Oxidation-Reduction and Corrosion
Key Takeaways
- Oxidation is loss of electrons (OIL); reduction is gain of electrons (RIG). Remember: OIL RIG.
- In a galvanic cell, the anode (oxidation) is negative and the cathode (reduction) is positive.
- Standard cell potential E°cell = E°cathode - E°anode; a positive value means the reaction is spontaneous.
- Corrosion is electrochemical degradation of metals; iron rust is the classic example (Fe → Fe²⁺ + 2e⁻).
- Corrosion prevention methods include galvanic protection (sacrificial anode), coatings, cathodic protection, and alloying.
- Oxidation states must be assigned to identify what is oxidized and what is reduced in a reaction.
Oxidation-Reduction and Corrosion
Oxidation-Reduction (Redox) Basics
OIL RIG — Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons)
Assigning Oxidation States
- Free elements: oxidation state = 0
- Monatomic ions: oxidation state = ion charge
- Oxygen: usually -2 (except in peroxides: -1)
- Hydrogen: usually +1 (except in metal hydrides: -1)
- Fluorine: always -1
- Sum of oxidation states = charge of the species
Balancing Redox Equations (Half-Reaction Method)
- Separate into oxidation and reduction half-reactions
- Balance atoms other than O and H
- Balance O by adding H₂O
- Balance H by adding H⁺
- Balance charge by adding electrons
- Multiply half-reactions so electrons cancel
- Add half-reactions together
Electrochemistry
Galvanic (Voltaic) Cells
Convert chemical energy to electrical energy (spontaneous).
| Component | Function |
|---|---|
| Anode | Oxidation occurs here (negative terminal) |
| Cathode | Reduction occurs here (positive terminal) |
| Salt Bridge | Maintains electrical neutrality between half-cells |
| External Wire | Electrons flow from anode to cathode |
Standard Cell Potential
A positive E°cell means the reaction is spontaneous (ΔG < 0).
Relationship to Gibbs free energy:
where n = moles of electrons transferred, F = Faraday's constant = 96,485 C/mol
Activity Series (Partial)
More reactive metals (stronger reducing agents) at top:
| Metal | E° (V) |
|---|---|
| Li | -3.04 |
| K | -2.93 |
| Ca | -2.87 |
| Na | -2.71 |
| Mg | -2.37 |
| Al | -1.66 |
| Zn | -0.76 |
| Fe | -0.44 |
| Ni | -0.26 |
| Sn | -0.14 |
| H₂ | 0.00 (reference) |
| Cu | +0.34 |
| Ag | +0.80 |
| Au | +1.50 |
A metal higher in the series can reduce (displace) ions of a metal lower in the series.
Corrosion
Corrosion is the electrochemical degradation of metals through oxidation.
Iron Corrosion (Rusting)
- Anode: Fe → Fe²⁺ + 2e⁻ (iron dissolves)
- Cathode: O₂ + 2H₂O + 4e⁻ → 4OH⁻ (oxygen reduced)
- Overall: Fe²⁺ + 2OH⁻ → Fe(OH)₂ → Fe₂O₃·nH₂O (rust)
Requirements: iron, oxygen, water (electrolyte)
Corrosion Prevention
| Method | Mechanism | Example |
|---|---|---|
| Sacrificial Anode | More reactive metal corrodes instead | Zinc on ship hulls |
| Cathodic Protection | Impressed current makes structure the cathode | Pipeline protection |
| Coatings | Physical barrier blocks electrolyte | Paint, galvanizing |
| Alloying | Adding elements that form protective oxide | Stainless steel (Cr forms Cr₂O₃) |
| Inhibitors | Chemicals slow corrosion reactions | Chromate additives |
| Design | Avoid dissimilar metal contact, allow drainage | Proper joint design |
Galvanic Corrosion
When two dissimilar metals are in contact in an electrolyte, the more reactive (anodic) metal corrodes preferentially. The farther apart the metals are in the galvanic series, the more severe the corrosion.
In the reaction Zn + Cu²⁺ → Zn²⁺ + Cu, which species is oxidized?
Zinc is often used as a sacrificial anode to protect steel structures because:
Given E°(Cu²⁺/Cu) = +0.34 V and E°(Zn²⁺/Zn) = -0.76 V, what is the standard cell potential for a Zn-Cu galvanic cell?