Electrochemistry is one of the highest-yield topics in NEET Chemistry, bridging physical chemistry's thermodynamics with the practical world of batteries, corrosion, and industrial electrolysis. The subject centers on electron transfer reactions — how they generate electricity spontaneously or are driven by external power.

Redox Foundations

Every electrochemical process is a redox reaction. Oxidation is the loss of electrons and corresponds to an increase in oxidation number; reduction is the gain of electrons corresponding to a decrease in oxidation number. Assigning oxidation numbers correctly is a prerequisite skill. The universal rules: free elements have oxidation number zero; F is always −1 (most electronegative element, no exception); O is usually −2 but is −1 in peroxides ($H_{2}O_{2}$, $Na_{2}O_{2}$) and +2 in $OF_{2}$; H is +1 except in metal hydrides (NaH, LiH, $CaH_{2}$) where it is −1; and the sum of all oxidation numbers must equal the charge on the species (zero for neutral molecules, non-zero for ions). These rules immediately allow identification of which species is oxidized and which is reduced in any reaction.

Galvanic Cells

Galvanic (voltaic) cells are spontaneous electrochemical cells where chemical energy is directly converted to electrical energy. The defining thermodynamic criterion is $\Delta G$ < 0 and E°cell > 0. The structural components are: an anode (where oxidation occurs) and a cathode (where reduction occurs), each immersed in their respective solutions, connected by an external circuit for electron flow and a salt bridge (KCl in agar) for ion flow to maintain electrical neutrality. The anode is the negative terminal (electrons flow away), and the cathode is the positive terminal (electrons arrive). In cell notation, the anode is written on the left: Zn(s)|$Zn^{2+}$(aq)||$Cu^{2+}$(aq)|Cu(s) represents the Daniell cell.

The standard cell EMF is calculated as E°cell = E°cathode − E°anode. For the Daniell cell: E°cell = (+0.34) − (−0.76) = +1.10 V. This formula always uses cathode minus anode. Standard electrode potentials E° are measured against the standard hydrogen electrode (SHE, E° = 0.00 V). The electrochemical series arranges half-reactions from most negative E° (strongest reducing agents: Li, K, Na, Al) to most positive E° (strongest oxidizing agents: Cu, Ag, Au, $F_{2}$). Any metal higher in the series (more negative E°) will displace any metal lower in the series from its salt solution.

Nernst Equation and Equilibrium

Real cells operate under non-standard conditions. The Nernst equation accounts for concentration effects: E = E° − (0.0592/n) log Q at 25°C, where n is the number of electrons transferred and Q is the reaction quotient. As a cell discharges, products accumulate and reactants deplete, increasing Q and decreasing EMF. At equilibrium, Q = K and E = 0, giving the crucial relationship: E° = (0.0592/n) log K. The Gibbs energy relationship $\Delta G$° = −nFE° (with F = 96500 C/mol) unifies thermodynamics and electrochemistry: spontaneous cells (E°cell > 0) have $\Delta G$° < 0 and K > 1.

Electrolytic Cells

Electrolytic cells are non-spontaneous ($\Delta G$ > 0, E°cell < 0) and require external electrical power to drive the chemical reaction. In contrast to galvanic cells, the anode is the positive terminal (connected to the external battery's positive terminal) and the cathode is the negative terminal. However, the fundamental reactions remain: oxidation at anode and reduction at cathode — the rule "AN OX RED CAT" is universal. Electrolysis of molten NaCl deposits Na at cathode ($Na^{+}$ + $e^{-}$ → Na) and produces $Cl_{2}$ at anode; in aqueous NaCl, $H_{2}$ forms at cathode preferentially (water is reduced rather than $Na^{+}$ due to overvoltage effects and more favorable reduction potential).

Faraday's Laws of Electrolysis

Michael Faraday established the quantitative laws governing electrolysis. The key formula is w = MIt/(nF), where w is mass deposited (g), M is molar mass (g/mol), I is current (A), t is time in seconds (critical — not minutes), n is electrons transferred per ion, and F = 96500 C/mol. For copper ($Cu^{2+}$ + 2$e^{-}$ → Cu, n = 2): 2 A for 1 hour deposits (63.5 × 2 × 3600)/(2 × 96500) = 2.37 g. 1 Faraday = 96500 C deposits 1 equivalent (M/n grams) of any substance.

Conductance

Conductance G = 1/R (unit: Siemens, S). Specific conductance κ = G × (l/A) where l/A is the cell constant. Molar conductivity Λm = κ × 1000/M (units: S·$cm^{2}$/mol) is the conductance of 1 mol of electrolyte. For strong electrolytes, Λm increases slightly on dilution following Λm = Λ°m − A√C (Debye-Hückel-Onsager). For weak electrolytes, Λm rises sharply on dilution as ionization increases; the degree of dissociation α = Λm/Λ°m. Kohlrausch's law states that limiting molar conductivity equals the sum of individual ionic conductivities: Λ°m = ν₊λ°₊ + ν₋λ°₋. This allows indirect determination of Λ°m for weak electrolytes that cannot be obtained by graph extrapolation.

Batteries, Fuel Cells, and Corrosion

Key batteries tested in NEET: dry cell (Leclanché, 1.5 V, non-rechargeable), lead storage battery (12 V total from six 2 V cells, rechargeable — both electrodes become $PbSO_{4}$ on discharge), mercury cell (1.35 V, constant EMF), and hydrogen-oxygen fuel cell (~70% efficiency with water as only product, used in space applications). Corrosion is an electrochemical process: iron acts as anode at anodic spots (Fe → $Fe^{2+}$ + 2$e^{-}$) and $O_{2}$ is reduced at cathodic spots ($O_{2}$ + 2$H_{2}$O + 4$e^{-}$ → 4$OH^{-}$), forming rust ($Fe_{2}O_{3}$·x$H_{2}$O). Prevention includes galvanization (Zn coating acts as sacrificial anode since E°($Zn^{2+}$/Zn) = −0.76 V < E°($Fe^{2+}$/Fe) = −0.44 V), cathodic protection, painting, and alloying.