Potassium ferrate.
Names
IUPAC name
Potassium ferrate(VI)
Other names
Potassium ferrate
Dipotassium ferrate
Identifiers
3D model (JSmol)
  • InChI=1S/Fe.2K.4O/q+6;2*+1;4*-2
    Key: REKHNDAXGYXSBT-UHFFFAOYSA-N
  • [K+].[K+].[O-][Fe]([O-])(=O)=O
Properties
K2FeO4
Molar mass 198.0392 g/mol
Appearance Dark purple solid
Density 2.829 g/cm3
Melting point >198 °C (decomposes)
soluble in 1M KOH
Solubility in other solvents[which?] reacts with most solvents
Structure
K2SO4 motif
Tetrahedral
0 D
Hazards
Occupational safety and health (OHS/OSH):
Main hazards
 Oxidizer
GHS labelling:
GHS03: Oxidizing[1]
Danger[1]
H272[1]
P210, P220, P221, P280, P370+P378, P501[1]
Flash point non-combustible
Safety data sheet (SDS) External SDS
Related compounds
Other anions
 K2MnO4
K2CrO4
K2RuO4
Other cations
 BaFeO4
Na2FeO4
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
checkY verify (what is checkY☒N ?)

Potassium ferrate is the chemical compound with the formula K2FeO4. This purple salt is paramagnetic, and is a rare example of an iron(VI) compound. In most of its compounds, iron has the oxidation state +2 or +3 (Fe2+ or Fe3+). Reflecting its high oxidation state, FeO2−4 is a powerful oxidizing agent.

Synthesis and structure

An aqueous solution of potassium ferrate(VI).

Georg Ernst Stahl (1660 – 1734) first discovered that the residue formed by igniting a mixture of potassium nitrate (saltpetre) and iron powder dissolved in water to give a purple solution. Edmond Frémy (1814 – 1894) later discovered that fusion of potassium hydroxide and iron(III) oxide in air produced a compound that was soluble in water:

8 KOH + 2 Fe2O3 + 3 O2 → 4 K2FeO4 + 4 H2O

The composition corresponded to that of potassium manganate. In the laboratory, K2FeO4 is prepared by oxidizing an alkaline solution of an iron(III) salt with concentrated chlorine bleach.:[2]

3 ClO− + 3 Fe(OH)3(H2O)3 + 4 K+ + 4 OH− → 3 Cl− + 2 K2FeO4 + 11 H2O

The salt is isostructural with K2MnO4, K2SO4, and K2CrO4. The solid consists of K+ and the tetrahedral FeO2−4 anion, with Fe-O distances of 1.66 Å.[3] The poorly soluble barium salt, BaFeO4, is also known.

Properties and applications

The main difficulty with the use of K2FeO4 is that it is often too reactive, as indicated by the fact that it decomposes in contact with water, especially in acidic water:[4]

4 K2FeO4 + 4 H2O → 3 O2 + 2 Fe2O3 + 8 KOH

At high pH, aqueous solutions are stable. The deep purple solutions are similar in appearance to potassium permanganate (KMnO4). It is stronger oxidizing agent than the latter. As a dry solid, K2FeO4 is stable.

Because the side products of its redox reactions are rust-like iron oxides, K2FeO4 has been described as a "green oxidant". It has been employed in waste-water treatment as an oxidant for organic contaminants and as a biocide. Conveniently, the resulting reaction product is iron(III) oxyhydroxide, an excellent flocculant. In organic synthesis, K2FeO4 oxidizes primary alcohols.[5] In contrast, related oxidants such as chromates are considered environmentally hazardous.[6]

K2FeO4 has also attracted attention as a potential cathode material in a "super iron battery."

Stabilised forms of potassium ferrate have been proposed for the removal of transuranic species, both dissolved and suspended, from aqueous solutions. Tonnage quantities were proposed to help remediate the effects of the Chernobyl disaster in Belarus. This new technique was successfully applied for the removal of a broad range of heavy metals.[citation needed]

Work on the use of potassium ferrate precipitation of transuranics and heavy metals was carried out in the Laboratories of IC Technologies Inc. in partnership with ADC Laboratories, in 1987 though 1992. The removal of the transuranic species were done on samples from various Dept. of Energy nuclear sites in the USA.[citation needed]

It has been proposed as a bleeding stopper for fresh wounds.[7][8]

References

  1. ^ a b c d "Potassium Ferrate". American Elements. Retrieved June 13, 2019.
  2. ^ Schreyer, J. M.; Thompson, G. W.; Ockerman, L. T. (1953). "Potassium Ferrate(VI)" (PDF). In Bailar, John C. (ed.). Inorganic Syntheses. Vol. IV. pp. 164–168.
  3. ^ Hoppe, M. L.; Schlemper, E. O.; Murmann, R. K. "Structure of Dipotassium Ferrate(VI)" Acta Crystallographica 1982, volume B38, pp. 2237-2239. doi:10.1107/S0567740882008395.
  4. ^ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
  5. ^ Green, J. R. "Potassium Ferrate" Encyclopedia of Reagents for Organic Synthesis 2001, John Wiley. doi:10.1002/047084289X.rp212.
  6. ^ Sharma, Virender K. (2002). "Potassium ferrate(VI): an environmentally friendly oxidant". Advances in Environmental Research. 6 (2): 143–156. doi:10.1016/s1093-0191(01)00119-8. ISSN 1093-0191.
  7. ^ "How WoundSeal Works". WoundSeal. 2016.
  8. ^ WO application 2014153566, John Hen; Talmadge Kelly Keene & Mark Travi, "Hemostatic device and method", published 2014-09-25, assigned to Biolife, LLC 
Categories