For the cell reaction
2Fe³⁺(aq) + 2I^– (aq) $$ \to $$ 2Fe²⁺(aq) + I₂(aq)
$${E_{cell}^\Theta }$$ = 0.24 V at 298 K. The standard Gibbs energy ($$\Delta $$_rG^o) of the cell reaction is :
[Given that Faraday constant F = 96500 C mol^–1]

Consider the change in oxidation state of bromine corresponding to different emf values as shown in the given diagram :
Then the species undergoing disproportionation is :

In the electrochemical cell :
$$Zn\left| {ZnS{O_4}\left( {0.01\,M} \right)} \right|$$$$\left| {CuS{O_4}\left( {1.0M} \right)} \right|Cu,$$
the emf of this Daniell cell is E₁. When the concentration of ZnSO₄ is changed to 1.0 M and that of CuSO4 changed to 0.01 M, the emf changes to E₂. From the followings, which one is the relationship between E₁ and E₂? (Given, RT/F = 0.059)

If the E^o_cell for a given reaction has a negative value, which of the following gives the correct relationships for the values of $$\Delta $$G^o and K_eq ?