Oxidation-Reduction and Corrosion

Redox Fundamentals

Redox reactions transfer electrons. Remember OIL RIG — Oxidation Is Loss of electrons, Reduction Is Gain. The species that is oxidized is the reducing agent (it donates electrons); the species reduced is the oxidizing agent.

Assigning Oxidation States

1. Free element (Fe, O₂, H₂): oxidation state = 0
2. Monatomic ion: = the ion charge (Na⁺ is +1, S²⁻ is −2)
3. Oxygen: usually −2 (peroxides −1)
4. Hydrogen: usually +1 (metal hydrides −1)
5. Fluorine: always −1
6. Sum of oxidation states = overall charge of the species

Example: In MnO₄⁻, oxygen contributes 4 × (−2) = −8; total charge is −1, so Mn = +7.

Balancing Redox (Half-Reaction Method)

1. Split into oxidation and reduction half-reactions
2. Balance atoms other than O and H
3. Balance O with H₂O, then H with H⁺ (add OH⁻ to both sides for basic solutions)
4. Balance charge by adding electrons
5. Scale each half so electrons cancel, then add

Verify both mass and charge balance — a common exam trap is leaving a net charge imbalance.

Electrochemistry: Galvanic Cells

A galvanic (voltaic) cell converts chemical energy to electricity through a spontaneous redox reaction.

Standard Cell Potential

$E°_{cell} = E°_{cathode} - E°_{anode}$

A positive E°cell means the reaction is spontaneous (ΔG° < 0). The link to free energy is ΔG° = −nFE°cell, where n = moles of electrons transferred and F = Faraday's constant = 96,485 C/mol.

Worked example — Zn–Cu cell: given E°(Cu²⁺/Cu) = +0.34 V and E°(Zn²⁺/Zn) = −0.76 V. Copper is reduced (cathode), zinc is oxidized (anode): E°cell = 0.34 − (−0.76) = +1.10 V → spontaneous. With n = 2, ΔG° = −(2)(96,485)(1.10) = −212 kJ/mol.

Activity Series (selected E°, V)

A metal with a more negative E° is the stronger reducing agent and will displace ions of any metal below it.

Galvanic vs. Electrolytic Cells

Distinguish the two cell types — a common exam point of confusion:

In both cell types, oxidation always occurs at the anode and reduction at the cathode — only the sign convention flips. Faraday's law governs how much metal is deposited or dissolved: the mass is proportional to the charge passed, m = (Q·M)/(n·F), where Q = I·t (current × time), M = molar mass, n = electrons per ion, and F = 96,485 C/mol. This relation underlies electroplating thickness and corrosion-rate estimates.

Corrosion

Corrosion is the electrochemical degradation of a metal by oxidation. The classic case, the rusting of iron, sets up local galvanic cells on the metal surface:

• Anode: Fe → Fe²⁺ + 2e⁻ (iron dissolves)
• Cathode: O₂ + 2H₂O + 4e⁻ → 4OH⁻
• Overall product: hydrated iron(III) oxide, Fe₂O₃·nH₂O (rust)

Rusting needs iron, oxygen, and water (the electrolyte) all present. Galvanic corrosion accelerates when two dissimilar metals contact in an electrolyte — the more anodic (more negative E°) metal corrodes preferentially, and the farther apart the metals are in the series, the more severe the attack.

Corrosion Prevention

Galvanizing coats steel with zinc; because zinc is more anodic than iron, it both forms a barrier and provides sacrificial protection even where the coating is scratched. This dual action is why zinc is the workhorse of corrosion control.

Forms of Corrosion

Corrosion takes several recognizable forms the FE may name:

Passivation is the key to corrosion-resistant alloys: chromium in stainless steel reacts with oxygen to form a thin, adherent, self-healing Cr₂O₃ film that blocks further attack — which is why stainless requires roughly 11% or more chromium. Corrosion rate is often expressed in mils per year (mpy) and scales with the corrosion current density via Faraday's law, so anything that raises the cell's current speeds metal loss.