Redox Reactions

Redox reactions — short for reduction-oxidation — are chemical reactions in which electrons are transferred from one substance to another. One species loses electrons (oxidation), and another gains them (reduction). Both happen simultaneously; you cannot have one without the other. Redox chemistry powers batteries, drives respiration and photosynthesis, causes corrosion, and underlies most of industrial chemistry.

Oxidation and Reduction

The modern definition is purely about electron transfer:

• Oxidation is the loss of electrons.
• Reduction is the gain of electrons.
• OIL RIG is the standard mnemonic: Oxidation Is Loss, Reduction Is Gain.

Example: \( \text{Zn} + \text{Cu}^{2+} \rightarrow \text{Zn}^{2+} + \text{Cu} \). Zinc loses two electrons (oxidized). Copper(II) gains two electrons (reduced).

Oxidizing and Reducing Agents

• Oxidizing agent (oxidant): the species that gains electrons. By gaining them, it causes something else to be oxidized. In the example above, \( \text{Cu}^{2+} \) is the oxidizing agent.
• Reducing agent (reductant): the species that loses electrons. By losing them, it causes something else to be reduced. Zn is the reducing agent.

Common oxidizing agents: \( \text{O}_2 \), \( \text{Cl}_2 \), \( \text{H}_2\text{O}_2 \), \( \text{KMnO}_4 \), \( \text{K}_2\text{Cr}_2\text{O}_7 \). Common reducing agents: H₂, C (carbon, including coke), CO, Na, Mg, Al, NaBH₄.

Oxidation States

To track electron transfer in molecules where ions aren’t obvious, chemists assign oxidation states (oxidation numbers). The rules:

1. Free elements have oxidation state 0 (e.g., O₂, Cu, Fe).
2. Monatomic ions equal their charge (Na⁺ = +1, Cl⁻ = -1).
3. Oxygen is usually -2 (except in peroxides, where it’s -1).
4. Hydrogen is usually +1 (except in metal hydrides, where it’s -1).
5. The sum of oxidation states in a neutral compound equals 0; in an ion, it equals the charge.

Example: in \( \text{KMnO}_4 \), K = +1, four O = -8, so Mn = +7. After reaction in acidic solution, Mn²⁺ has oxidation state +2 — Mn was reduced by 5 units (gained 5 electrons).

Balancing Redox Equations

The half-reaction method splits the equation into two halves — one for oxidation and one for reduction — then balances atoms and charges separately before recombining.

Example: \( \text{MnO}_4^- + \text{Fe}^{2+} \rightarrow \text{Mn}^{2+} + \text{Fe}^{3+} \) in acid.

1. Reduction half: \( \text{MnO}_4^- + 8\text{H}^+ + 5e^- \rightarrow \text{Mn}^{2+} + 4\text{H}_2\text{O} \)
2. Oxidation half: \( \text{Fe}^{2+} \rightarrow \text{Fe}^{3+} + e^- \)
3. Multiply the oxidation half by 5 so electrons cancel.
4. Combined: \( \text{MnO}_4^- + 5\text{Fe}^{2+} + 8\text{H}^+ \rightarrow \text{Mn}^{2+} + 5\text{Fe}^{3+} + 4\text{H}_2\text{O} \).

Real-World Examples

• Rusting. \( 4\text{Fe} + 3\text{O}_2 \rightarrow 2\text{Fe}_2\text{O}_3 \). Iron is oxidized; oxygen is reduced.
• Combustion. \( \text{CH}_4 + 2\text{O}_2 \rightarrow \text{CO}_2 + 2\text{H}_2\text{O} \). Carbon is oxidized from -4 to +4; oxygen is reduced from 0 to -2.
• Cellular respiration. Glucose is oxidized to CO₂; oxygen is reduced to water. The released energy drives ATP synthesis.
• Photosynthesis. The reverse — CO₂ is reduced to glucose; water is oxidized to O₂. Sunlight provides the energy.
• Batteries. Every battery is a controlled redox reaction. A lithium-ion battery oxidizes lithium at the anode and reduces a transition-metal oxide at the cathode while electrons flow through the external circuit.

Related study notes: Electrolysis, Electronegativity, Cellular Respiration, Photosynthesis.

Frequently Asked Questions