I came across a compound $\ce{[FeO4]^{2-}}$ but $\ce{Fe}$ has +6 charge according to my calculations. How this is possible ? Or it is possible but my periodic table is kind of non-detailed one. (On the table I got $d$-metals possible charges).


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If you write out the electron configuration of Fe, you will find it has $3\text{d}^6 4\text{s}^2$ in its valence level, meaning it can theoretically take on any oxidation state from +1 to +8 (and even some negative states, for that matter).

In practice it is extremely uncommon to find any oxidation states other than +2 or +3 for iron - which explains why a simple periodic table containing only the most common naturally occurring states will omit them - but they are possible, just not very stable or long-lived. The polyatomic ion you have observed is called ferrate.

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    $\begingroup$ I've always enjoyed citing catalase as an example of a ubiquitous compound where iron takes on uncommon oxidation values. $\endgroup$ Commented Jun 1, 2013 at 20:40
  • $\begingroup$ Actually iron(VI) is not all that exotic, being in use commercially in water treatment. If we regard each oxide ligand as donating a pi electron pair as well as a sigma pair (a reasonable model for oxide ligands to electron-deficient metal-ion centers), then ferrate(VI) meets the 18-electron rule. $\endgroup$ Commented May 28, 2023 at 21:11

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