Electrochemical Cell

How the model works

Each half-cell has a standard reduction potential E°, its tendency to gain electrons compared with the hydrogen electrode. In a galvanic cell:

• The half-cell with the higher E° is reduced: it is the cathode (+).
• The other is oxidised: it is the anode (−).
• Electrons flow through the wire from anode to cathode. In the salt bridge, anions drift towards the anode and cations towards the cathode to keep both solutions neutral.

E°cell = E°cathode − E°anode
E = E°cell − (RT / nF) · ln Q

The second line is the Nernst equation. Q is the reaction quotient, here the ratio [anode ion]^(n/nₐ) / [cathode ion]^(n/n_c) for the balanced reaction. As the cell discharges, Q grows and the voltage drops; at equilibrium E = 0 and the battery is “flat”.

Standard potentials at 25 °C from common tables. Ideal solutions; no internal resistance or overpotential. The voltage shown is the open-circuit value.

Guided experiments

1. The Daniell cell. Zinc and copper, both 1 M. Prediction: zinc is the anode, copper the cathode, and E° = 0.34 − (−0.76) = 1.10 V. Zn | Cu
2. Swap the electrodes. Put copper on the left and zinc on the right. Prediction: the voltage is still 1.10 V. Anode and cathode are decided by chemistry, not by which side you draw them on. Cu | Zn
3. Dilute the anode. Zn²⁺ = 0.001 M, Cu²⁺ = 1 M. Prediction: E rises to 1.10 + (0.0592/2)·3 ≈ 1.19 V, because a lower product concentration favours the forward reaction. Dilute Zn²⁺

Challenge

A silver–copper cell uses Ag⁺/Ag (E° = +0.80 V) and Cu²⁺/Cu (+0.34 V). Write the overall reaction and give E°cell. Which metal dissolves?

References

• OpenStax, Chemistry 2e, §17.2 “Galvanic Cells”, §17.3 “Electrode and Cell Potentials” and §17.4 “Potential, Free Energy, and Equilibrium”.
• CRC Handbook of Chemistry and Physics, standard reduction potentials table.