde, then carboxylic acid	–
Secondary alcohol	ketone	–
Tertiary alcohol	no reaction (same conditions)	–
Carboxylic acid	–	aldehyde, then primary alcohol
Ketone	–	secondary alcohol
Alkene / alkyne	–	+ H₂ (Ni catalyst): alkane / alkene, then alkane

Equation patterns

• RCH₂OH + [O] → RCHO + H₂O
• RCHO + [O] → RCOOH
• R₂CHOH + [O] → R₂CO + H₂O
• RCOOH + 2[H] → RCHO + H₂O
• RCHO + 2[H] → RCH₂OH
• R₂CO + 2[H] → R₂CHOH

Distillation vs reflux

• Aldehyde wanted: distil it off as it forms (lower boiling point, no hydrogen bonding between its molecules).
• Carboxylic acid wanted: reflux with excess oxidizing agent, so nothing escapes before full oxidation.

Hydride ions: reducing agents for carbonyl compounds act as sources of H⁻, which adds to the δ+ carbonyl carbon. Reagent names and mechanisms are not assessed.

Hydrogenation: each mole of C=C uses 1 mol H₂; each mole of C≡C uses 2 mol H₂ to reach the alkane. Adding hydrogen is reduction and lowers the degree of unsaturation.

Must-know distinctions

• Oxidized species vs oxidizing agent: opposite roles.
• Voltaic vs electrolytic anode: both oxidation; signs opposite.
• Primary vs secondary cell: irreversible vs reversible electrode reactions.
• Fuel cell vs primary cell: reactants supplied continuously vs stored inside.
• Distillation vs reflux: remove the product as it forms vs keep everything in the flask.
• Hydrogenation vs bromine addition to an alkene: reduction vs electrophilic addition.

Quick self-test

1. State the oxidation state of vanadium in VO₂⁺ and of manganese in K₂MnO₄.
2. Name Cu₂O and CrCl₃ using oxidation numbers.
3. In 2Fe³⁺ + 2I⁻ → 2Fe²⁺ + I₂, identify the oxidizing agent.
4. Write the half-equation for SO₄²⁻ changing to SO₂ in acidic solution.
5. Which reacts with aqueous potassium bromide: chlorine or iodine? Give the ionic equation.
6. Write the ionic equation for zinc with dilute hydrochloric acid.
7. A cell is made from manganese and copper half-cells. Manganese is oxidized more easily. Which electrode is negative, and which way do electrons flow in the wire?
8. On discharge, one electrode of a nickel–cadmium cell reacts as Cd + 2OH⁻ → Cd(OH)₂ + 2e⁻. Write the reaction at this electrode during charging.
9. Molten magnesium chloride is electrolysed. Give the product at each electrode.
10. Name the organic product when propan-2-ol is oxidized, and when 2-methylpropan-1-ol is heated under reflux with excess oxidizing agent.
11. Name the product of reducing butanal, and of reducing cyclohexanone.
12. How many moles of hydrogen convert 0.25 mol of propyne into propane?

Answers

1. V: x + 2(−2) = +1, so +5. Mn: 2(+1) + x + 4(−2) = 0, so +6.
2. Copper(I) oxide; chromium(III) chloride.
3. Fe³⁺ (it is reduced from +3 to +2).
4. SO₄²⁻ + 4H⁺ + 2e⁻ → SO₂ + 2H₂O (charge 0 on each side).
5. Chlorine: Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂.
6. Zn + 2H⁺ → Zn²⁺ + H₂
7. Manganese is the negative anode; electrons flow from manganese to copper.
8. Cd(OH)₂ + 2e⁻ → Cd + 2OH⁻ (now reduction).
9. Magnesium at the cathode; chlorine at the anode.
10. Propanone; 2-methylpropanoic acid.
11. Butan-1-ol; cyclohexanol.
12. 2 × 0.25 = 0.50 mol.

Where marks are usually lost

• Quoting a total for two atoms (for example +12 for Cr₂) instead of the value per atom.
• Naming the reduced species as the reducing agent.
• Adding electrons to the wrong side of a half-equation, or leaving them in an overall equation.
• Balancing atoms but not charge.
• Giving the anode in an electrolytic cell a negative sign.
• Saying electrons, rather than ions, travel through the salt bridge or electrolyte.
• Writing a charging reaction that is not the exact reverse of the discharge reaction.
• Stating that a tertiary alcohol forms a ketone.
• Using distillation when the question asks for the carboxylic acid.
• Forgetting that an alkyne needs two moles of H₂ per mole to reach the alkane.

Official syllabus

Reactivity 3.2 (Electron transfer reactions), understandings 3.2.1 to 3.2.11 for SL and HL, summarized from the Chemistry guide of the International Baccalaureate Diploma Programme, February 2023, first assessment 2025.