Potassium ferrate

Names
IUPAC name

Potassium ferrate(VI)

Other names

Potassium ferrate
Dipotassium ferrate

Identifiers

CAS Number

* 39469-86-8

3D model (JSmol)

* Interactive image

ChemSpider |

* 28296166

PubChem CID

* 53493006

InChI

* InChI=1S/Fe.2K.4O/q+6;2*+1;4*-2

Key: REKHNDAXGYXSBT-UHFFFAOYSA-N

SMILES

* [K+].[K+].[O-][Fe]([O-])(=O)=O

Properties

Chemical formula

| K
2FeO

Molar mass | 198.0392 g/mol
Appearance  | Dark purple solid
Density | 2.829 g/cm3
Melting point | >198 °C (decomposes)

Solubility in water

| soluble in 1M KOH
Solubility in other solvents[which?] | reacts with most solvents
Structure

Crystal structure

| K
2SO
4 motif

Coordination geometry

| Tetrahedral

Dipole moment

| 0 D
Hazards
Main hazards | Oxidizer
GHS pictograms |
GHS Signal word | Danger

GHS hazard statements

| H272

GHS precautionary statements

| P210, P220, P221, P280, P370+378, P501
Flash point | non-combustible
Related compounds

Other anions

| K
2MnO

2CrO

2RuO

Other cations

| BaFeO

2FeO

Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

Y verify (what is YN ?)
Infobox references

Potassium ferrate is an inorganic compound with the formula K
2FeO
4. 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.

## Contents

* 1 Synthesis
* 1.1 Dry oxidation
* 1.2 Wet oxidation
* 1.3 Electrochemical synthesis
* 2 Properties
* 3 Applications
* 4 History
* 5 References

## Synthesis

Generally, there are three ways to produce hexavalent iron: dry oxidation, wet oxidation, and electrochemical synthesis.[citation needed] 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, oxidizing environment, most commonly created with potassium nitrate or potassium peroxide and potassium hydroxide.[1] The combination of high temperature (200-800 °C) and oxidizing agents 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.[2]

### Wet oxidation

In the wet oxidation method, K
2FeO
4 is prepared by oxidizing an alkaline solution of an iron(III) salt. Generally, this method employs ferric (FeIII) salts such as iron(III) hydroxide as the source of iron ions, an oxidizing agent such as hypochlorite (in the form of Ca(ClO)2 or NaClO) or chlorine (Cl2) and a strong alkali such as sodium hydroxide, sodium carbonate (NaOH, Na2CO3) or potassium hydroxide (KOH) to increase the pH of the solution.[3][4][5]

### Electrochemical synthesis

Electrochemical methods used to synthesize potassium ferrate usually consist of an iron anode which electrolyzes a KOH solution.[2]

## Properties

An aqueous solution of potassium ferrate(VI).

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 K
2MnO
4, K
2SO
4, and K
2CrO
4. The solid consists of K+
and the tetrahedral FeO2−
4 anion, with Fe-O distances of 1.66 Å.[6] Potassium ferrate decomposes rapidly in neutral and acidic water, e.g.:[7]

4 K
2FeO
4 \+ 4 H
2O → 3 O
2 \+ 2 Fe

3 \+ 8 KOH

In alkaline solution and as a dry solid, K
2FeO
4 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).[8]

## Applications

Like sodium ferrate, K
2FeO
4 generally does not generate environmentally toxic by-products and can be used in water treatment processes.[citation needed] 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.[9][10] In organic synthesis, K
2FeO
4 oxidizes primary alcohols.[11] K
2FeO
4 has also attracted attention as a potential cathode material in a "super iron battery."[12]

Stabilised forms of potassium ferrate have been proposed for the removal of transuranium elements, both dissolved and suspended, from aqueous solutions.[13] Tonnage quantities were proposed to help remediate the effects of the Chernobyl disaster in Ukraine . 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 through 1992. The removal of the transuranium elements was demonstrated on samples from various Dept. of Energy nuclear sites in the USA.{{citation needed|date=December 20 Because the side products of its redox reactions are rust-like iron oxides, K
2FeO
4 has been described as an "environmentally friendly" oxidant. In contrast, related oxidants such as chromates are considered environmentally hazardous.[14]

