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R3.2

Electron transfer reactions

10 SL · +5 HL (SL 3.2.1-3.2.11) (HL 3.2.12-3.2.16)

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Parent topic: What are the mechanisms of chemical change?

Guiding question

Guiding questionWhat happens when electrons are transferred?

The one big idea

Redox is electron transfer. Oxidation state tracks who lost and gained electrons; half-equations separate the two halves. Voltaic cells harvest the energy of spontaneous redox; electrolytic cells drive non-spontaneous redox with external energy. At HL, standard electrode potentials quantify it all.

What this rests on

Our editorial judgement.The dependencies in this section are this site’s own assessment of what a sub-topic rests on. They have notbeen verified by a chemistry teacher and are not part of the official IB guide. By contrast, section 8 (Linking questions) reproduces the IB’s own wording verbatim.

What rests on this

Core concepts that must be mastered

  • Oxidation and reduction, half-equations(3.2.1, 3.2.2)Oxidation/reduction can be described four ways: by electron transfer, by oxidation-state change, by oxygen gain/loss, and by hydrogen loss/gain. You must deduce oxidation states in a compound or ion (including variable oxidation states of transition element ions and most main-group non-metals) and identify the oxidised/reduced species and the oxidising/reducing agents. The oxidising agent oxidises something else and is itself reduced; the reducing agent reduces something else and is itself oxidised — the travel-agent analogy (the agent arranges the trip but does not go). Oxidation state is not the same as ionic charge: in SO₄²⁻, S is +6 by bookkeeping, but there is no actual S⁶⁺ ion. Half-equations separate oxidation and reduction. You must deduce them and combine into overall equations in acidic or neutral solutions. Balance atoms first, then oxygen with H₂O, then hydrogen with H⁺, then charge with e⁻. Equalise the electrons before adding the half-equations.
  • Relative ease of oxidation/reduction and electrodes(3.2.3–3.2.5)The relative ease of oxidation/reduction is predictable from periodic position. Metal/metal-ion reactions demonstrate relative ease of oxidation: a more reactive metal displaces a less reactive one. The reactivity series need not be learned — data will be supplied. Down group 17, halogens become weaker oxidising agents while halide ions become stronger reducing agents: F₂ is the strongest oxidiser, I⁻ is the strongest reducer of the halides. Acids react with reactive metals to release hydrogen: Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g). Oxidation occurs at the anode, reduction at the cathode — always, in every cell. The sign differs by cell type: the anode is negative in a voltaic cell (it supplies electrons) and positive in an electrolytic cell (it attracts anions). Anchor on the process (oxidation at the anode), not the sign.
  • Voltaic, secondary and electrolytic cells(3.2.6–3.2.8)A primary (voltaic) cell converts energy from spontaneous redox reactions to electrical energy. Construction includes metal/metal-ion half-cells, anode, cathode, external circuit, and salt bridge. Electrons flow anode → cathode in the external circuit; ions move across the salt bridge to maintain electrical neutrality — electrons do not flow through the salt bridge. Secondary (rechargeable) cells involve reversible redox reactions: you must deduce charging reactions from given discharge electrode reactions and vice versa. You must discuss the advantages and disadvantages of fuel cells, primary cells and secondary cells. An electrolytic cell converts electrical energy to chemical energy, driving non-spontaneous reactions. You must deduce the products of electrolysis of a molten salt — the cation is reduced at the cathode, the anion is oxidised at the anode. For molten NaCl: Na⁺ + e⁻ → Na (cathode); 2Cl⁻ → Cl₂ + 2e⁻ (anode).
  • Oxidation and reduction of organic functional groups(3.2.9–3.2.11)Functional groups may undergo oxidation. You must deduce equations for oxidation of primary and secondary alcohols, including the two-step oxidation of primary alcohols (alcohol → aldehyde → carboxylic acid). Primary alcohol oxidation requires distillation to stop at the aldehyde, or reflux to continue to the carboxylic acid. Tertiary alcohols are not oxidised under similar conditions. Specific oxidising agents and mechanisms are not assessed. Functional groups may undergo reduction: carboxylic acids reduce to primary alcohols via the aldehyde, and ketones reduce to secondary alcohols. The role of hydride ions (H⁻) is included; specific reducing agents and mechanisms are not assessed. Reduction of unsaturated compounds by H₂ addition lowers the degree of unsaturation: alkenes → alkanes, alkynes → alkenes (or alkanes with excess H₂).
  • HL: Standard electrode potentials and Gibbs energy(3.2.12–3.2.14)The hydrogen half-cell H⁺(aq) + e⁻ ⇌ ½H₂(g) is assigned E⦵ = 0 by convention and is used to measure standard electrode potentials. You must interpret standard reduction potential data in terms of ease of oxidation/reduction — a more negative E⦵ means a stronger reducing agent (the metal is easily oxidised); a more positive E⦵ means a stronger oxidising agent (the species is easily reduced). E⦵cell is calculated from standard electrode potentials: E⦵cell = E⦵(cathode) − E⦵(anode), using reduction potentials. Positive E⦵cell ⇒ the reaction is spontaneous forward. Remember: E⦵ is intensive — you do not multiply it by the number of electrons when balancing half-equations. ΔG⦵ = −nFE⦵cell connects electrochemistry to thermodynamics. Here n is the number of moles of electrons transferred and F is the Faraday constant (in the data booklet). This is why ΔG⦵ scales with stoichiometry but E⦵ does not — the n is in this equation, not in the E⦵ calculation. Positive E⦵cell means negative ΔG⦵, which means spontaneous. A positive E⦵cell says nothing about rate — it is a thermodynamic statement, not a kinetic one.

