Electrode potentials and electrochemical cells (A-level only) — AQA A-Level Chemistry Revision

    This topic explores redox reactions occurring in electrochemical cells, where electron transfer between reducing and oxidising agents happens indirectly vi

    Topic Synopsis

    This topic explores redox reactions occurring in electrochemical cells, where electron transfer between reducing and oxidising agents happens indirectly via an external circuit. It covers the measurement of electrode potentials using the standard hydrogen electrode, the use of the electrochemical series to predict reaction feasibility, and the commercial application of these principles in batteries and fuel cells.

    Key Concepts & Core Principles

    Exam Tips & Revision Strategies

    Common Misconceptions & Mistakes to Avoid

    Examiner Marking Points

    Electrode potentials and electrochemical cells (A-level only)

    AQA
    A-Level

    This topic explores redox reactions occurring in electrochemical cells, where electron transfer between reducing and oxidising agents happens indirectly via an external circuit. It covers the measurement of electrode potentials using the standard hydrogen electrode, the use of the electrochemical series to predict reaction feasibility, and the commercial application of these principles in batteries and fuel cells.

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    Objectives
    5
    Exam Tips
    5
    Pitfalls
    0
    Key Terms
    8
    Mark Points

    Topic Overview

    Electrode potentials and electrochemical cells form a core part of AQA A-Level Chemistry, bridging physical and inorganic chemistry. This topic explores how the tendency of a species to gain or lose electrons can be measured and used to predict the direction of redox reactions. You'll learn to construct electrochemical cells, measure standard electrode potentials, and use them to calculate cell potentials (EMF) under standard conditions. Understanding this allows you to predict whether a reaction is feasible and to design cells that generate electricity, such as batteries and fuel cells.

    The topic is crucial for explaining real-world applications like corrosion prevention (sacrificial protection), the operation of rechargeable batteries, and the function of hydrogen-oxygen fuel cells. It also introduces the concept of the electrochemical series, which ranks species by their reducing or oxidising power. Mastery of this area requires a solid grasp of redox reactions, equilibrium, and thermodynamics, as electrode potentials are essentially a measure of thermodynamic tendency. This knowledge is assessed through both calculations and explanations of cell diagrams, including the use of salt bridges and standard conditions.

    Key Concepts

    Core ideas you must understand for this topic

    • →Standard electrode potential (E°) is the potential difference between a half-cell and the standard hydrogen electrode under standard conditions (298 K, 1 mol dm⁻³ solutions, 1 atm pressure for gases).
    • →The standard hydrogen electrode (SHE) is the reference half-cell with a potential of 0.00 V, consisting of H⁺(aq) (1 mol dm⁻³) and H₂(g) (1 atm) at a platinum electrode.
    • →Cell potential (EMF) is calculated as E°cell = E°(cathode) – E°(anode) or E°(right) – E°(left) in cell diagram notation, where the more positive potential is the cathode (reduction).
    • →The electrochemical series lists species in order of their standard electrode potentials; more negative potentials indicate stronger reducing agents, while more positive potentials indicate stronger oxidising agents.
    • →The Nernst equation (E = E° – (RT/nF) ln Q) allows calculation of electrode potentials under non-standard conditions, accounting for concentration, temperature, and pressure.

    What You Need to Demonstrate

    Key skills and knowledge for this topic

    • IUPAC convention for writing half-equations for electrode reactions
    • Conventional representation of cells (cell diagrams)
    • Standard electrode potential (Eθ) definition (298 K, 100 kPa, 1.00 mol dm⁻³)
    • Calculation of cell EMF from standard electrode potentials
    • Prediction of redox reaction direction using Eθ values
    • Simplified electrode reactions in lithium cells
    • Electrode reactions in alkaline hydrogen-oxygen fuel cells
    • Explanation of how electrode reactions generate electric current

    Marking Points

    Key points examiners look for in your answers

    • IUPAC convention for writing half-equations for electrode reactions
    • Conventional representation of cells (cell diagrams)
    • Standard electrode potential (Eθ) definition (298 K, 100 kPa, 1.00 mol dm⁻³)
    • Calculation of cell EMF from standard electrode potentials
    • Prediction of redox reaction direction using Eθ values
    • Simplified electrode reactions in lithium cells
    • Electrode reactions in alkaline hydrogen-oxygen fuel cells
    • Explanation of how electrode reactions generate electric current

    Examiner Tips

    Expert advice for maximising your marks

    • 💡Always remember that the more positive electrode potential is the one that undergoes reduction
    • 💡Ensure cell diagrams are written with the oxidation half-cell on the left and reduction on the right
    • 💡Practice calculating EMF as E(right) - E(left)
    • 💡Be prepared to deduce the overall reaction from given half-equations in commercial cells
    • 💡Remember that fuel cells do not require recharging, unlike secondary cells
    • 💡Always write half-equations with the correct number of electrons and balance them before combining to get the overall cell reaction. Use the rule: electrons must cancel out. For example, for Zn²⁺/Zn and Cu²⁺/Cu, the overall reaction is Zn + Cu²⁺ → Zn²⁺ + Cu.
    • 💡When calculating E°cell, identify which half-cell is the cathode (reduction) and which is the anode (oxidation). The cathode has the more positive (or less negative) E° value. Then use E°cell = E°(cathode) – E°(anode).
    • 💡For questions on feasibility, remember that a reaction with a positive E°cell is thermodynamically feasible under standard conditions, but kinetics may be slow. Also, if the E°cell is very small (e.g., <0.3 V), the reaction may not proceed appreciably due to activation energy.

    Common Mistakes

    Pitfalls to avoid in your exam answers

    • Incorrectly identifying the positive and negative electrodes in a cell diagram
    • Failing to use the correct IUPAC convention for cell notation
    • Confusing the direction of electron flow with the direction of ion flow
    • Incorrectly calculating EMF by subtracting the wrong potential values
    • Neglecting standard conditions when discussing electrode potentials
    • Misconception: The standard hydrogen electrode is always the anode or cathode. Correction: The SHE acts as a reference; its role (anode or cathode) depends on the other half-cell. If the other half-cell has a more positive potential, the SHE is the anode (oxidation occurs).
    • Misconception: A positive cell potential always means the reaction is spontaneous. Correction: While a positive E°cell indicates a spontaneous reaction under standard conditions, non-standard conditions (e.g., concentration changes) can alter the potential. Use the Nernst equation to check feasibility under actual conditions.
    • Misconception: The salt bridge allows electrons to flow between half-cells. Correction: The salt bridge allows ions to flow, maintaining electrical neutrality. Electrons flow through the external circuit, not the salt bridge.

    Frequently Asked Questions

    Common questions students ask about this topic

    Before You Start

    Prior knowledge that will help with this topic

    • •Redox reactions: understanding oxidation states, half-equations, and balancing redox reactions.
    • •Equilibrium: familiarity with Le Chatelier's principle and equilibrium constants, as electrode potentials relate to equilibrium positions.
    • •Thermodynamics: basic concepts of Gibbs free energy (ΔG = –nFE) to connect cell potential to spontaneity.

    Likely Command Words

    How questions on this topic are typically asked

    Calculate
    Predict
    Deduce
    Explain
    Write

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