## 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:[2]

8 KOH + 2 Fe

3 \+ 3 O
2 → 4 K
2FeO
4 \+ 4 H

## References

1. ↑ Madarász, J., Zbořil, R., Homonnay, Z., Sharma, V. K., & Pokol, G. (2006). Thermal decomposition of iron(VI) oxides, K2FeO4 and BaFeO4, in an inert atmosphere. Journal of Solid State Chemistry, 179(5), 1426–1433. https://doi.org/10.1016/j.jssc.2006.01.054
2. ↑ 2.0 2.1 2.2 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. https://linkinghub.elsevier.com/retrieve/pii/S2213343717301161.
3. ↑ 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. Bibcode: 1998EnvTe..19.1157W. http://www.tandfonline.com/doi/abs/10.1080/09593331908616776.
4. ↑ Munyengabe, Alexis; Zvinowanda, Caliphs (2019). "Production, Characterization and Application of Ferrate(VI) in Water and Wastewater Treatments". Brazilian Journal of Analytical Chemistry 6 (25). doi:10.30744/brjac.2179-3425.RV-19-2019. http://www.brjac.com.br/artigos/2019-V6-N25/brjac-19-2019.pdf.
5. ↑ Schreyer, J. M.; Thompson, G. W.; Ockerman, L. T. (1953). "Potassium Ferrate(VI)". in Bailar, John C.. Inorganic Syntheses. IV. pp. 164–168. https://sites.lsa.umich.edu/jbuss/wp-content/uploads/sites/811/2020/08/inorganic-synthesis04-1.pdf#page=174.
6. ↑ 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.
7. ↑ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
8. ↑ 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. PMID 16213564. Bibcode: 2006Chmsp..63..212J. https://linkinghub.elsevier.com/retrieve/pii/S0045653505010209.
9. ↑ "How WoundSeal Works". WoundSeal.com - Stop Bleeding Instantly. WoundSeal. 2016. http://woundseal.com/wp/how-it-works.
10. ↑ ; Talmadge Kelly Keene & Mark Travi"Hemostatic device and method" WO patent application 2014153566, published 2014-09-25, assigned to Biolife, LLC
11. ↑ Green, J. R. "Potassium Ferrate" Encyclopedia of Reagents for Organic Synthesis 2001, John Wiley. doi:10.1002/047084289X.rp212.
12. ↑ 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. https://linkinghub.elsevier.com/retrieve/pii/S1388248109006341.
13. ↑ 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. Bibcode: 2016JRNC..310..347P. http://link.springer.com/10.1007/s10967-016-4867-5.
14. ↑ 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. Bibcode: 2002AdvER...6..143S.

* v
* t
* e

Potassium compounds

* KAsO2
* KBr
* KBrO3
* KCN
* KCNO
* KCl
* KClO
* KClO3
* KClO4
* KF
* KH
* KHCO2
* KHCO3
* KHF2
* KHS
* KHSO3
* KHSO4
* KH2AsO4
* KH2PO3
* KH2PO4
* KI
* KIO3
* KIO4
* KMnO4
* KNH2
* KN3
* K3N
* KNO2
* KNO3
* KOCN
* KOH
* KO2
* KPF6
* KSCN
* KCH3COO
* K2Al2O4
* K2CO3
* K2CrO4
* K2Cr2O7
* K2FeO4
* K2HPO4
* K2MnO4
* K2O
* K2O2
* K2PtCl4
* K2PtCl6
* K2S
* K2SeO4
* K2SO3
* K2SO4
* KHSO5
* K2S2O5
* K2S2O7
* K2S2O8
* K2Po
* K2SiO3
* K3[Fe(CN)6]
* K3[Fe(C2O4)3]
* K4[Fe(CN)6]
* K3PO4
* K4Mo2Cl8

* v
* t
* e

Iron compounds

Fe(-II)|

* Na2Fe(CO)4

Fe(0)|

* Fe(CO)5
* Fe2(CO)9
* Fe3(CO)12
* Fe(CO)3CH3COC2H2C6H6

Fe(I)|

* FeH

| Organoiron(I) compounds|

* (C5H5FeCO)2(CO)2

|

Fe(0,II)|

* Fe3C

Fe(II)|

* FeBr2
* FeCl2
* FeF2
* FeI2
* FeH2
* FeMoO4
* FeO
* Fe(OH)2
* FeS
* FeSO4
* FeSe
* FeSeO4
* FeSi2
* FeTiO3
* FeCO3
* FeC2O4
* Fe(C2H3O2)2
* Fe(C3H5O3)2
* FeC6H6O7
* FeC12H22O14
* Fe2CrO4
* Fe3(PO4)2
* H2Fe(CO)4
* Fe(HCO3)2
* Mg2FeH6

| Organoiron(II) compounds|

* Fe(C5H5)2
* Fe(C5H4P(C6H5)2)2
* C4H4Fe(CO)3
* C4H6Fe(CO)3

|

Fe(0,III)|

* Fe3P

Fe(II,III)|

* Fe3O4
* Fe3S4

Fe(III)|

* FeBr3
* FeCl3
* FeF3
* FeI3
* Fe(NO3)3
* Fe(acac)3
* FeOCl
* FeO(OH)
* FePO4
* Fe4(P2O7)3
* Fe2(CrO4)3
* Fe2(C2O4)3
* Fe2O3
* Fe2S3
* Fe2(SO4)3
* FeBO3
* FeB

| Organoiron(III) compounds|

* Fe(C5H5)2BF4
* C6H8O7⋅xFe3+⋅yNH3

|

Fe(VI)|

* K2FeO4
* BaFeO4

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*[CID]: Compound ID
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*[P210]: Keep away from heat/sparks/open flames/hot surfaces – No smoking
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*[P221]: Take any precaution to avoid mixing with combustibles
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