Central Topic: Electrochemistry

Branch 1: Redox Fundamentals

• Oxidation = loss of $e^{-}$, increase in oxidation number
• Reduction = gain of $e^{-}$, decrease in oxidation number
• OX number rules → F always −1; O usually −2 (exceptions: peroxides −1, $OF_{2}$ +2); H usually +1 (exception: metal hydrides −1)
• Balancing → half-reaction method (balance atoms then charge)

Branch 2: Galvanic Cells

• Structure: anode | anode solution || cathode solution | cathode
• Spontaneous: $\Delta G$ < 0, E°cell > 0
• Anode: oxidation, negative terminal
• Cathode: reduction, positive terminal
• Salt bridge: maintains electrical neutrality
• Sub-branch: Daniell Cell (Zn/Cu, E° = 1.10 V)
• Sub-branch: Standard Electrode Potential + Electrochemical Series

Branch 3: Quantitative Electrochemistry

• E°cell = E°cathode − E°anode
• Nernst: E = E° − (0.0592/n) log Q
• Equilibrium: E = 0, E° = (0.0592/n) log K
• $\Delta G$° = −nFE°
• Faraday's law: w = MIt/(nF); 1F = 96500 C

Branch 4: Electrolytic Cells

• Non-spontaneous: requires external power
• Anode: positive, cathode: negative (opposite of galvanic)
• Oxidation at anode, reduction at cathode (same as galvanic)
• Applications: electroplating, electrolysis of water/NaCl, refining metals

Branch 5: Conductance

• G = 1/R (S); specific conductance κ (S/cm); molar conductivity Λm (S·$cm^{2}$/mol)
• Strong electrolytes: Λm = Λ°m − A√C (slight increase on dilution)
• Weak electrolytes: sharp increase on dilution; α = Λm/Λ°m
• Kohlrausch's law: Λ°m = ν₊λ°₊ + ν₋λ°₋

Branch 6: Batteries and Fuel Cells

• Dry cell: Zn/$MnO_{2}$, 1.5 V, non-rechargeable
• Lead storage: Pb/$PbO_{2}$, 2 V/cell × 6 = 12 V, rechargeable
• Mercury cell: Zn-Hg/HgO, 1.35 V, constant voltage
• $H_{2}$-$O_{2}$ fuel cell: ~70% efficiency, water by-product

Branch 7: Corrosion

• Electrochemical process: Fe oxidizes at anode, $O_{2}$ reduces at cathode
• Rust: $Fe_{2}O_{3}$·x$H_{2}$O
• Prevention: galvanization (Zn coating), cathodic protection (Mg anode), painting, alloying (stainless steel)