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{{chembox |
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{{Chembox |
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| verifiedrevid = 398461022 |
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| verifiedrevid = 428792085 |
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| Name = Potassium ferrate |
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| Name = Potassium ferrate. |
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| IUPACName = Potassium ferrate(VI) |
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| IUPACName = Potassium ferrate(VI) |
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| OtherNames = Potassium ferrate<br /> Dipotassium ferrate |
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| OtherNames = Potassium ferrate<br /> Dipotassium ferrate |
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| ImageFile1 = Potassium ferrate.svg |
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| ImageFile1 = Potassium ferrate.svg |
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| ImageFile2 = K2FeO4-xtal-1982-CM-3D-balls.png |
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| ImageFile2 = K2FeO4-xtal-1982-CM-3D-balls.png |
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| Section2 = {{Chembox Properties |
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|Section1 = {{Chembox Identifiers |
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| CASNo = 39469-86-8 |
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| Formula = K<sub>2</sub>FeO<sub>4</sub> |
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| PubChem = 53493006 |
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| MolarMass = 198.0392 g/mol |
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| StdInChI=1S/Fe.2K.4O/q+6;2*+1;4*-2 |
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| Appearance = Dark purple solid |
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| StdInChIKey = REKHNDAXGYXSBT-UHFFFAOYSA-N |
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| Density = 2.829 g/cm<sup>3</sup>, solid |
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| SMILES = ..()(=O)=O |
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| Solubility = soluble in 1M KOH |
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}} |
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| Solvent = other solvents |
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|Section2 = {{Chembox Properties |
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| SolubleOther = reacts with most solvents |
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| Formula = {{chem2|K2FeO4}} |
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| MeltingPt = >198 °C (decomposition temp) |
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| MolarMass = 198.0392 g/mol |
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}} |
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| Appearance = Dark purple solid |
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| Section3 = {{Chembox Structure |
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| Density = 2.829 g/cm<sup>3</sup> |
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| Coordination = Tetrahedral |
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| Solubility = soluble in 1M KOH |
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| CrystalStruct = K<sub>2</sub>SO<sub>4</sub> motif |
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| Solvent = other solvents{{which|date=October 2022}} |
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| Dipole = 0 ] |
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| SolubleOther = reacts with most solvents |
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}} |
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| MeltingPt = >198 °C (decomposes)}} |
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| Section7 = {{Chembox Hazards |
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|Section3={{Chembox Structure |
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| ExternalMSDS = |
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| Coordination = Tetrahedral |
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| MainHazards = oxidizer |
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| CrystalStruct = {{chem2|K2SO4}} motif |
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| FlashPt = non-combustible |
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| Dipole = 0 ]}} |
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| RPhrases = 8 |
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|Section7={{Chembox Hazards |
