This question involves the Nernst Equation. If we are given a reaction involving two elements undergoing a redox reaction a specific E'cell exists for standard conditions. This E'cell states whether a electrochemical reaction is spontaneous or not spontaneous. Now, in these kinds of reactions we don't start out a equilibrium. Because of this we use the Nernst Equation that has "Q" that describes the concentrations of the elements involved in the redox reaction. We can start out with a larger amount of reactants (Q less than 1) that equates the Ecell to be more positive, thus, more spontaneous. My question is this: how can a system in equilibrium still have a positive Ecell? I would think the reaction would proceed until the Ecell is zero.
To clarify as it may sound confusing: a redox reaction with nickel and zine has a cell potential of 51 at standard conditions. However, when more reactants are present the cell potential is greater meaning that it is more spontaneous. When equilibrium is finally achieved the cell potential is at its standard, 51. Wouldn't it go to zero? Why would a positive standard potential be at equilibrium when it surely is still spontaneous?
To clarify as it may sound confusing: a redox reaction with nickel and zine has a cell potential of 51 at standard conditions. However, when more reactants are present the cell potential is greater meaning that it is more spontaneous. When equilibrium is finally achieved the cell potential is at its standard, 51. Wouldn't it go to zero? Why would a positive standard potential be at equilibrium when it surely is still spontaneous?
