The correct option is A increase by 0.010V
Given,
Cu2++e→Cu+;−ΔGo1=Eo1F×1
Cu++e→Cu;−ΔGo2=Eo2F×1
----------------------------
by adding these equations
Cu2++2e→Cu;−ΔGo3=−[ΔGo1+ΔGo2]=−[0.15F+0.5F]
or 2×Eo3F=0.65F
(as ΔG=−nFEo)
∴Eo3=0.325V
So for reverse reaction.
∴Cu→Cu2++2eEo=−0.325V
Cu→Cu2++2eEo=−0.325V
2Ag++2e→2AgEo=0.799V
∴Eocell=EoOPCu/Cu2++EoRPAg+/Ag
=−0.325+0.799=0.474V
On passing the current through the given cell Ag will oxidise to Ag+ and Cu2+ will be reduced to Cu.
As same charge is passing throuth both solutions so
According to Faraday's law:
wE=i.t96500=9.65×360096500
=0.36 eq. of Ag+=0.36 mol of Ag+
=0.36 eq. of Cu2+
=0.362 mole of Cu2+ (as mole = equivalent/valency)
∴[Ag+]=1+0.36=1.36M
[Cu2+]=1.0−0.18=0.82M
For Ag|Ag+||Cu2+|Cu cell,
$E^o_{cell} = - 0.474 $
And according to Nernst equation,
E=EoOPAg+EoRPCu+0.0592log[Cu2+][Ag+]2
Initially Cu2+=1M and Ag+=1M
∴Ecell=Eocell=−0.474
After electrolysis Cu2+=0.82M,Ag+=1.36M
∴Ecell=Eocell+0.0592log0.82(1.36)2
=Eocell−0.010=−0.474−0.010=−0.484V
Eocell is negative for Ag+|Ag||Cu2+|Cu and is −0.474V
As for spontaneous reaction E should be positive and it was 0.474 V in first case and 0.484 V in II case so it increases by 0.010 V.