Redox Reactions & Electrochemistry: Core Rules and Key Points

Oxidation = loss of electrons = increase in oxidation number; reduction = gain of electrons = decrease in oxidation number.
Oxidation number rules: Free element = 0; F = −1 (always, no exception); O = −2 (except −1 in peroxides $H_{2}O_{2}$ , +2 in $OF_{2}$ ); H = +1 (except −1 in metal hydrides); sum = charge on species.
Galvanic cell: Spontaneous ( $\Delta G$ < 0, E°cell > 0). Chemical → Electrical energy. Anode is NEGATIVE terminal; cathode is POSITIVE terminal.
Electrolytic cell: Non-spontaneous ( $\Delta G$ > 0). Electrical → Chemical energy. Anode is POSITIVE terminal; cathode is NEGATIVE terminal.
Universal rule (both cell types): Oxidation at anode; reduction at cathode. ("AN OX RED CAT")
E°cell formula: E°cell = E°cathode − E°anode (ALWAYS cathode minus anode).
Standard electrode potentials: Measured vs. SHE (E° = 0.00 V). More negative E° = stronger reducing agent; more positive E° = stronger oxidizing agent.
Key E° values: Li = −3.04 V (strongest reducing agent), Zn = −0.76 V, Fe = −0.44 V, $H_{2}$ = 0.00 V, Cu = +0.34 V, Ag = +0.80 V, $F_{2}$ = +2.87 V (strongest oxidizing agent).
Daniell cell: Zn| $Zn^{2+}$ || $Cu^{2+}$ |Cu; E°cell = 0.34 − (−0.76) = 1.10 V; n = 2.
Nernst equation (25°C): E = E° − (0.0592/n) log Q. As cell discharges, Q↑, EMF↓.
At equilibrium: E = 0, Q = K; therefore E° = (0.0592/n) log K.
Gibbs energy: $\Delta G$ ° = −nFE°; F = 96500 C/mol; positive E°cell ↔ negative $\Delta G$ ° ↔ spontaneous.
Faraday's law: w = MIt/(nF); t MUST be in seconds; n is electrons per ion ( $Ag^{+}$ : n=1; $Cu^{2+}$ : n=2; $Al^{3+}$ : n=3).
Conductance: G = 1/R (S); κ = G × cell constant (S/cm); Λm = κ × 1000/M (S· $cm^{2}$ /mol).
Kohlrausch's law: Λ°m = ν₊λ°₊ + ν₋λ°₋. For $CH_{3}COOH$ : Λ°m = Λ°m( $CH_{3}COONa$ ) + Λ°m(HCl) − Λ°m(NaCl).
Weak electrolytes: α = Λm/Λ°m; Λm rises sharply on dilution; Λ°m cannot be obtained by graph extrapolation.
Strong electrolytes: Λm = Λ°m − A√C; slight increase on dilution; Λ°m obtainable by extrapolation.
Corrosion: Fe oxidizes at anode (Fe → $Fe^{2+}$ + 2 $e^{-}$ ); $O_{2}$ reduced at cathode ( $O_{2}$ + 2 $H_{2}$ O + 4 $e^{-}$ → 4 $OH^{-}$ ); rust = $Fe_{2}O_{3}$ ·x $H_{2}$ O. Prevented by galvanization (Zn sacrificial anode), cathodic protection (Mg blocks), painting.
Lead storage battery: 12 V (6 × 2 V cells); rechargeable; both electrodes become $PbSO_{4}$ on discharge; $H_{2}$$SO_{4}$ dilutes on discharge.