Electrode potentials and electrochemical cells (A-level only) revision notes

Electrode Potentials & Electrochemical Cells

1. Standard Hydrogen Electrode (SHE)

• The SHE is the reference electrode against which all other electrode potentials are measured.
• It consists of a platinum electrode immersed in a 1 M H⁺ solution, with a hydrogen gas atmosphere at 1 atm.
• The half‑reaction is:

• By definition, the potential of the SHE is 0.00 V under standard conditions.
• The platinum surface is inert and provides a large surface area for the H⁺/H₂ equilibrium.

2. Standard Electrode Potential (E°)

• Definition: The potential difference between a given electrode and the SHE when both are at 1 M concentration, 1 atm pressure, and 25 °C.
• E° values are tabulated for common redox couples.
• Positive E° indicates a tendency to gain electrons (be reduced) relative to SHE; negative E° indicates a tendency to lose electrons (be oxidised).
• E° is independent of concentration, pressure, and temperature – it is a property of the redox couple.

3. Calculating Cell Potentials (E°cell)

• A galvanic cell consists of two half‑cells: an anode (oxidation) and a cathode (reduction).
• The standard cell potential is:

• Rule of thumb: E°cell = E°cathode – E°anode.
• A positive E°cell predicts a spontaneous galvanic reaction; a negative value indicates a non‑spontaneous reaction that would require external energy.
• Example:

• Always write the half‑cell with the higher E° as the cathode.

4. Electrochemical Cell Diagrams

• Cell notation follows the convention:

• Anode is written on the left, cathode on the right.
• The vertical line (|) separates the two half‑cells; the double vertical line (||) indicates the salt bridge or membrane.
• Example of a Zn/Cu cell:

• The direction of electron flow is from left to right (anode to cathode). The ionic flow in the salt bridge balances charge.
• When interpreting a diagram, identify which species are oxidised and which are reduced.

5. Practical Measurement of EMF

• EMF (E) is the measured potential difference between the two electrodes of a galvanic cell.
• Use a high‑impedance voltmeter to minimise current draw.
• Procedure:

1. Assemble the cell with fresh electrodes and a suitable electrolyte.
2. Connect the voltmeter leads to the respective electrodes.
3. Record the voltage reading – this is the EMF.

• Compare the measured EMF with the calculated E°cell to assess cell performance and identify any over‑potential or internal resistance.

6. Predicting Feasibility of Redox Reactions

• Feasibility is judged by the sign of E°cell.
• A positive E°cell indicates that the reaction will proceed spontaneously as written.
• A negative E°cell suggests that the reaction is non‑spontaneous; it could be driven by applying an external voltage.
• The magnitude of E°cell also gives an idea of the driving force – larger values correspond to stronger spontaneous reactions.

7. Limitations of Standard‑Condition Predictions

• Standard potentials are measured under idealised conditions (1 M, 1 atm, 25 °C). Real systems often deviate.
• Concentration effects: Changing ion concentrations alters the actual potential via the Nernst equation.
• Temperature effects: Higher temperatures can shift equilibrium positions.
• Electrode surface area and material: Surface roughness or contamination can introduce over‑potentials.
• Therefore, while E°cell is a useful guide, actual cell potentials may differ.

8. Fuel Cells vs Rechargeable and Non‑Rechargeable Cells

• Fuel cells: Continuous supply of fuel (e.g. hydrogen) and oxidiser (e.g. oxygen) generates electricity; they are typically galvanic and can be considered a type of electrochemical cell.
• Rechargeable cells: Store chemical energy and can be re‑charged by applying a voltage that reverses the redox reactions (electrolytic mode). Example: Li‑ion batteries.
• Non‑rechargeable cells: Designed for one‑time use; once the reactants are consumed, the cell cannot be re‑charged. Example: Alkaline batteries.
• Key differences lie in the ability to reverse the reaction, the design of the electrodes, and the management of ion transport.

9. Effect of Concentration and Conditions on Cell Potentials

• The Nernst equation relates the actual electrode potential to concentration:

• As concentration of the oxidised species increases, the potential of the cathode rises.
• Dilution of the electrolyte generally reduces the cell potential.
• In practice, adjusting concentrations can be used to optimise cell performance or to study kinetic effects.
• Temperature changes affect both the equilibrium constant and the Gibbs free energy, thereby shifting E°cell.

10. Summary of Key Points

• The SHE provides a universal reference (0.00 V).
• Standard electrode potentials are intrinsic properties of redox couples.
• Cell potentials are calculated by subtracting anode E° from cathode E°.
• Cell notation and electron flow direction are essential for diagram interpretation.
• EMF measurement is a practical way to verify theoretical predictions.
• Feasibility is determined by the sign of E°cell, but real‑world conditions can modify the outcome.
• Fuel cells, rechargeable, and non‑rechargeable cells differ mainly in reversibility and fuel supply.
• Concentration and temperature must be considered when applying standard potentials to real systems.

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Further Reading: Refer to the AQA A Level Chemistry specification for tables of standard electrode potentials and detailed experimental procedures for EMF measurement.

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