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| SPhrases = 17-36 |
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| ExternalSDS = |
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}} |
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| GHSPictograms = {{GHS03}}<ref name=sds>{{cite web|title =Potassium Ferrate|url=https://www.americanelements.com/potassium-ferrate-39469-86-8|publisher = ]|access-date = June 13, 2019}}</ref> |
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| Section8 = {{Chembox Related |
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| GHSSignalWord = Danger<ref name=sds /> |
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| OtherAnions = ]<br /> ]<br /> K<sub>2</sub>RuO<sub>4</sub> |
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| HPhrases = {{H-phrases|272}}<ref name=sds /> |
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| OtherCations = BaFeO<sub>4</sub><br /> Na<sub>2</sub>FeO<sub>4</sub>}} |
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| PPhrases = {{P-phrases|210|220|221|280|370+378|501}}<ref name=sds /> |
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}} |
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| MainHazards = Oxidizer |
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| FlashPt = non-combustible}} |
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|Section8={{Chembox Related |
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| OtherAnions = ]<br />]<br />{{chem2|K2RuO4}} |
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| OtherCations = ]<br />]}}}} |
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'''Potassium ferrate''' is an inorganic compound with the ] '''{{chem2|K2FeO4}}'''. It is the potassium salt of ]. Potassium ferrate is a powerful oxidizing agent with applications in green chemistry, organic synthesis, and cathode technology. |
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==Synthesis== |
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'''Potassium ferrate''' is the ] with the ] K<sub>2</sub>FeO<sub>4</sub>. This purple ] is ], and is a rare example of an ](VI) compound. In most of its compounds, iron has the oxidation state +2 or +3 (Fe<sup>2+</sup> or Fe<sup>3+</sup>). Reflecting its high oxidation state, ] is a powerful ]. |
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Generally, there are three ways to produce hexavalent iron: dry oxidation, wet oxidation, and electrochemical synthesis.<ref name=":0">{{Cite journal |last=Talaiekhozani |first=Amirreza |last2=Bagheri |first2=Marzieh |last3=Talaei |first3=Mohammad Reza |last4=Jaafarzadeh |first4=Nematollah |date=2016 |title=An Overview on Production and Applications of Ferrate(VI) |url=https://brieflands.com/articles/jjhs-15057.html |journal=Jundishapur Journal of Health Sciences |volume=8 |issue=3 |doi=10.17795/jjhs-34904 |issn=2345-4075}}</ref> The methods used to produce potassium ferrate are similar to those used to produce ] and ]. |
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=== Dry oxidation === |
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K<sub>2</sub>FeO<sub>4</sub> has attracted interest for applications in "]" because the by-products of its use, iron oxides, are environmentally innocuous. In contrast, some related oxidants such as ] are considered environmentally hazardous. However, the main difficulty with the use of K<sub>2</sub>FeO<sub>4</sub> is that it is often too reactive, as indicated by the fact that it decomposes in contact with water.<ref>Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.</ref> |
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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.<ref name=":0" /><ref name=":1">{{Cite journal |last=Talaiekhozani |first=Amirreza |last2=Talaei |first2=Mohammad Reza |last3=Rezania |first3=Shahabaldin |date=2017-04-01 |title=An overview on production and application of ferrate (VI) for chemical oxidation, coagulation and disinfection of water and wastewater |url=https://linkinghub.elsevier.com/retrieve/pii/S2213343717301161 |journal=Journal of Environmental Chemical Engineering |volume=5 |issue=2 |pages=1828–1842 |doi=10.1016/j.jece.2017.03.025 |issn=2213-3437}}</ref> |
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=== Wet oxidation === |
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: 4 K<sub>2</sub>FeO<sub>4</sub> + 4 H<sub>2</sub>O → 3 O<sub>2</sub> + 2 Fe<sub>2</sub>O<sub>3</sub> + 8 KOH |
