8
$\begingroup$

Why does $\ce{FeF6}$ not exist? There are hexavalent iron compounds, so that is not the problem. There is the $\ce{[Fe(CN)6]^3-}$ ion, so there doesn't seem to be a steric problem.

$\endgroup$
4
  • 1
    $\begingroup$ Maybe reading about bonding characteristics of Potassium Ferrate may help. $\endgroup$ – Rishi Apr 17 at 4:43
  • 1
    $\begingroup$ No, it doesn't, as far as it is possible to prove a negative. Nor do CrF6, MnF7 or CoF5. And note your steric argument is not correct, Fe3+ is appreciably bigger than Fe6+, as far as those terms have any meaning in chemical environments. $\endgroup$ – Ian Bush Apr 17 at 8:51
  • 2
    $\begingroup$ Not an answer. But $\ce{FeF3}$ as a crystal already is something like $\ce{FeF6}$, in the sense that each iron is bound to six fluorine atoms; the lattice is linked via $\ce{-F-}$ bridges. The existence of $\ce{[Fe(CN)6]^3-}$ is more of an argument for $\ce{[FeF6]^3-}$ than $\ce{FeF6}$. $\endgroup$ – Linear Christmas Apr 17 at 12:50
  • $\begingroup$ Food for thought: The dimer of iron(III) fluoride doesn't exist(unlike iron(III) chloride), but the anion, $\ce{[Fe2F6]^-}$ does exist consisting of $\ce{Fe^2+}$ and $\ce{Fe^3+}$ ions and it form salts like $\ce{[NH4]Fe2F6}$. $\endgroup$ – Nilay Ghosh Apr 22 at 15:04
4
$\begingroup$

Besides steric factors related to the small size of the transition metal core, we could be seeing an electronic effect described in this answer. Iron is fairly early in the transition metal series, so when pushed beyond the $+3$ oxidation state it has few $d$ electrons in the central core. As explained in the referenced answer, this makes the iron strongly pi-electron accepting, and therefore more favorable for combining with a stronger pi-donor ligand such as oxide instead of fluoride.

$\endgroup$
5
  • $\begingroup$ Thank you! So: what about FeO2F2 or FeOF4..? $\endgroup$ – Thomas Forster Apr 18 at 9:37
  • $\begingroup$ You may want to look at en.wikipedia.org/wiki/High-valent_iron. Organic compounds do allow a wider variety of ligands, but all the examples given there (with iron in oxidation states of +4 to +6) still involve pi-donor ligands from Groups 15 and 16, primarily oxygen and nitrogen. $\endgroup$ – Oscar Lanzi Apr 18 at 9:48
  • $\begingroup$ @ThomasForster Not exactly FeO2F2 but there are similar types of compounds: $\ce{Fe3(OF)2}$, $\ce{Li2Fe4OF8}$, $\ce{FeSbO2F2}$. $\endgroup$ – Nilay Ghosh Apr 22 at 14:57
  • $\begingroup$ @nilay good finds, but no Fe(VI). They are just Fe(II) and Fe(III). $\endgroup$ – Oscar Lanzi Apr 22 at 15:34
  • $\begingroup$ This shows that iron can't achieve +6 o.s. that easily. $\endgroup$ – Nilay Ghosh Apr 23 at 3:33

Your Answer

By clicking “Post Your Answer”, you agree to our terms of service, privacy policy and cookie policy

Not the answer you're looking for? Browse other questions tagged or ask your own question.