Standard electrode potential, E⦵

· Standard electrode potential, E⦵ = the e.m.f. / potential difference of a half-cell compared with the standard hydrogen electrode (SHE) under standard conditions.
· It is always written as a reduction potential: oxidised species + electrons ⇌ reduced species.
· Standard conditions: 298 K, 101 kPa, 1.00 mol dm⁻³ aqueous ions.
· A more positive E⦵ means the species on the left is a stronger oxidising agent and is more easily reduced.
· A more negative E⦵ means the species on the right is a stronger reducing agent and is more easily oxidised.
· E⦵ is not multiplied when a half-equation is multiplied to balance electrons.

Standard hydrogen electrode, SHE

· SHE is the reference half-cell assigned E⦵ = 0.00 V.
· Components: Pt electrode, H₂(g) at 101 kPa, H⁺(aq) at 1.00 mol dm⁻³, 298 K.
· Half-equation: 2H⁺(aq) + 2e⁻ ⇌ H₂(g).
· Platinum is used because it is inert, conducts electricity and provides a surface for electron transfer.
· The SHE is connected to another half-cell using a salt bridge and a high-resistance voltmeter.
· The voltmeter reading gives the standard electrode potential of the other half-cell.

This diagram shows the standard hydrogen electrode used as the reference electrode. It helps students visualise why H₂ gas, H⁺ ions and an inert platinum electrode are required. Source

Measuring electrode potentials

· For a metal / metal ion half-cell: place metal electrode in 1.00 mol dm⁻³ solution of its ions, e.g. Cu(s) | Cu²⁺(aq).
· For a non-metal / ion half-cell: use an inert Pt electrode, e.g. Cl₂(g) | Cl⁻(aq) | Pt.
· For ions of the same element in different oxidation states: use Pt electrode in a solution containing both ions, e.g. Fe³⁺(aq), Fe²⁺(aq) | Pt.
· Use a salt bridge to complete the circuit and allow ion movement.
· Use a high-resistance voltmeter to measure e.m.f. with minimal current flow.
· Cell notation examples: Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s) and Pt | Fe²⁺(aq), Fe³⁺(aq) || Cu²⁺(aq) | Cu(s).

This image shows how two half-cells are connected to measure a cell potential. The salt bridge allows ions to move, completing the circuit without mixing the solutions directly. Source

Standard cell potential, E⦵cell

· Standard cell potential, E⦵cell = potential difference between two standard half-cells.
· Formula: E⦵cell = E⦵cathode – E⦵anode.
· The cathode is where reduction occurs.
· The anode is where oxidation occurs.
· The half-cell with the more positive E⦵ is usually the cathode.
· Electrons flow externally from anode to cathode.
· In cell notation, the left side is usually the anode and the right side is usually the cathode.
· For a spontaneous simple cell: E⦵cell is positive.

Predicting feasibility using E⦵cell

· If E⦵cell > 0, the reaction is feasible / spontaneous under standard conditions.
· If E⦵cell < 0, the reaction is not feasible in the written direction under standard conditions.
· The oxidising agent is the species that is reduced and usually has the more positive E⦵.
· The reducing agent is the species that is oxidised and usually comes from the half-equation with the less positive / more negative E⦵.
· Feasibility predictions from E⦵ values apply to standard conditions only; changes in concentration can change the actual electrode potential.

Constructing redox equations from E⦵ data

· Write both half-equations as reduction equations first.
· Identify the half-equation with the more positive E⦵: this stays as reduction.
· Reverse the other half-equation: this becomes oxidation.
· Balance electrons by multiplying half-equations if needed.
· Add the half-equations and cancel electrons and any repeated species.
· Do not multiply E⦵ values when balancing electrons.
· Check that the final equation gives a positive E⦵cell for the feasible direction.

Relative oxidising and reducing strength

· Strongest oxidising agents have the most positive E⦵ values and are reduced most readily.
· Strongest reducing agents are found on the right-hand side of the most negative E⦵ half-equations and are oxidised most readily.
· Reactivity of metals as reducing agents increases as their E⦵ values become more negative.
· Reactivity of non-metals as oxidising agents increases as their E⦵ values become more positive.
· Use E⦵ data to compare elements, compounds and ions as oxidising or reducing agents.

The Nernst equation

· The Nernst equation predicts how electrode potential changes when conditions are not standard.
· CIE form at 298 K: E = E⦵ + (0.059 / z) log([oxidised species] / [reduced species]).
· z = number of electrons in the half-equation.
· For Cu²⁺(aq) + 2e⁻ ⇌ Cu(s): E = E⦵ + (0.059 / 2) log[Cu²⁺] because pure solids are not included.
· Increasing [oxidised species] makes E more positive.
· Increasing [reduced species] makes E less positive.
· For Fe³⁺(aq) + e⁻ ⇌ Fe²⁺(aq): E = E⦵ + 0.059 log([Fe³⁺] / [Fe²⁺]).
· The Nernst equation explains why changing ion concentration can change cell voltage and even affect feasibility.

ΔG⦵ and E⦵cell

· Equation: ΔG⦵ = –nE⦵cellF.
· ΔG⦵ = standard Gibbs free energy change.
· n = number of moles of electrons transferred in the overall redox equation.
· F = Faraday constant.
· If E⦵cell is positive, then ΔG⦵ is negative, so the reaction is feasible.
· If E⦵cell is negative, then ΔG⦵ is positive, so the reaction is not feasible in the written direction.

Common exam traps

· Do not confuse E⦵ with E⦵cell: one is for a half-cell, the other is for a complete cell.
· Do not reverse signs unless you actually reverse a half-equation.
· Do not multiply E⦵ values when multiplying half-equations.
· Always identify anode = oxidation and cathode = reduction.
· Remember electron flow is from anode to cathode in the external circuit.
· In Nernst calculations, omit pure solids and pure liquids from concentration terms.
· Use the correct z value for the number of electrons transferred in the half-equation.