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In the wet oxidation method, {{chem2|K2FeO4}} is prepared by ] an ] solution of an iron(III) salt. Generally, this method employs either ferrous (Fe<sup>II</sup>) or ferric (Fe<sup>III</sup>) salts as the source of iron ions, calcium, ] (Ca(ClO)<sub>2</sub>, NaClO), ] (Na<sub>2</sub>S<sub>2</sub>O<sub>3</sub>) or chlorine (Cl<sub>2</sub>) as oxidizing agents and, finally, sodium hydroxide, sodium carbonate (NaOH, NaCO<sub>3</sub>) or potassium hydroxide (KOH) to increase the pH of the solution.<ref>{{Cite journal |last=White |first=D. A. |last2=Franklin |first2=G. S. |date=November 1998 |title=A Preliminary Investigation into the Use of Sodium Ferrate in Water Treatment |url=http://www.tandfonline.com/doi/abs/10.1080/09593331908616776 |journal=Environmental Technology |volume=19 |issue=11 |pages=1157–1161 |doi=10.1080/09593331908616776 |issn=0959-3330}}</ref><ref>{{Cite journal |last=Munyengabe |first=Alexis |last2=Zvinowanda |first2=Caliphs |date=2019 |title=Production, Characterization and Application of Ferrate(VI) in Water and Wastewater Treatments |url=http://www.brjac.com.br/artigos/2019-V6-N25/brjac-19-2019.pdf |journal=Brazilian Journal of Analytical Chemistry |volume=6 |issue=25 |doi=10.30744/brjac.2179-3425.RV-19-2019}}</ref><ref>{{Cite book |last=Schreyer |first=J. M. |title=Inorganic Syntheses |last2=Thompson |first2=G. W. |last3=Ockerman |first3=L. T. |date=1953 |editor-last=Bailar |editor-first=John C. |volume=IV |pages=164–168 |chapter=Potassium Ferrate(VI) |chapter-url=https://sites.lsa.umich.edu/jbuss/wp-content/uploads/sites/811/2020/08/inorganic-synthesis04-1.pdf#page=174}}</ref> For example:<blockquote>{{chem2|3 ClO− + 3 Fe(OH)3(H2O)3 + 4 K+ + 4 OH− → 3 Cl− + 2 K2FeO4 + 11 H2O}}</blockquote> |
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=== Electrochemical synthesis === |
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==Synthesis and structure== |
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Electrochemical methods used to synthesize potassium ferrate usually consist of an iron anode which electrolyzes a KOH solution.<ref name=":1" /> |
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] (1660 – 1734) first discovered that the residue formed by igniting a mixture of ] (saltpetre) and ] powder dissolved in water to give a purple solution. ] (1814 – 1894) later discovered that fusion of ] and iron(III) oxide in air produced a compound that was soluble in water. The composition corresponded to that of ]. In the laboratory, K<sub>2</sub>FeO<sub>4</sub> is prepared by oxidizing an ] solution of an iron(III) salt with concentrated ].<ref>Schreyer, J. M.; Thompson, G. W.; Ockerman, L. T. "Potassium Ferrate(VI)" Inorganic Syntheses, 1953 volume IV, pages 164-168.</ref> |
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==Properties== |
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The salt is isostructural with ], ], and ]. The solid consists of K<sup>+</sup> and the tetrahedral FeO<sub>4</sub><sup>2−</sup> anion, with Fe-O distances of 1.66 Å.<ref>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}}.</ref> The poorly soluble barium salt, ], is also known. |
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] |
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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 ] and is ] with ], ], and ]. The solid consists of {{chem2|K+}} and the tetrahedral {{chem2|FeO4(2−)}} anion, with Fe-O distances of 1.66 Å.<ref>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}}.</ref> Potassium ferrate decomposes rapidly in neutral and acidic water, e.g.:<ref>Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. {{ISBN|0-12-352651-5}}.</ref> |
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:{{chem2|4 K2FeO4 + 4 H2O → 3 O2 + 2 Fe2O3 + 8 KOH}} |
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In alkaline solution and as a dry solid, {{chem2|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).<ref>{{Cite journal |last=Jiang |first=Jia-Qian |last2=Wang |first2=S. |last3=Panagoulopoulos |first3=A. |date=2006-04-01 |title=The exploration of potassium ferrate(VI) as a disinfectant/coagulant in water and wastewater treatment |url=https://linkinghub.elsevier.com/retrieve/pii/S0045653505010209 |journal=Chemosphere |volume=63 |issue=2 |pages=212–219 |doi=10.1016/j.chemosphere.2005.08.020 |issn=0045-6535}}</ref> |
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== Applications == |
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==Properties and applications== |
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Like ], {{chem2|K2FeO4}} generally does not generate environmentally toxic by-products and can be used in water treatment processes.<ref name=":0" /> It can act as: |
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As a dry solid, K<sub>2</sub>FeO<sub>4</sub> is stable. It decomposes with evolution of O<sub>2</sub> in neutral water, and especially rapidly in acidic water. At high ], aqueous solutions are stable. The deep purple solutions are similar in appearance to ] ({{chem|KMnO|4}}). It is stronger oxidizing agent than the latter. |
