Oxidation-Reduction Reactions and Electron Balance
Redox reactions couple electron loss to electron gain. Oxidation states provide a bookkeeping system for identifying and balancing that transfer.
Learning goals
- Assign oxidation states using standard rules.
- Identify oxidation, reduction, oxidizing agent, and reducing agent.
- Balance electron loss and gain with oxidation-state or half-reaction methods.
- Check conservation of atoms and total charge.
- Recognize disproportionation and comproportionation.
1. Oxidation-state rules
- An element in its standard elemental form has oxidation state 0.
- A monatomic ion has oxidation state equal to its charge.
- Fluorine is −1 in compounds. Oxygen is usually −2; peroxides are a common −1 exception.
- Hydrogen is usually +1 with nonmetals and −1 in many metal hydrides.
- The sum is 0 in a neutral compound and equals the ion charge in a polyatomic ion.
Oxidation state is formal electron bookkeeping; it is not always the measured charge on an atom in a covalent bond.
2. Coupled processes and agents
| Process | Electron change | Oxidation-state change |
|---|---|---|
| Oxidation | loss of electrons | increases |
| Reduction | gain of electrons | decreases |
The reducing agent donates electrons and is oxidized. The oxidizing agent accepts electrons and is reduced. These labels describe the whole reacting species, not the free electron itself.
Zn → Zn²+ + 2e−
Cu²+ + 2e− → Cu
Zn + Cu²+ → Zn²+ + Cu
Cu²+ + 2e− → Cu
Zn + Cu²+ → Zn²+ + Cu
3. Electron-balance workflow
- Assign oxidation states and identify atoms that change.
- Write electron loss and electron gain.
- Multiply changes so electrons lost equal electrons gained.