University College - Chemistry Chem 112A Lecture Notes - Lecture 11: Nernst Equation, Reaction Rate, Rate Equation

• School: WUSTL
• Department: University College - Chemistry
• Course: University College - Chemistry Chem 112A
• Professor: Mabbs Richard

Document Summary

Nernst equation: n_red = e^o_red - rt/(nf) ln(q) q for reaction written as reduction. Process spontaneous when deltae_cell is positive (ie oxidation spontaneous at anode). Anode: 6h2o -> o2 + 4h3o+ + 4e- [1/2o2 + 2e- + 2h3o+ -> 3h2o e^o_red = 1. 229v] Cathode: 2 electrons + 2h2o -> h2 + 2oh- e^o = -. 8281. Anode: e_red = e^o_red - rt/(nf)ln(a_h2o^2)/(a_o2^( ) * a_h3o+^2)) E- don"t spontaneously flow from anode to cathode. Apply an external potential to drive the reaction. Electrolytic cell: e- flow from anode to cathode because of applied potential. Cathode reactions (ph , 298k, 1 atm gas, [nicl2]=. 1m. 2h2o + 2e- -> 2 oh- + h2 e^o_red = -. 41v. Ni^2+ + 2e- -> ni e^o_red = -. 23v. Anode: 3h2o -> 2h3o+ + 1/2o2 + 2e- e_red = 0. 82v. Electrolysis begins when deltae_external > |deltae_cell| at anode: H2o oxidized to o2 if e_red (ions) > 0. 82v (ph = 7)