Redox Processes

Rules: element = 0, O = −2, H = +1, sum = 0 (compound) or charge (ion).

Oxidation: increase in oxidation state / loss of electrons Reduction: decrease in oxidation state / gain of electrons OIL RIG: Oxidation Is Loss, Reduction Is Gain

Half-equation method:

1. Write separate oxidation and reduction half-equations
2. Balance atoms (use H₂O for O, H⁺ for H)
3. Balance charges with electrons
4. Equalise electrons and combine
5. In basic solution: add OH⁻ to neutralise H⁺

Convert chemical energy → electrical energy from spontaneous redox reactions.

• Anode: oxidation occurs (negative terminal)
• Cathode: reduction occurs (positive terminal)
• Salt bridge: completes circuit, maintains electrical neutrality
• Electron flow: anode → external circuit → cathode

$E^\ominus_{cell} = E^\ominus_{cathode} - E^\ominus_{anode}$

• Positive $E^\ominus_{cell}$ → spontaneous
• More positive $E^\ominus$ → stronger oxidising agent
• More negative $E^\ominus$ → stronger reducing agent

Uses electricity to drive non-spontaneous redox reactions.

• Cations → cathode (reduction)
• Anions → anode (oxidation)

For aqueous solutions:

• Cathode: metal deposited if less reactive than H; otherwise H₂
• Anode: halogen if halide present; otherwise O₂

• Redox: electron transfer; OIL RIG
• Voltaic cells: spontaneous redox → electricity; anode = oxidation, cathode = reduction
• $E^\ominus_{cell} = E^\ominus_{cathode} - E^\ominus_{anode}$; positive = spontaneous
• Electrolysis: electricity drives non-spontaneous reactions