A galvanic cell is Zn(s) | Zn^2+(0.01 M) || Cu^2+(1.0 M) | Cu(s) at 298 K with E^°cell=1.10 V. Take F=96500 C mol^-1 and 2.303RT/F=0.059 V.

Nernst equation for $n=2$: $E=E^\circ-\frac{0.059}{2}\log\frac{[\mathrm{Zn^{2+}}]}{[\mathrm{Cu^{2+}}]}=1.10-0.0295\log(0.01)=1.10+0.059=1.159$ V, so A is correct. Then $\Delta G=-nFE=-2(96500)(1.159)=-223.7$ kJ; the value $-212$ kJ is $\Delta G^\circ$, so B is wrong. At equilibrium $\log K=nE^\circ/0.059=2.2/0.059=37.3$, so $K\approx10^{37}$ and C (which halves the exponent by forgetting $n$) is wrong. If both concentrations are 0.01 M the ratio is 1, the log term vanishes and $E=E^\circ=1.10$ V, not 1.16 V, so D is wrong. Each option therefore tests a different slip: ignoring the concentration term, mixing standard and actual free energy, dropping n in the equilibrium constant, or mishandling the reaction quotient.