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REDOX & Electrochemistry - LSU Chemistry

REDOX & Electrochemistry - LSU Chemistry


REDOX 26Example: What about the reaction of Cu withnitric acid? Will this be a spontaneous rxn?Nitric acid is not like the other strong acids (HCl,HBr, HI, H 2 SO 4 ) in that it is a good oxidizing aciddue to the presence of the NO 3− anion, which is notjust a simple inert counter-anion for H + . Thecombination of NO 3− and H + makes for a ratherstrong oxidizing mixture.3[Cu(s) Cu 2+ (aq) + 2e - ] −0.342[NO 3−(aq) + 4H + (aq) + 3e -NO(g) + 2H 2 O]+0.963Cu(s) + 8H + (aq) + 2NO 3−(aq)3Cu 2+ (aq) + 2NO(g) + 4H 2 O+0.62VIf you get nitric acid on your skin, you will not onlyfeel the burning of the acid (H + ), but your skin willbe oxidized to a yellow-brown color! Soconcentrated nitric acid is doubly dangerous!

REDOX 27Although one can add half-cell rxns to yield overallredox equations, one can not simply add two halfcellrxns to yield another half-cell rxn. Forexample, consider the addition of the following twohalf cell rxns to generate a third half cell rxn:Cu 2+ (aq) + 1e - Cu + (aq) +0.16 VCu + (aq) + 1e - Cu(s) +0.52 VCu 2+ (aq) + 2e - Cu(s) +0.68 VThe correct half-cell potential for this rxn is 0.34 V,which is exactly half of what we incorrectlyattempted to calculate above. This factor of ½comes from the fact that all the half-cell potentialvalues are normalized to a single e- value even ifmultiple e- are used in the half cell rxn.As you might expect, the electrochemical potentialfor a rxn is directly related to the ΔG for a reaction(only with an opposite sign relationship!):ΔGº = −nFEºn = # of electrons being transferred, F = Faraday’sconstant (96.5 kJ/Vmol), Eº = standard potential.

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