Potassium ferrate
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| Potassium ferrate | |
|---|---|
| IUPAC name | Potassium ferrate(VI) |
| Other names | Potassium ferrate Dipotassium ferrate |
| Properties | |
| Molecular formula | K2FeO4 |
| Molar mass | 198.0392 g/mol |
| Appearance | Dark purple solid |
| Density | 2.829 g/cm3, solid |
| Melting point |
>198 °C (decomposition temp) |
| Solubility in water | soluble in 1M KOH |
| Solubility in other solvents | reacts with most solvents |
| Structure | |
| Crystal structure | K2SO4 motif |
| Coordination geometry |
Tetrahedral |
| Dipole moment | 0 D |
| Hazards | |
| Main hazards | oxidizer |
| R-phrases | 8 |
| S-phrases | 17-36 |
| Flash point | non-combustible |
| Related compounds | |
| Other anions | K2MnO4 K2CrO4 K2RuO4 |
| Other cations | BaFeO4 Na2FeO4 |
| Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa) Infobox disclaimer and references |
|
Potassium ferrate is the chemical compound with the formula K2FeO4. This purple, paramagnetic salt is a rare example of an Fe(VI) compound. In most of its compounds, iron has the oxidation state II (i.e., Fe2+) or III (Fe3+). Reflecting its high oxidation state, FeO42− is a powerful oxidant. K2FeO4 has attracted interest for applications in "Green Chemistry" because the by-products of its use, iron oxides, are environmentally innocuous. In contrast, some related oxidants such as chromate are considered environmentally hazardous. The main difficulty with the use of K2FeO4 is that it is often too reactive, as indicated by the fact that it decomposes in water.[1]
[edit] Redox properties and applications
As a dry solid, K2FeO4 is stable. It decomposes with evolution of O2 in neutral and especially rapidly in acidic water. At high pH, aqueous solutions are stable. The deep purple solutions are similar in appearance to potassium permanganate. K2FeO4 is a stronger oxidizing agent than KMnO4.
Because the side products of its redox reactions are rust-like iron oxides, K2FeO4 has been described as a "green oxidant." Indeed it has been employed for waste-water treatment as an oxidant for organic contaminants and as a biocide. Conveniently, the reduced iron(III) oxyhydroxide is an excellent flocculant.
In organic synthesis, K2FeO4 oxidizes primary alcohols.[2]
It has also attracted attention as a potential cathode material in a "Super iron battery."
[edit] Synthesis and structure
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 the compound produced was soluble in water corresponding to the composition of potassium manganate. In the laboratory, K2FeO4 is prepared by oxidizing an alkaline solution of an iron(III) salt with concentrated chlorine bleach.[3]
The salt is isostructural with K2MnO4, K2SO4, and K2CrO4. The solid consists of K+ and the tetrahedral FeO42− anion, with Fe-O distances of 1.66 Å.[4] The poorly soluble barium salt, BaFeO4, is also known.
[edit] References
- ^ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
- ^ Green, J. R. “Potassium Ferrate” Encyclopedia of Reagents for Organic Synthesis 2001, John Wiley. doi:10.1002/047084289X.rp212.
- ^ Schreyer, J. M.; Thompson, G. W.; Ockerman, L. T. "Potassium Ferrate(VI)" Inorganic Syntheses, 1953 volume IV, pages 164-168.
- ^ 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.
