Explain why magnesium is a more effective reducing agent than iron, using standard electrode potentials to support your answer.

IB DP Chemistry Higher Level (2023 syllabus) — R3.2 Electron transfer reactions · Explain · 4 marks · View as Markdown

Written & reviewed by James Millett — Biology (Imperial College London), PGCE Science (University of Cambridge).

Standard electrode potentials (E°) at 298 K: Mg²⁺(aq) + 2e⁻ ⇌ Mg(s), E° = −2.37 V; Fe²⁺(aq) + 2e⁻ ⇌ Fe(s), E° = −0.44 V.

Model answer (4 marks)

A reducing agent is a species that donates electrons (is oxidised) in a redox reaction.
A more negative standard electrode potential indicates a greater tendency to be oxidised (lose electrons).
Magnesium has a more negative E° (−2.37 V) than iron (−0.44 V), so magnesium has a greater tendency to be oxidised.
Therefore magnesium donates electrons more readily than iron and is a stronger reducing agent.

Examiner tips

Use the definition of a reducing agent first. Mention that a more negative E° means a stronger tendency to lose electrons. State the actual E° values and compare them. Conclude that magnesium is the stronger reducing agent.

Common mistakes

Confusing a more positive E° with a stronger reducing agent. Failing to mention that the species is oxidised. Not citing the actual E° values in the answer.

Mark scheme (4 marks)

A reducing agent is a species that donates electrons (is oxidised) in a redox reaction.
A more negative standard electrode potential indicates a greater tendency to be oxidised (lose electrons).
Magnesium has a more negative E° (−2.37 V) than iron (−0.44 V), so magnesium has a greater tendency to be oxidised.
Therefore magnesium donates electrons more readily than iron and is a stronger reducing agent.

Key terms in this question

reducing agent · standard electrode potential