Ferrate(VI) is the inorganic anion with the chemical formula [FeO<sub>4</sub>]2−. It is photosensitive, contributes a pale violet colour to compounds and solutions containing it and is one of the strongest water-stable oxidizing species known. Although it is classified as a weak base, concentrated solutions containing ferrate(VI) are corrosive and attack the skin and are only stable at high pH. It is similar to the somewhat more stable permanganate.
See also: Ferrate.
The term ferrate is normally used to mean ferrate(VI), although it can refer to other iron-containing anions, many of which are more commonly encountered than salts of [FeO<sub>4</sub>]2−. These include the highly reduced species disodium tetracarbonylferrate, and salts of the iron(III) complex tetrachloroferrate [FeCl<sub>4</sub>]− in 1-Butyl-3-methylimidazolium tetrachloroferrate. Although rarely studied, ferrate(V) [FeO<sub>4</sub>]3− and ferrate(IV) [FeO<sub>4</sub>]4− oxyanions of iron also exist. These too are called ferrates.[1]
Ferrate(VI) salts are formed by oxidizing iron in an aqueous medium with strong oxidizing agents under alkaline conditions, or in the solid state by heating a mixture of iron filings and powdered potassium nitrate.[2]
For example, ferrates are produced by heating iron(III) hydroxide with sodium hypochlorite in alkaline solution:
2 + 3 + 4 → 2 + 5 + 3
The anion is typically precipitated as the barium(II) salt, forming barium ferrate.
Fe(VI) is a strong oxidizing agent over the entire pH range, with a reduction potential (Fe(VI)/Fe(III) couple) varying from +2.2 V to +0.7 V versus SHE in acidic and basic media respectively.
+ 8 + 3 e− + 4 ; E0 = +2.20 V (acidic medium)
+ 4 + 3 e− + 5 ; E0 = +0.72 V (basic medium)
Because of this, the ferrate(VI) anion is unstable at neutral or acidic pH values, decomposing to iron(III):[3] The reduction goes through intermediate species in which iron has oxidation states +5 and +4. These anions are even more reactive than ferrate(VI). In alkaline conditions ferrates are more stable, lasting for about 8 to 9 hours at pH 8 or 9.[4]
Aqueous solutions of ferrates are pink when dilute, and deep red or purple at higher concentrations.[5] [6] The ferrate ion is a stronger oxidizing agent than permanganate,[7] and oxidizes ammonia to molecular nitrogen.[8]
The ferrate(VI) ion has two unpaired electrons and is thus paramagnetic. It has a tetrahedral molecular geometry, isostructural with the chromate and permanganate ions.
Ferrates are excellent disinfectants, and are capable of removing and destroying viruses.[9] They are also of interest as potential as an environmentally friendly water treatment chemical, as the byproduct of ferrate oxidation is the relatively benign iron(III).[10]
Sodium ferrate is a useful reagent with good selectivity and is stable in aqueous solution of high pH, remaining soluble in an aqueous solution saturated with sodium hydroxide.