Thermodynamics and Electrochemistry
Thermodynamics and electrochemistry
Entropy S measures dispersal of energy and the number of microstates. ΔS is positive when gas moles increase, temperature rises, or a solid dissolves into ions. The second law says ΔSuniv = ΔSsys + ΔSsurr > 0 for a spontaneous process, with ΔSsurr = −ΔHsys/T. Gibbs free energy combines both: ΔG = ΔH − TΔS. ΔG < 0 is spontaneous, ΔG = 0 at equilibrium, and ΔG > 0 is nonspontaneous. ΔG° = −RT ln K connects thermodynamics to equilibrium.
Electrochemistry: in a galvanic cell, oxidation occurs at the anode (negative electrode) and reduction at the cathode (positive electrode). E°cell = E°cathode − E°anode, using standard reduction potentials; reversing a half-reaction changes the sign, not the magnitude, and multiplying a half-reaction does not change E°. ΔG° = −nFE°, with n = moles of electrons and F = 96,485 C/mol. The Nernst equation E = E° − (0.0592/n)log Q applies at 25 °C. Electrolytic cells use external voltage to drive nonspontaneous reactions; the anode is still oxidation.
Traps: a positive E°cell means spontaneous under standard conditions, but Q can change the actual voltage. Do not multiply E° when balancing electrons. In electrolysis, charge = current × time, and moles of electrons = It/F.
Worked example: Zn|Zn2+||Cu2+|Cu has E° = 0.34 − (−0.76) = 1.10 V. With n = 2, ΔG° = −2(96,485)(1.10) = −2.12 × 10^5 J = −212 kJ/mol.
Rule of thumb: oxidation at the anode, reduction at the cathode-'an ox, red cat.'