  • HL: Electrolysis of aqueous solutions and electroplating (3.2.15, 3.2.16): In electrolysis of aqueous solutions, competing reactions can occur at each electrode, including oxidation and reduction of water. You must deduce products of aqueous electrolysis from standard electrode potentials. For example, electrolysis of aqueous NaCl: at the cathode, water is reduced in preference to Na⁺ (because reducing water is easier), so hydrogen is evolved; at the anode, Cl⁻ is oxidised in preference to water (because concentrated Cl⁻ tips the competition). Effects of concentration and electrode nature are limited to the two named electrolytes. Electroplating is the electrolytic coating of an object with a thin metallic layer — you must deduce the electrode reactions during electroplating. The object to be plated is the cathode; the plating metal is the anode.

  • How this sub-topic connectsElectron transfer is the second of the three mechanism sub-topics. Oxidation state connects to the periodic table (S3.1: variable oxidation states of transition elements). At HL, E⦵ data connects to thermodynamics via ΔG⦵ = −nFE⦵cell (R1.4) and to equilibrium via ΔG⦵ = −RT lnK (R2.3). Fuel cells (R1.3) are a type of voltaic cell. Organic oxidation and reduction connect to functional group chemistry (S3.2: alcohols, aldehydes, ketones, carboxylic acids). Electroplating and electrolysis connect to industrial chemistry and to the Nature of Science theme. The halogen reactivity trend connects to periodic trends (S3.1).

Quantitative non-negotiables

  • Deduce oxidation states in a compound or ion, including variable oxidation states of transition elements.
  • Deduce redox half-equations and overall equations in acidic or neutral solutions.
  • Identify anode/cathode and their signs in voltaic and electrolytic cells.
  • Deduce products of electrolysis of a molten salt (SL).
  • Deduce equations for oxidation of primary and secondary alcohols, and reduction of carboxylic acids and ketones.
  • At HL: calculate E⦵cell from standard electrode potentials and predict spontaneity.HL
  • At HL: calculate ΔG⦵ = −nFE⦵cell from E⦵ data.HL
  • At HL: deduce products of aqueous electrolysis from standard electrode potentials.HL

Common failure modes

  • M-61 — The oxidising agent is the thing that gets oxidised.

