The question asks: Recall that you observed very little corrosion occurring on the iron nail immersed in NaOH(aq) solution. This observation is difficult to explain from an electrochemistry perspective since electrochemistry principles predict a spontaneous reaction that should cause corrosion. Explain why there was no corrosion on this nail from an equilibrium perspective using the half-reaction shown below: O2(g) + 2 H2O(l) + 4 e- --> 4 OH-(aq) Originally I had the thought to regard the electron concentration as an indicator of equilibrium, thinking that since the electron concentration is shifted reactant side, no reaction will occur. Now I'm beginning to second guess that hypothesis because the half reaction doesn't explicitly deal with the production of an iron oxide, therefore I find it a presumptuous to say the electron concentration in the OH-(aq) half reaction dictates a successful redox reaction. Any thoughts?