Topic 3 of 7
Cell direction and real conditions
Use the potential difference and recognise its limits.
A-Level 9476 (2026-2027)
The spontaneous cell sends electrons from anode to cathode
Use reduction potentials consistently, then balance the chemical equation.
The standard cell potential E°cell is the cell electromotive force, or maximum potential difference measured at negligible current, when all components are in standard states at the stated temperature. For a chosen reaction, E°cell = E°cathode - E°anode, with both tabulated values kept as reduction potentials.
At the cathode, reduction consumes electrons. At the anode, oxidation releases them. In a spontaneous galvanic cell, electrons flow through the wire from the negative anode to the positive cathode. Ions carry charge through the electrolyte and salt bridge; electrons do not flow through the bridge.
A zinc-copper galvanic cell
Zinc is the negative anode and dissolves to zinc ions. Electrons move through the external wire to the positive copper cathode, where copper ions are reduced. Bridge anions enter the zinc side and bridge cations enter the copper side to maintain electrical neutrality.
Worked example
Predict the direction and voltage
Use E°(Zn2+/Zn) = -0.76 V and E°(Cu2+/Cu) = +0.34 V.
- The copper reduction is more favourable, so Cu2+ is reduced at the cathode.
- Reverse the zinc half-equation: Zn → Zn2+ + 2e-. Zinc is the anode and reducing agent.
- E°cell = 0.34 - (-0.76) = +1.10 V.
- The half-equations already transfer two electrons each: Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s).
Electrons flow from zinc to copper. The written reaction is thermodynamically spontaneous under standard conditions. Zinc mass falls and copper mass rises.
| E°cell | Standard thermodynamic conclusion |
|---|---|
| Positive | Forward reaction is favourable: ΔG° < 0. |
| Negative | Forward reaction is unfavourable; the reverse is favoured under standard conditions. |
| Zero | ΔG° = 0; there is no standard thermodynamic driving force in either direction. |
A standard prediction is not a guarantee of a rapid real reaction
Concentration affects equilibrium; activation barriers affect the time taken.
Electrode potentials depend on the actual composition. Think of each half-equation as a reduction equilibrium: a change favouring reduction generally makes that electrode potential more positive. A change favouring the reverse makes it less positive. These are qualitative predictions at fixed temperature, with other conditions held constant.
| Half-equation | Change | Effect on reduction potential |
|---|---|---|
| Cu2+ + 2e- ⇌ Cu | Increase [Cu2+]. | More positive: reduction of Cu2+ is favoured. |
| Fe3+ + e- ⇌ Fe2+ | Increase [Fe3+], or decrease [Fe2+]. | More positive; the oxidised-to-reduced concentration ratio rises. |
| 2H+ + 2e- ⇌ H2 | Decrease [H+] at fixed hydrogen pressure. | Less positive: hydrogen-ion reduction is less favoured. |
In the Zn/Cu cell, diluting Cu2+ lowers the copper reduction potential. Increasing Zn2+ raises the zinc reduction potential. Either change reduces Ecell = ECu - EZn. Far enough from standard conditions, an E° prediction can cease to describe the actual direction. Complex formation and precipitation matter because they change the free ion concentrations.
Even a positive actual cell potential says nothing about how quickly a reaction proceeds. Large activation barriers, slow electrode reactions or a protective surface film can make a thermodynamically favoured change difficult to observe. Different temperature, gas pressure, pH or chemical species can also invalidate a direct application of tabulated E° values. Use the exact relevant half-equation and the conditions, not only the names of the elements.