Section 1: Redox Reactions and Oxidation Numbers

Redox reactions involve simultaneous oxidation (electron loss, oxidation number increase) and reduction (electron gain, oxidation number decrease). Oxidation numbers are assigned using standard rules: free elements = 0; F = −1 (always); O = −2 usually (exceptions: −1 in peroxides $H_{2}O_{2}$/$Na_{2}O_{2}$, +2 in $OF_{2}$); H = +1 usually (exception: −1 in metal hydrides NaH/LiH/$CaH_{2}$); sum = molecular charge. Common examples: Cr is +6 in $K_{2}Cr_{2}O_{7}$; Mn is +7 in $KMnO_{4}$; S is +6 in $SO_{4}^{2-}$ and $H_{2}SO_{4}$; N is +5 in $NO_{3}^{-}$. A species is the reducing agent if it gets oxidized (loses electrons) and the oxidizing agent if it gets reduced (gains electrons).

Section 2: Galvanic Cells and Standard Electrode Potentials

Galvanic cells convert chemical energy to electrical energy spontaneously ($\Delta G$ < 0, E°cell > 0). The anode is the site of oxidation (negative terminal in galvanic cells) and the cathode is the site of reduction (positive terminal). The Daniell cell (Zn|$Zn^{2+}$||$Cu^{2+}$|Cu) has E°cell = 0.34 − (−0.76) = 1.10 V. Standard electrode potentials E° are measured vs. SHE (E° = 0.00 V). The electrochemical series orders metals from strongest reducing agent (Li, E° = −3.04 V) to strongest oxidizing agent ($F_{2}$, E° = +2.87 V). Metal A displaces metal B from its salt solution if E°(A) < E°(B).

Section 3: Nernst Equation and Thermodynamics

Under non-standard conditions: E = E° − (0.0592/n) log Q at 25°C. Q increases as cell discharges, decreasing EMF. At equilibrium: E = 0, Q = K, giving E° = (0.0592/n) log K. Thermodynamic relationships: $\Delta G$° = −nFE° (F = 96500 C/mol). Together: $\Delta G$° = −RT ln K = −nFE°. For Daniell cell: K = 10^37.2 (essentially irreversible under standard conditions).

Section 4: Electrolytic Cells and Faraday's Laws

Electrolytic cells require external power ($\Delta G$ > 0). Anode is positive, cathode is negative in electrolytic cells — opposite polarity from galvanic. Both cell types: oxidation at anode, reduction at cathode. Faraday's first law: w = ZIt where Z = M/(nF). Combined: w = MIt/(nF). Key values: n = 1 for Ag; n = 2 for Cu, Zn, Fe, Pb; n = 3 for Al. Always use t in seconds. 1 Faraday deposits M/n grams = 1 equivalent.

Section 5: Conductance and Kohlrausch's Law

Conductance G = 1/R (S); specific conductance κ = G × cell constant (S/cm); molar conductivity Λm = κ × 1000/M (S·$cm^{2}$/mol). Strong electrolytes: Λm = Λ°m − A√C (linear; extrapolation to C = 0 gives Λ°m). Weak electrolytes: non-linear graph; use Kohlrausch's law instead: Λ°m = ν₊λ°₊ + ν₋λ°₋. Application: Λ°m($CH_{3}COOH$) = Λ°m($CH_{3}COONa$) + Λ°m(HCl) − Λ°m(NaCl) = 390.8 S·$cm^{2}$/mol (typical values). Degree of dissociation: α = Λm/Λ°m.

Section 6: Batteries, Fuel Cells, and Corrosion

Battery types for NEET: (1) Dry cell — 1.5 V, Zn/$MnO_{2}$, not rechargeable. (2) Lead storage — 12 V (6 × 2 V), Pb/$PbO_{2}$, rechargeable; both become $PbSO_{4}$ on discharge; electrolyte ($H_{2}SO_{4}$) dilutes on discharge. (3) Mercury cell — 1.35 V, constant EMF, hearing aids. (4) Nickel-cadmium — 1.2 V, rechargeable, power tools. (5) $H_{2}$-$O_{2}$ fuel cell — continuous feed, ~70% efficiency, water only by-product. Corrosion: electrochemical process; Fe oxidizes at anode, $O_{2}$ reduces at cathode; rust = $Fe_{2}O_{3}$·x$H_{2}$O. Prevention: galvanization (Zn), cathodic protection (Mg), painting, stainless steel (Cr alloy).