24.2 Standard electrode potentials E ⦵ , standard cell potentials E ⦵ cell and the Nernst

A standard electrode (reduction) potential E° is the potential of a half-cell measured relative to the standard hydrogen electrode under standard conditions, written as a reduction.

A standard cell potential is the difference between the cathode and anode reduction potentials: E°cell = E°cathode − E°anode. Do not multiply E° by stoichiometric coefficients.

If E°(Cu²⁺/Cu)=+0.34 V and E°(Zn²⁺/Zn)=−0.76 V, E°cell for Zn|Zn²⁺||Cu²⁺|Cu is +1.10 V.

E° values are not oxidation potentials unless the half-equation is reversed, and their numerical values are not scaled by balancing.

The standard hydrogen electrode uses a platinum surface, H₂ gas at standard pressure, H⁺ solution at standard concentration and a specified temperature. Its reduction potential is defined as 0.00 V.

The platinum is an inert conductor and catalyst; it is not consumed. Connecting another half-cell to the SHE lets its E° be measured directly relative to the zero reference.

If a metal half-cell connected to the SHE gives E°cell=−0.76 V when the metal is the reduction side, its standard reduction potential is −0.76 V.

The SHE is not a container of hydrogen ions only; gas pressure, concentration, temperature and the platinum electrode are part of the standard setup.

To measure E° for a metal/ion pair, connect the half-cell to the SHE through a salt bridge and high-resistance voltmeter. For two oxidation states of one element, use an inert platinum electrode in each solution.

Maintain standard concentrations, gas pressure and temperature, record polarity and write both reduction half-equations. The measured cell voltage gives the difference in reduction potentials.

Fe³⁺/Fe²⁺ uses Pt|Fe³⁺,Fe²⁺, while Zn²⁺/Zn uses Zn metal in Zn²⁺ solution. Each is paired with the SHE to obtain a reference value.

A salt bridge completes ion flow but does not supply electrons, and an inert Pt electrode is required when no solid conductor is present.

For a galvanic cell, E°cell = E°cathode − E°anode, using both values as reduction potentials. The cathode is where reduction occurs and the anode where oxidation occurs.

Balance half-equations separately, but never multiply E° values by stoichiometric coefficients. A positive E°cell indicates a thermodynamically favourable standard-cell direction.

Cu²⁺/Cu = +0.34 V and Zn²⁺/Zn = −0.76 V give E°cell = +1.10 V for Zn oxidation and Cu²⁺ reduction.

Do not add two reduction potentials or reverse a value without also reversing the half-equation direction.

The more positive reduction potential is the cathode; the less positive half-cell is oxidised at the anode. Electrons flow through the external circuit from anode to cathode.

A positive E°cell means the written galvanic direction is feasible under standard conditions. Conventional current is opposite to electron flow, so name the quantity explicitly.

In a Zn/Cu cell, Zn is the negative anode, Cu is the positive cathode, and electrons travel from Zn to Cu through the wire.

The salt bridge carries ions, not electrons, and a positive cathode does not mean positive ions are being reduced in every context.

A species on the left of a reduction half-equation is an oxidising agent because it accepts electrons. More positive E° means a stronger tendency to be reduced and therefore a stronger oxidising agent.

The reduced form is the corresponding reducing agent; a more negative E° for its reduction half-equation generally indicates a stronger reducing form. Compare like half-equations and conditions.

Cl₂/Cl⁻ has a more positive E° than I₂/I⁻, so chlorine is the stronger oxidising agent and iodide is the stronger reducing agent.

Do not rank the products of a half-equation as oxidising agents; identify which side gains electrons.

Write one oxidation and one reduction half-equation, multiply them so electrons cancel, then add and simplify species. Atoms and charge must balance in the final equation.

Use E° values or the stated reaction to choose direction. Do not multiply electrode potentials, only the half-equation coefficients.

Zn → Zn²⁺ + 2e⁻ and Cu²⁺ + 2e⁻ → Cu combine to Zn + Cu²⁺ → Zn²⁺ + Cu.

Do not add electrons to the final equation or cancel species that occur on only one side.

For a reduction half-cell Ox + ze⁻ ⇌ Red, increasing [Ox] generally makes reduction more favourable and increasing [Red] makes it less favourable, at fixed temperature and other standard conditions.

The sign and size of the change depend on the reaction quotient. Qualitative reasoning should identify the species in the numerator and denominator before predicting the direction.

Increasing [Cu²⁺] raises the reduction potential of Cu²⁺/Cu; increasing [Cu(s)] is not meaningful because a pure solid has effectively constant activity.

Do not treat every concentration increase as raising E, and do not include pure solids in the concentration ratio.

At 25 °C, E = E° + (0.059/z) log([oxidised]/[reduced]) for the stated reduction half-equation. z is the number of electrons transferred.

Use activities/concentrations only for aqueous species in the simplified syllabus form, keep the ratio in the correct direction, and use base-10 logarithms.

For Fe³⁺ + e⁻ ⇌ Fe²⁺, increasing [Fe³⁺]/[Fe²⁺] makes E more positive. For Cu²⁺/Cu, the solid copper is omitted from the ratio.

Do not use z from the overall cell equation if the half-equation transfers a different number of electrons, and do not include pure solids.

For an electrochemical cell, ΔG° = −nFE°cell, where n is electrons transferred and F is the Faraday constant. A positive E°cell therefore corresponds to negative ΔG°.

Use consistent units and the cell direction. The relation describes thermodynamic driving force under standard conditions, not reaction speed.

If n=2 and E°cell=1.10 V, ΔG°≈−2×96500×1.10 J mol⁻¹≈−212 kJ mol⁻¹.

Do not omit n or treat a negative E°cell as a spontaneous forward cell reaction.