# Greater ligand stabilization of the lower oxidation state increasing redox potential

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!