Potassium ferrate
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| Names | |
|---|---|
| IUPAC name
Potassium ferrate(VI)
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| Other names
Potassium ferrate
Dipotassium ferrate | |
| Identifiers | |
3D model (JSmol)
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PubChem CID
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| 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
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Oxidizer |
| GHS labelling: | |
 [1]
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| 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
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K2MnO4 K2CrO4 K2RuO4 |
Other cations
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BaFeO4 Na2FeO4 |
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Potassium ferrate is an inorganic compound with the formula K2FeO4. 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
[edit]Generally, there are three ways to produce hexavalent iron: dry oxidation, wet oxidation, and electrochemical synthesis.[2] The methods used to produce potassium ferrate are similar to those used to produce sodium ferrate and barium ferrate.
Dry oxidation
[edit]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 lead 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.[2][3]
Wet oxidation
[edit]In the wet oxidation method, K2FeO4 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, NaCO3) or potassium hydroxide (KOH) to increase the pH of the solution.[4][5][6] For example:
3 ClO− + 3 Fe(OH)3(H2O)3 + 4 K+ + 4 OH− → 3 Cl− + 2 K2FeO4 + 11 H2O
Electrochemical synthesis
[edit]Electrochemical methods used to synthesize potassium ferrate usually consist of an iron anode which electrolyzes a KOH solution.[3]
Properties
[edit]solution.png)
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 K2MnO4, K2SO4, and K2CrO4. The solid consists of K+ and the tetrahedral FeO2−4 anion, with Fe-O distances of 1.66 Å.[7] Potassium ferrate decomposes rapidly in neutral and acidic water, e.g.:[8]
- 4 K2FeO4 + 4 H2O → 3 O2 + 2 Fe2O3 + 8 KOH
In alkaline solution and as a dry solid, K2FeO4 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).[9]
Applications
[edit]Like sodium ferrate, K2FeO4 generally does not generate environmentally toxic by-products and can be used in water treatment processes.[2] 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.[10][11] In organic synthesis, K2FeO4 oxidizes primary alcohols.[12] K2FeO4 has also attracted attention as a potential cathode material in a "super iron battery."[13]
Stabilised forms of potassium ferrate have been proposed for the removal of transuranium elements, both dissolved and suspended, from aqueous solutions.[14] Tonnage quantities were proposed to help remediate the effects of the Chernobyl disaster in Belarus [citation needed]. 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.[citation needed]
Because the side products of its redox reactions are rust-like iron oxides, K2FeO4 has been described as an "environmentally friendly" oxidant. In contrast, related oxidants such as chromates are considered environmentally hazardous.[15]
History
[edit]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:[3]
- 8 KOH + 2 Fe2O3 + 3 O2 → 4 K2FeO4 + 4 H2O
References
[edit]- ^ a b c d "Potassium Ferrate". American Elements. Retrieved June 13, 2019.
- ^ a b c Talaiekhozani, Amirreza; Bagheri, Marzieh; Talaei, Mohammad Reza; Jaafarzadeh, Nematollah (2016). "An Overview on Production and Applications of Ferrate(VI)". Jundishapur Journal of Health Sciences. 8 (3). doi:10.17795/jjhs-34904. ISSN 2345-4075.
- ^ a b c Talaiekhozani, Amirreza; Talaei, Mohammad Reza; Rezania, Shahabaldin (2017-04-01). "An overview on production and application of ferrate (VI) for chemical oxidation, coagulation and disinfection of water and wastewater". Journal of Environmental Chemical Engineering. 5 (2): 1828–1842. doi:10.1016/j.jece.2017.03.025. ISSN 2213-3437.
- ^ White, D. A.; Franklin, G. S. (November 1998). "A Preliminary Investigation into the Use of Sodium Ferrate in Water Treatment". Environmental Technology. 19 (11): 1157–1161. doi:10.1080/09593331908616776. ISSN 0959-3330.
- ^ Munyengabe, Alexis; Zvinowanda, Caliphs (2019). "Production, Characterization and Application of Ferrate(VI) in Water and Wastewater Treatments" (PDF). Brazilian Journal of Analytical Chemistry. 6 (25). doi:10.30744/brjac.2179-3425.RV-19-2019.
- ^ 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.
- ^ 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.
- ^ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
- ^ Jiang, Jia-Qian; Wang, S.; Panagoulopoulos, A. (2006-04-01). "The exploration of potassium ferrate(VI) as a disinfectant/coagulant in water and wastewater treatment". Chemosphere. 63 (2): 212–219. doi:10.1016/j.chemosphere.2005.08.020. ISSN 0045-6535.
- ^ "How WoundSeal Works". WoundSeal. 2016.
- ^ WO application 2014153566, John Hen; Talmadge Kelly Keene & Mark Travi, "Hemostatic device and method", published 2014-09-25, assigned to Biolife, LLC
- ^ Green, J. R. "Potassium Ferrate" Encyclopedia of Reagents for Organic Synthesis 2001, John Wiley. doi:10.1002/047084289X.rp212.
- ^ Wang, Suqin; Yang, Zhanhong; Liu, Dongren; Yi, Shi; Chi, Weiwei (2010-03-01). "Evaluation of potassium ferrate(VI) cathode material coated with 2,3-Naphthalocyanine for alkaline super iron battery". Electrochemistry Communications. 12 (3): 367–370. doi:10.1016/j.elecom.2009.12.036. ISSN 1388-2481.
- ^ Petrov, Vladimir G.; Perfiliev, Yury D.; Dedushenko, Sergey K.; Kuchinskaya, Tatiana S.; Kalmykov, Stepan N. (October 2016). "Radionuclide removal from aqueous solutions using potassium ferrate(VI)". Journal of Radioanalytical and Nuclear Chemistry. 310 (1): 347–352. doi:10.1007/s10967-016-4867-5. ISSN 0236-5731.
- ^ 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.