    Confidence: verified

    Why it’s wrong: The oxidising agent oxidises something else and is therefore itself reduced.

    Correction: Use the 'travel agent' analogy: the agent arranges the trip but doesn't go.

  • M-62 — Oxidation state is the same as ionic charge.

    Confidence: verified

    Why it’s wrong: They coincide for monatomic ions but not otherwise.

    Correction: Oxidation state is a formal bookkeeping value assigned by electronegativity even in purely covalent species (S in SO₄²⁻ is +6, though there's no S⁶⁺ ion).

  • M-63 — Fluorine is the best oxidising agent, so fluoride is the best reducing agent.

    Confidence: verified

    Why it’s wrong: Inverted. Down group 17 the halogens become weaker oxidising agents while the halide ions become stronger reducing agents.

    Correction: F₂ is the strongest oxidiser; I⁻ is the strongest reducer of the halides.

  • M-64 — The anode is always positive.

    Confidence: verified

    Why it’s wrong: Oxidation always occurs at the anode, but the anode is negative in a voltaic cell and positive in an electrolytic cell.

    Correction: Anchor on the process (oxidation at the anode), not the sign.

  • M-65 — You multiply E⦵ by the number of electrons when balancing half-equations.

    Confidence: verified

    Why it’s wrong: E⦵ is intensive — it does not scale with stoichiometry.

    Correction: ΔG⦵ does scale; that's why ΔG⦵ = −nFE⦵ has the n in it.

  • M-66 — A positive E⦵cell means the reaction is fast.

    Confidence: verified

    Why it’s wrong: Positive E⦵cell means spontaneous (ΔG⦵ negative).

    Correction: It says nothing about rate.

  • M-67 — Electrons flow through the salt bridge.

    Confidence: verified

    Why it’s wrong: Electrons flow through the external wire.

    Correction: The salt bridge carries ions to maintain electrical neutrality.

  • M-68 — Electrolysis of aqueous NaCl gives sodium at the cathode.

    Confidence: verified

    Why it’s wrong: Water is reduced in preference (hydrogen is evolved), because reducing water is easier than reducing Na⁺.

    Correction: Reactivity 3.2.15 frames this as competing electrode reactions.

What "HL standard" actually looks like

HLHL adds the standard hydrogen electrode (E⦵ = 0 by convention), standard electrode potentials and E⦵cell calculations, spontaneity prediction from E⦵ data, ΔG⦵ = −nFE⦵cell, electrolysis of aqueous solutions with competing electrode reactions (including oxidation/reduction of water), and electroplating. +5 HL hours.

See the command terms reference →

Linking questions

Official IB.The questions in this section are the IB’s own linking questions, reproduced verbatim from the guide. They are not this site’s editorial judgement — see section 3 for that distinction.

  1. Official IB

    Why is the formation of an ionic compound from its elements a redox reaction?

    Structure 2.1.2 → R3.2 Electron transfer reactions

  2. Official IB

    What trends in reactivity of metals can be predicted from the periodic table?

    Structure 2.3.1 → R3.2 Electron transfer reactions

  3. Official IB

    How can oxidation states be used to analyse redox reactions?

    Structure 3.1.6 → R3.2 Electron transfer reactions

  4. Official IB

    How can functional group reactivity be used to determine a reaction pathway between compounds, e.g. converting ethene into ethanoic acid?

    Structure 3.2.2 → R3.2 Electron transfer reactions

  5. Official IB

    Which species are the oxidizing and reducing agents in a combustion reaction?

    Reactivity 1.3.1 → R3.2 Electron transfer reactions

  6. Official IB

    What are some of the environmental, economic, ethical and social implications of burning fossil fuels?

    Reactivity 1.3.3 → R3.2 Electron transfer reactions

  7. Official IB

    What are the main differences between a fuel cell and a primary (voltaic) cell?

