Standard electrode potentials E
Cambridge A-Level Chemistry (9701) · Unit 24: Electrochemistry · 11 flashcards
Standard electrode potentials E is topic 24.2 in the Cambridge A-Level Chemistry (9701) syllabus , positioned in Unit 24 — Electrochemistry , alongside Electrolysis. In one line: E⦵ is the potential difference of a half-cell connected to a standard hydrogen electrode (SHE) under standard conditions (298K, 1 atm/100kPa, 1 mol dm⁻³ solution of ions). It represents the tendency of a species to be reduced.
Marked as A2 Level: examined at A Level in Paper 4 (A Level Structured Questions) and Paper 5 (Planning, Analysis and Evaluation). It is not tested on the AS-only papers (Papers 1, 2 and 3).
The deck below contains 11 flashcards — 1 definition, 7 key concepts and 3 calculations — covering the precise wording mark schemes reward. Use the definition card to lock down command-word answers (define, state), then move on to the concept and calculation cards to handle explain, describe, calculate and compare questions.
Standard electrode (reduction) potential, E⦵
E⦵ is the potential difference of a half-cell connected to a standard hydrogen electrode (SHE) under standard conditions (298K, 1 atm/100kPa, 1 mol dm⁻³ solution of ions). It represents the tendency of a species to be reduced.
What the Cambridge 9701 syllabus says
2025-2027 syllabus, in our words · A2 LevelWhat the syllabus expects you to be able to do in this topic, in plain English. For the exact wording, see the official syllabus (PDF).
- Define (a) the standard electrode potential, also known as the standard reduction potential, and (b) the standard cell potential.
- Describe the reference electrode known as the standard hydrogen electrode.
- Describe how standard electrode potentials are measured for: (a) a metal or non-metal in contact with its ions in aqueous solution; (b) ions of one element in two different oxidation states.
- Calculate the standard potential of a cell from the two standard electrode potentials involved.
- Apply standard cell potentials in order to: (a) deduce which electrode is the positive pole and which the negative, and hence explain or deduce the direction of electron movement through a simple cell's external circuit; (b) predict whether a reaction is feasible.
- From E⦵ values, deduce how strong elements, compounds and ions are, relative to one another, as reducing agents or as oxidising agents.
- Combine the relevant half-equations to build redox equations.
- Predict qualitatively how an electrode potential, E, shifts when the aqueous ion concentrations change.
- Use the Nernst equation, for example E = E⦵ + (0.059/z) log([oxidised species]/[reduced species]), to make quantitative predictions of how an electrode potential depends on aqueous ion concentrations, e.g. for Cu²⁺(aq) + 2e⁻ ⇌ Cu(s) and Fe³⁺(aq) + e⁻ ⇌ Fe²⁺(aq).
- Understand the relationship ΔG⦵ = −nFE⦵cell and apply it.
Cambridge syllabus keywords to use in your answers
These are the official Cambridge 9701 terms tagged to this section. Mark schemes credit responses that use the exact term — weave them into your answers verbatim rather than paraphrasing.
Tips to avoid common mistakes in Standard electrode potentials E
- › Standard cell potential is the potential difference between two half-cells under standard conditions; electrode potential refers to a single half-cell against SHE.
- › Check that n is the number of moles of electrons and F is the Faraday constant (96500 C mol–1).
- › Carefully memorize and apply the Nernst equation: E = Eo + (0.059/z) log [oxidised species] / [reduced species].
Define standard electrode (reduction) potential, E⦵.
E⦵ is the potential difference of a half-cell connected to a standard hydrogen electrode (SHE) under standard conditions (298K, 1 atm/100kPa, 1 mol dm⁻³ solution of ions). It represents the tendency of a species to be reduced.
Describe the key features of the standard hydrogen electrode (SHE).
The SHE consists of a platinum electrode immersed in a 1 mol dm⁻³ solution of H⁺ ions, with hydrogen gas bubbled through at 1 atm/100kPa and a temperature of 298K. The platinum electrode acts as a surface for the equilibrium: 2H⁺(aq) + 2e⁻ ⇌ H₂(g). By definition, the E⦵ value is 0.00V.
How can you measure the standard electrode potential of a metal in contact with its ions?
Set up a half-cell with the metal electrode in a 1 mol dm⁻³ solution of its ions under standard conditions. Connect this half-cell to the SHE via a salt bridge. The voltmeter reading gives the standard electrode potential (E⦵) of the metal.
How can you measure the standard electrode potential of ions of the same element in different oxidation states (e.g., Fe²⁺/Fe³⁺)?
Construct a half-cell containing a platinum electrode in a solution containing both ions at 1 mol dm⁻³ concentration each, under standard conditions. Connect this half-cell to the SHE via a salt bridge. The voltmeter reading gives the standard electrode potential (E⦵) for the ion pair.
Calculate the standard cell potential (E⦵cell) for a cell composed of Zn²⁺/Zn (E⦵ = -0.76V) and Cu²⁺/Cu (E⦵ = +0.34V) half-cells.
E⦵cell = E⦵(reduction) - E⦵(oxidation). In this case, copper is reduced and zinc is oxidized. E⦵cell = (+0.34V) - (-0.76V) = +1.10V.
Given E⦵cell = +1.10V for a Zn/Cu cell, deduce the polarity of each electrode and the direction of electron flow.
Since E⦵cell is positive, the reaction is spontaneous. Copper is the positive electrode (cathode, reduction) and zinc is the negative electrode (anode, oxidation). Electrons flow from the zinc electrode to the copper electrode in the external circuit.
How can standard electrode potentials be used to predict the feasibility of a redox reaction?
A reaction is feasible (spontaneous) if the calculated E⦵cell is positive. A positive E⦵cell indicates that the reaction will proceed as written under standard conditions.
How does a more positive E⦵ value relate to the strength of an oxidizing agent?
A more positive E⦵ value indicates a greater tendency for reduction, meaning the species is a stronger oxidizing agent. It has a greater ability to accept electrons.
Write the balanced redox equation for the reaction between Fe²⁺(aq) and MnO₄⁻(aq) in acidic solution, given the half-equations.
First, balance the half equations. Then, multiply each half-equation by a suitable factor so that the number of electrons is the same in both. Add the equations together and cancel the electrons. Result should be: 5Fe²⁺ + MnO₄⁻ + 8H⁺ → 5Fe³⁺ + Mn²⁺ + 4H₂O
State the Nernst equation and explain how it predicts the change of electrode potential with changes in ion concentrations.
The Nernst equation is: E = E⦵ + (0.059/z) log ([oxidised species]/[reduced species]). Increasing the concentration of the oxidised species increases the electrode potential (E) making it more positive and favouring reduction. The reverse is true if you increase the concentration of the reduced species.
State the equation that relates the change in Gibbs Free Energy to the Standard Cell Potential and define each of the terms.
ΔG⦵ = –nFE⦵cell, where ΔG⦵ is the standard Gibbs free energy change, n is the number of moles of electrons transferred in the balanced equation, F is the Faraday constant (96500 C mol⁻¹), and E⦵cell is the standard cell potential.
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Standard electrode potentials E sits alongside these A-Level Chemistry decks in the same syllabus unit. Each uses the same spaced-repetition system, so progress in one informs the next.
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