According to Atkins, P.W., De Paula, J. and Keeler, J., 2006, Atkins' Physical Chemistry,

$$ \begin{align} \ce{[Fe(CN)6]^3- + e- &<=> [Fe(CN)6]^4-} &\quad E^\circ &= \pu{+0.36 V} \label{rxn:1}\tag{R1} \\ \ce{Fe^3+ + e- &<=> Fe^2+} &\quad E^\circ &= \pu{+0.77 V} \label{rxn:2}\tag{R2} \end{align} $$

According to my lecture notes, if a ligand stabilizes a lower oxidation state more than the higher oxidation state, the redox potential increases i.e. becomes more positive. This makes sense as stabilizing $\ce{Fe^2+}$ in reaction \eqref{rxn:1} would have the position of equilibrium to the right, increasing electron charge transfer and hence the redox potential, meaning a more positive $E^\circ$ and hence more negative $ΔG$ as

$$ΔG = -nFE^\circ$$

Indeed, $\ce{Fe^2+}$ is more stabilised than $\ce{Fe^3+}$ by $\ce{CN-}$ as $\ce{Fe^2+}$ is a borderline soft/hard acid, $\ce{Fe^3+}$ is a hard acid and $\ce{CN^-}$ is a soft base.

I don't understand why $E^\circ$ becomes more positive going from reaction \eqref{rxn:1} to reaction \eqref{rxn:2}. Please help me to understand why!


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