Potassium ferrate
Template:Chembox Potassium ferrate is an inorganic compound with the formula Template:Chem2. It is the potassium salt of ferric acid. Potassium ferrate is a powerful oxidizing agent with applications in green chemistry, organic synthesis, and cathode technology.
Synthesis
Generally, there are three ways to produce hexavalent iron: dry oxidation, wet oxidation, and electrochemical synthesis.Script error: No such module "Unsubst". The methods used to produce potassium ferrate are similar to those used to produce sodium ferrate and barium ferrate.
Dry oxidation
The dry oxidation method entails heating or melting iron oxides in an alkaline, oxygenated environment. The combination of high temperature (200 °C - 800 °C) and oxygen presents an explosion hazard that has led many researchers to believe this method of production is not suitable from a safety viewpoint, although many attempts have been made to overcome this problem.[1]
Wet oxidation
In the wet oxidation method, Template:Chem2 is prepared by oxidizing an alkaline solution of an iron(III) salt. Generally, this method employs either ferrous (FeII) or ferric (FeIII) salts as the source of iron ions, calcium, sodium hypochlorite (Ca(ClO)2, NaClO), sodium thiosulfate (Na2S2O3) or chlorine (Cl2) as oxidizing agents and, finally, sodium hydroxide, sodium carbonate (NaOH, Na2CO3) or potassium hydroxide (KOH) to increase the pH of the solution.[2][3][4] For example:
Electrochemical synthesis
Electrochemical methods used to synthesize potassium ferrate usually consist of an iron anode which electrolyzes a KOH solution.[1]
Properties
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Potassium ferrate is a dark purple crystalline solid that dissolves in water to form a reddish-purple solution. The salt is paramagnetic and is isostructural with [[Potassium manganate|Template:Chem2]], [[Potassium sulfate|Template:Chem2]], and [[Potassium chromate|Template:Chem2]]. The solid consists of Template:Chem2 and the tetrahedral Template:Chem2 anion, with Fe-O distances of 1.66 Å.[5] Potassium ferrate decomposes rapidly in neutral and acidic water, e.g.:[6]
In alkaline solution and as a dry solid, Template:Chem2 is stable. Under the acidic conditions, the oxidation–reduction potential of the ferrate(VI) ions (2.2 V) is greater than that of ozone (2.0 V).[7]
Applications
Like sodium ferrate, Template:Chem2 generally does not generate environmentally toxic by-products and can be used in water treatment processes.Script error: No such module "Unsubst". It can act as:
- Oxidizing agent: promoting the oxidation of organic species in metal complexes.
- Coagulator: allows removal of inorganic pollution compounds such as heavy metals, inorganic salts, trace elements and metal complexes.
- Disinfectant: destroys human pathogens including viruses, spores, bacteria and protozoa.
In addition, potassium ferrate can be used as a bleeding stopper for fresh wounds.[8][9] In organic synthesis, Template:Chem2 oxidizes primary alcohols.[10] Template:Chem2 has also attracted attention as a potential cathode material in a "super iron battery."[11]
Stabilised forms of potassium ferrate have been proposed for the removal of transuranium elements, both dissolved and suspended, from aqueous solutions.[12] Tonnage quantities were proposed to help remediate the effects of the Chernobyl disaster in Belarus Script error: No such module "Unsubst".. This new technique was successfully applied for the removal of a broad range of heavy metals. Work on the use of potassium ferrate precipitation of transuranium elements 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 transuranium elements was demonstrated on samples from various Dept. of Energy nuclear sites in the USA.Script error: No such module "Unsubst".
Because the side products of its redox reactions are rust-like iron oxides, Template:Chem2 has been described as an "environmentally friendly" oxidant. In contrast, related oxidants such as chromates are considered environmentally hazardous.[13]
History
In 1702, Georg Ernst Stahl (1660 – 1734) observed that the ignition product of potassium nitrate (saltpetre) and iron powder displayed a red-purple color in an aqueous solution, which was eventually attributed to hexavalent potassium ferrate. Eckenberg and Becquerel in 1834 reported that a red-purple color appeared during heating of a mixture of potassium hydroxide and iron ore. In 1840, Edmond Frémy (1814 – 1894) discovered that fusion of potassium hydroxide and iron(III) oxide in air produced a high-capacity iron compound that was soluble in water:[1]
References
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- ↑ Hoppe, M. L.; Schlemper, E. O.; Murmann, R. K. "Structure of Dipotassium Ferrate(VI)" Acta Crystallographica 1982, volume B38, pp. 2237-2239. Script error: No such module "CS1 identifiers"..
- ↑ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. Template:ISBN.
- ↑ Script error: No such module "Citation/CS1".
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- ↑ <templatestyles src="Citation/styles.css"/>Template:Citation/make link, John Hen; Talmadge Kelly Keene & Mark Travi, "Hemostatic device and method", published Script error: No such module "auto date formatter"., assigned to Biolife, LLCScript error: No such module "Check for unknown parameters".
- ↑ Green, J. R. "Potassium Ferrate" Encyclopedia of Reagents for Organic Synthesis 2001, John Wiley. Script error: No such module "CS1 identifiers"..
- ↑ Script error: No such module "Citation/CS1".
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