Ferrate(VI)
| Ferrate(VI) | |
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 Solutions of ferrate (left)
and permanganate (right) |
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Ferrate(VI) |
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Tetraoxoironbis(olate)[citation needed] |
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Other names
Ferrate(2-) |
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| Identifiers | |
| PubChem | 25000034 |
| ChemSpider | 21865127  |
| ChEBI | CHEBI:30992 |
| Jmol-3D images | Image 1 Image 2 |
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| Properties | |
| Molecular formula | FeO42- |
| Molar mass | 119.843 g mol-1 |
| Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa) | |
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| Infobox references | |
Ferrate(VI) is the inorganic anion with the chemical formula [FeO4]2-. It is photosensitive, contributes a pale violet colour to compounds and solutions containing it and is one of the strongest water-stable oxidising species known. Although it is classified as a weak base, concentrated solutions containing ferrate(VI) are corrosive and attack the skin and are only stable at high pH.
Nomenclature[edit]
The term ferrate is normally used to mean ferrate(VI), although it can refer to other iron-containing anions, many of which are more commonly encountered than salts of [FeO4]2-. These include the highly reduced species disodium tetracarbonylferrate Na2[Fe(CO)4] and salts of the iron(III) complex tetrachloroferrate [FeCl4]-. Although rarely studied, ferrate(V) [FeO4]3- and ferrate(IV) [FeO4]4- oxyanions of iron also exist. These too are called ferrates.[1]
Synthesis[edit]
Ferrate(VI) salts are formed by oxidizing iron in an aqueous medium with strong oxidizing agents under alkaline conditions, or in the solid state by heating a mixture of iron filings and powdered potassium nitrate.[2]
For example, ferrates are produced by heating iron(III) hydroxide with sodium hypochlorite in alkaline solution:[3]
- 2 Fe(OH)
3 + 3 OCl−
+ 4 OH−
→ 2 [FeO
4]2−
+ 5 H
2O + 3 Cl−
The anion is typically precipitated as the barium(II) salt, forming barium ferrate.[3]
Properties[edit]
The ferrate(VI) anion is unstable at neutral[2] or acidic pH values, decomposing to iron(III):[3]
- [FeO
4]2−
+ 3 -
e + 8 H+
Fe3+
+ 4 H
2O
The reduction goes through intermediate species in which iron has oxidation states +5 and +4.[4] These anions are even more reactive than ferrate(VI).[5] In alkaline conditions ferrates are more stable, lasting for about 8 to 9 hours at pH 8 or 9.[5]
Aqueous solutions of ferrates are pink when dilute, and deep red or purple at higher concentrations.[4][6] The ferrate ion is a stronger oxidizing agent than permanganate,[7] and will oxidize chromium(III) to dichromate,[8] and ammonia to molecular nitrogen.[9]
Ferrates are excellent disinfectants, and are capable of removing and destroying viruses.[10]
The ferrate(VI) ion has two unpaired electrons, and is thus paramagnetic. It has a tetrahedral molecular geometry.[4]
See also[edit]
References[edit]
- ^ Graham Hill; John Holman (2000). Chemistry in context (5th ed.). Nelson Thornes. p. 202. ISBN 0-17-448276-0.
- ^ a b R. K. Sharma (2007). Text Book Of Coordination Chemistry. Discovery Publishing House. pp. 124–125. ISBN 81-8356-223-X.
- ^ a b c Gary Wulfsberg (1991). Principles of descriptive inorganic chemistry. University Science Books. pp. 142–143. ISBN 0-935702-66-0.
- ^ a b c Egon Wiberg; Nils Wiberg; Arnold Frederick Holleman (2001). Inorganic chemistry. Academic Press. pp. 1457–1458. ISBN 0-12-352651-5.
- ^ a b Gary M. Brittenham (1994). Raymond J. Bergeron, ed. The Development of Iron Chelators for Clinical Use. CRC Press. pp. 37–38. ISBN 0-8493-8679-9.
- ^ John Daintith, ed. (2004). Oxford dictionary of chemistry (5th ed.). Oxford University Press. p. 235. ISBN 0-19-860918-3.
- ^ Kenneth Malcolm Mackay; Rosemary Ann Mackay; W. Henderson (2002). Introduction to modern inorganic chemistry (6th ed.). CRC Press. pp. 334–335. ISBN 0-7487-6420-8.
- ^ Amit Arora (2005). Text Book Of Inorganic Chemistry. Discovery Publishing House. pp. 691–692. ISBN 81-8356-013-X.
- ^ Karlis Svanks (June 1976). "Oxidation of Ammonia in Water by Ferrates(VI) and (IV)" (PDF). Water Resources Center, Ohio State University. p. 3. Retrieved 2010-05-04.
- ^ Stanley E. Manahan (2005). Environmental chemistry (8th ed.). CRC Press. p. 234. ISBN 1-56670-633-5.
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