    Reactivity 1.3.5 → R3.2 Electron transfer reactions

  8. Official IB

    How can electrochemical data also be used to predict the spontaneity of a reaction?

    Reactivity 1.4.3 → R3.2 Electron transfer reactions

  9. Official IB

    When is it useful to use half-equations?

    Reactivity 2.1.1 → R3.2 Electron transfer reactions

  10. Official IB

    How could we classify the reaction that occurs when hydrogen gas is released from the reaction between an acid and a metal?

    Reactivity 3.1.7 → R3.2 Electron transfer reactions

  11. Official IB

    What are some of the similarities and differences between indicators used in acid–base titrations and in redox titrations?

    Reactivity 3.1.14 → R3.2 Electron transfer reactions

  12. Official IB

    What are the advantages and limitations of using oxidation states to track redox changes?

    Reactivity 3.2.1 → S3.1 The periodic table: Classification of elements

  13. Official IB

    The surface oxidation of metals is often known as corrosion. What are some of the consequences of this process?

    Reactivity 3.2.1 → S2.3 The metallic model

  14. Official IB

    Why does metal reactivity increase, and non-metal reactivity decrease, down the main groups of the periodic table?

    Reactivity 3.2.3 → S3.1 The periodic table: Classification of elements

  15. Official IB

    Electrical energy can be derived from the combustion of fossil fuels or from electrochemical reactions. What are the similarities and differences in these reactions?

    Reactivity 3.2.6 → R1.3 Energy from fuels

  16. Official IB

    Secondary cells rely on electrode reactions that are reversible. What are the common features of these reactions?

    Reactivity 3.2.7 → R2.3 How far? The extent of chemical change

  17. Official IB

    Under what conditions can ionic compounds act as electrolytes?

    Reactivity 3.2.8 → S2.1 The ionic model

  18. Official IB

    How does the nature of the functional group in a molecule affect its physical properties, such as boiling point?

    Reactivity 3.2.9 → S3.2 Functional groups: Classification of organic compounds

  19. Official IB

    What is the difference between combustion and oxidation of an alcohol?

    Reactivity 3.2.9 → R1.3 Energy from fuels

  20. Official IB

    Why is there a colour change when an alcohol is oxidized by a transition element compound?

    Reactivity 3.2.9 → S3.1 The periodic table: Classification of elements

  21. Official IBwording unverified against the guide

    How can oxidation states be used to show that the following molecules are given in increasing order of oxidation: CH₄, CH₃OH, HCHO, HCOOH, CO₂?

    Reactivity 3.2.10 → S3.1 The periodic table: Classification of elements

  22. Official IB

    Why are some reactions of alkenes classified as reduction reactions while others are classified as electrophilic addition reactions?

    Reactivity 3.2.11 → R3.4 Electron-pair sharing reactions

  23. Official IB

    How can thermodynamic data also be used to predict the spontaneity of a reaction?

    Reactivity 3.2.14 → R1.4 Entropy and spontaneity

3 further official linking questions target Tool, Inquiry or Nature of Science strands and are not drawn as edges on the map:

Editorial The decision to surface these off-graph questions here is our editorial judgement — the IB does not prescribe where they should appear.

  1. Official IB

    Why are some redox titrations described as "self-indicating"?

    Reactivity 3.2.2 → Tools 1

  2. Official IB

    What observations can be made when metals are mixed with aqueous metal ions, and solutions of halogens are mixed with aqueous halide ions?

    Reactivity 3.2.3 → Tools 1

  3. Official IB

    How is an electrolytic cell used for electroplating?

    Reactivity 3.2.16 → Tools 1

Bridge: GCSE → IB HL

Derived. These are the bridge items tagged to this sub-topic— places where the GCSE model gets redefined, genuinely new territory, or carry-over strengths. The tagging is this site’s editorial judgement.

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