Oxidation is loss of electrons (oxidation number increases) and reduction is gain of electrons (oxidation number decreases); both must occur together in a redox reaction. Galvanic cells exploit spontaneous redox reactions to generate electrical current, with oxidation at the negative anode and reduction at the positive cathode. The standard cell EMF is calculated as E°cell = E°cathode − E°anode, always "cathode minus anode." The Nernst equation (E = E° − (0.0592/n) log Q at 25°C) adjusts EMF for non-standard concentrations, and at equilibrium E = 0 giving the relationship E° = (0.0592/n) log K. The Gibbs energy relationship $\Delta G$° = −nFE° connects electrochemistry directly to thermodynamics, with F = 96500 C/mol. Electrolytic cells are non-spontaneous — they require external power — and have reversed polarity (anode is positive, cathode is negative), though oxidation and reduction still occur at anode and cathode respectively. Faraday's law quantifies electrolysis: w = MIt/(nF), with time strictly in seconds and n = electrons per ion ($Ag^{+}$: 1; $Cu^{2+}$: 2; $Al^{3+}$: 3). Molar conductivity (Λm = κ × 1000/M) increases on dilution for all electrolytes, sharply for weak ones; degree of dissociation α = Λm/Λ°m and Kohlrausch's law (Λ°m = ν₊λ°₊ + ν₋λ°₋) are used to find limiting values. Corrosion is an electrochemical process producing rust ($Fe_{2}O_{3}$·x$H_{2}O$) that is prevented by galvanization, cathodic protection (Mg blocks), or painting. The lead storage battery (12 V, rechargeable) and the $H_{2}$-$O_{2}$ fuel cell (~70% efficiency) are the key devices tested in NEET alongside the dry cell (1.5 V, non-rechargeable) and mercury cell (1.35 V, constant EMF).