How to find the oxidation state of complexes with bridging ligands? [closed]








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![image of the di-bridged binuclear iron complex](https://i.stack.imgur.com/0m0JA.jpg)


How can I find the oxidation states of $\ce{Fe, OH-, NH2-}$ in the given complex?




Just follow the normal rules for determining oxidation states. Any bond between elements is cleaved *heterolyticly* giving only the electronegative element all electrons of that bond. Then count electrons.


All oxygens in there will have $\mathrm{-II}$, all nitrogens $\mathrm{-III}$, all hydrogens (they are either connected to oxygen or to nitrogen) $\mathrm{+I}$.


You also know that your cationic complex $\ce{[Fe(H2O)4\mu{-}OH\mu{-}NH2Fe(NH3)4]^4+}$ has to have a $4+$ charge due to the two sulphate counterions. That allows us to determine the average oxidation state of the iron ions:


$$\begin{align}5\times (-2) + 5 \times (-3) + 23\times (+1) + 2\,x &= +4\\
-10 - 15 + 23 + 2\,x &= +4\\
-2 + 2\,x &= +4\\
2\,x &= + 6\\
x &= +3\end{align}$$


Thus, these are most likely $\mathrm{+III}$. You can also take it down to a per-complex level, whereby you would equally distribute the sulphate charges to the two irons giving both a formal $2+$. Then, there are two bridging ligands each as monoanions, so they require each iron cation to have another positive charge resulting in $3+$ or iron(III).


Note that while iron(III) is the most likely configuration, it does not have to be that. It may also be an iron(II,IV) complex theoretically. Be prepared for unexpected exceptions.


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