Oxidation-Reduction and Electrochemistry
Oxidation numbers and balancing
Oxidation is loss of electrons (oxidation number rises); reduction is gain. Rules: elements are 0; O is usually -2 (except in peroxides, -1); H is +1 with nonmetals; a monatomic ion equals its charge; the sum equals the overall charge. The oxidizing agent is itself reduced; the reducing agent is itself oxidized. Balance half-reactions for atoms and charge, adding H2O and H+ in acidic solution, then equalize electrons and add.
Galvanic cells
At the anode oxidation occurs (both start with vowels); at the cathode reduction occurs. Electrons flow through the external wire from anode to cathode; the salt bridge keeps charge neutrality. Standard cell potential: E°cell = E°cathode - E°anode, using reduction potentials. A positive E°cell means the reaction is spontaneous and ΔG° = -nFE°cell is negative. The more positive the reduction potential, the stronger the oxidizing agent.
Electrolysis
An external source drives a nonspontaneous reaction. Faraday's law: n(e-) = It/F with F = 96485 C/mol, and the moles of metal deposited follow the electron ratio from the half-reaction.
Worked example
A current of 2.0 A passes through AgNO3(aq) for 30 min. Charge q = 2.0 x 1800 = 3600 C; n(e-) = 3600/96485 = 0.0373 mol. Since Ag+ + e- -> Ag, moles Ag = 0.0373, and mass = 0.0373 x 107.9 = 4.03 g.
Traps
Reversing anode and cathode in a cell diagram. Subtracting potentials in the wrong order. Forgetting the stoichiometric electron count, e.g. Cu2+ + 2e- -> Cu needs two faradays per mole. Writing a spontaneous cell with a negative E°cell.