Ferrate(VI)

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Ferrate(VI)
Aromatic skeletal formula of ferrate
Ferrate and permanganate solution.jpg
Solutions of ferrate (left)
and permanganate (right)
Identifiers
PubChem 25000034
ChemSpider 21865127 YesY
ChEBI CHEBI:30992 YesY
Jmol-3D images Image 1
Image 2
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
2
O
+ 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+
is in equilibrium with Fe3+
+ 4 H
2
O

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]

  1. ^ Graham Hill; John Holman (2000). Chemistry in context (5th ed.). Nelson Thornes. p. 202. ISBN 0-17-448276-0. 
  2. ^ a b R. K. Sharma (2007). Text Book Of Coordination Chemistry. Discovery Publishing House. pp. 124–125. ISBN 81-8356-223-X. 
  3. ^ a b c Gary Wulfsberg (1991). Principles of descriptive inorganic chemistry. University Science Books. pp. 142–143. ISBN 0-935702-66-0. 
  4. ^ a b c Egon Wiberg; Nils Wiberg; Arnold Frederick Holleman (2001). Inorganic chemistry. Academic Press. pp. 1457–1458. ISBN 0-12-352651-5. 
  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. 
  6. ^ John Daintith, ed. (2004). Oxford dictionary of chemistry (5th ed.). Oxford University Press. p. 235. ISBN 0-19-860918-3. 
  7. ^ 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. 
  8. ^ Amit Arora (2005). Text Book Of Inorganic Chemistry. Discovery Publishing House. pp. 691–692. ISBN 81-8356-013-X. 
  9. ^ 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. 
  10. ^ Stanley E. Manahan (2005). Environmental chemistry (8th ed.). CRC Press. p. 234. ISBN 1-56670-633-5.