From Wikipedia, the free encyclopedia
  (Redirected from Ferrate)
Jump to: navigation, search
Aromatic skeletal formula of ferrate
Ferrate and permanganate solution.jpg
Solutions of ferrate (left)
and permanganate (right)
PubChem 25000034
ChemSpider 21865127 YesY
ChEBI CHEBI:30992 YesY
Jmol-3D images Image 1
Image 2
Molecular formula FeO42-
Molar mass 119.843 g mol-1
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.


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]


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 OCl
+ 4 OH
→ 2 [FeO
+ 5 H
+ 3 Cl

The anion is typically precipitated as the barium(II) salt, forming barium ferrate.[3]


The ferrate(VI) anion is unstable at neutral[2] or acidic pH values, decomposing to iron(III):[3]

+ 3 -
+ 8 H+
is in equilibrium with Fe3+
+ 4 H

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]


  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.