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* ]: promoting the oxidation of organic species in metal complexes. |
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Because the side products of its redox reactions are rust-like iron oxides, K<sub>2</sub>FeO<sub>4</sub> has been described as a "green oxidant." It has been employed in ] as an oxidant for organic contaminants and as a biocide. Conveniently, the resulting reaction product is iron(III) oxyhydroxide, an excellent ]. |
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* ]: allows removal of inorganic pollution compounds such as heavy metals, inorganic salts, trace elements and metal complexes. |
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* ]: destroys human pathogens including viruses, spores, bacteria and protozoa. |
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In addition, potassium ferrate can be used as a ] ] for fresh wounds.<ref>{{cite web |year=2016 |title=How WoundSeal Works |url=http://woundseal.com/wp/how-it-works |publisher=WoundSeal}}</ref><ref>{{cite patent|country=WO|number=2014153566|title=Hemostatic device and method|status=application|pubdate=2014-09-25|invent1=John Hen|invent2=Talmadge Kelly Keene|invent3=Mark Travi|assign1=Biolife, LLC}}</ref> In ], {{chem2|K2FeO4}} ] primary ]s.<ref>Green, J. R. "Potassium Ferrate" Encyclopedia of Reagents for Organic Synthesis 2001, John Wiley. {{doi|10.1002/047084289X.rp212}}.</ref> {{chem2|K2FeO4}} has also attracted attention as a potential ] material in a "]."<ref>{{Cite journal |last=Wang |first=Suqin |last2=Yang |first2=Zhanhong |last3=Liu |first3=Dongren |last4=Yi |first4=Shi |last5=Chi |first5=Weiwei |date=2010-03-01 |title=Evaluation of potassium ferrate(VI) cathode material coated with 2,3-Naphthalocyanine for alkaline super iron battery |url=https://linkinghub.elsevier.com/retrieve/pii/S1388248109006341 |journal=Electrochemistry Communications |volume=12 |issue=3 |pages=367–370 |doi=10.1016/j.elecom.2009.12.036 |issn=1388-2481}}</ref> |
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In ], K<sub>2</sub>FeO<sub>4</sub> ] primary ]s.<ref>Green, J. R. “Potassium Ferrate” Encyclopedia of Reagents for Organic Synthesis 2001, John Wiley. {{DOI|10.1002/047084289X.rp212}}.</ref> |
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Stabilised forms of potassium ferrate have been proposed for the removal of ]s, both dissolved and suspended, from ]s.<ref>{{Cite journal |last=Petrov |first=Vladimir G. |last2=Perfiliev |first2=Yury D. |last3=Dedushenko |first3=Sergey K. |last4=Kuchinskaya |first4=Tatiana S. |last5=Kalmykov |first5=Stepan N. |date=October 2016 |title=Radionuclide removal from aqueous solutions using potassium ferrate(VI) |url=http://link.springer.com/10.1007/s10967-016-4867-5 |journal=Journal of Radioanalytical and Nuclear Chemistry |volume=310 |issue=1 |pages=347–352 |doi=10.1007/s10967-016-4867-5 |issn=0236-5731}}</ref> Tonnage quantities were proposed to help remediate the effects of the ] in ] {{citation needed|date=March 2017}}. 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|date=December 2020}} |
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K<sub>2</sub>FeO<sub>4</sub> has also attracted attention as a potential ] material in a "]." |
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Because the side products of its ] reactions are rust-like iron oxides, {{chem2|K2FeO4}} has been described as an "]" ]. In contrast, related oxidants such as ]s are considered environmentally hazardous.<ref>{{Cite journal |last=Sharma |first=Virender K. |date=2002 |title=Potassium ferrate(VI): an environmentally friendly oxidant |journal=Advances in Environmental Research |volume=6 |issue=2 |pages=143–156 |doi=10.1016/s1093-0191(01)00119-8 |issn=1093-0191}}</ref> |
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== History == |
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In 1702, ] (1660 – 1734) observed that the ignition product of ] (saltpetre) and ] 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, ] (1814 – 1894) discovered that fusion of ] and ] in air produced a high-capacity iron compound that was soluble in water:<ref name=":1" /> |
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:{{chem2|8 KOH + 2 Fe2O3 + 3 O2 → 4 K2FeO4 + 4 H2O}} |
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==References== |
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==References== |
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{{Iron compounds}} |
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