Ignore the moles of electrons for now. What is the reaction?

(Could be it is just a lousy wording, but knowing the context first won't hurt).

The cell is as follows...

Cd | Cd(2+)(aq) || Ni(2+)(aq) | Ni

The question itself asks me to calculate the redox reaction, the standard cell potential, DeltaG, the Equilibrium Constant, and also the cell potential if the Cadmium cell is 0.010M and the Nickel cell is 1 M.

So...

Ox = Cd -> Cd(2+) + 2e

Red = Ni(2+) + 2e -> Ni

The reduction potential for cadmium and nickel -0.40V and -0.25V respectively.

So, Ecell = -0.25 + 0.4 = 0.15

DeltaG = -2(96485)(0.15)

= -28945J / -28.9 kJ

K = e^{2(96485)(0.15) / 8.314(298)}

= 118,540 (or 1.2 x 10^5)

Using Nernst Equation...

0.15 - [8.314(298) / 2(96485)] ln [1/0.01]

= 0.15 - (0.012839156)(4.605170186)

= 0.15 - 0.0591265

= 0.0908V

I'm not sure if I'm calculating all of this right - the equilibrium constant seems quite high for my liking, but not too sure...