Physical chemistry
Galvanic cells and electrolysis
A galvanic cell converts spontaneous redox energy into electrical work; electrolysis uses an external source to drive a nonspontaneous reaction.
In a Daniell cell, Zn is oxidized at the anode and Cu²⁺ reduced at the cathode. Electrons flow through the external circuit from anode to cathode; the salt bridge maintains charge balance.
Model and quantities
Read the relation together with assumptions about state, experimental conditions, and sign conventions. Keep units consistent and check dimensions before interpreting a result.
Definition: Faraday’s law of electrolysis
Charge Q corresponds to Q/F moles of electrons. The amount formed depends on charge and the number of electrons transferred per particle.
Quantities in the relation are defined for the reaction or system at hand. In particular, distinguish standard-state quantities from actual conditions and do not infer a mechanism from a general expression alone.
Example: Worked example
A 2.0 A current flows for 965 s through Cu²⁺ solution. For Cu²⁺+2e⁻→Cu and F≈96500 C/mol, find the amount of Cu formed.
Solution
Q=It=1930 C; n(e⁻)=1930/96500=0.0200 mol. Two moles of electrons produce one mole Cu, so n(Cu)=0.0100 mol.
| Concept | Description | Unit / note |
|---|---|---|
| Key relation | Use under stated conditions | Check units and sign convention |
| Measured quantity | Relates a state or process | Compare data with model |
| Scope | Model specific conditions | Check assumptions first |
Cell emf is the potential difference at near-zero current; a real load causes a voltage drop from internal resistance. Electrolysis products also depend on overpotential and competing reactions.
In the worked example, which result follows from the given data?
Which statement is consistent with this lesson?
References
- Peter Atkins, Julio de Paula (2014). Physical Chemistry