Ferrate

Ferrate
IUPAC name

Ferrate(VI)
Systematic name

Dioxoironbis(olate)
Other names

Ferrate(2-)
Identifiers
PubChem 25000034
ChemSpider 21865127 Y
ChEBI CHEBI:30992 Y
Jmol-3D images Image 1
Image 2
Properties
Molecular formula FeO42-
Molar mass 119.843 g mol-1
Exact mass 119.914600621 g mol-1
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Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

In chemistry, ferrate(VI) refers either to the anion [FeO4]2−, in which iron is in the +6 oxidation state, or to a salt containing this anion. The term ferrate is often used to mean ferrate(VI), although according to IUPAC naming conventions, it may also refer to other iron-containing oxyanions, such as ferrate(V) and ferrate(IV).[1]

Contents

Synthesis

Ferrate(VI) salts are formed by oxidizing iron in 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]

Ferrates can also be produced by heating iron(III) hydroxide with sodium hypochlorite in alkaline solution:[3]

2 Fe(OH)3 + 3 OCl− + 4 OH− → 2 [FeO4]2− + 5 H2O + 3 Cl−

The yield may be increased by precipitating the resulting [FeO4]2− ion with Ba2+, forming barium ferrate.[3]

The ferrate(VI) ion has two unpaired electrons, and is thus paramagnetic. It has tetrahedral molecular geometry.[4]

Properties

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

FeO2−
4 + 3 e− + 8 H+ Fe3+ + 4 H2O

The reduction goes through intermediate species in which iron has oxidation states +5 and +4.[4] These have been found to be 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 Cr3+ to CrO2−
4,[8] and ammonia to elemental nitrogen.[9]

Ferrates are excellent disinfectants, and are capable of removing heavy metals, phosphates, and destroying viruses.[10]

References

  1. ^ Graham Hill; John Holman (2000). Chemistry in context (5th ed.). Nelson Thornes. p. 202. ISBN 0174482760. 
  2. ^ a b R. K. Sharma (2007). Text Book Of Coordination Chemistry. Discovery Publishing House. pp. 124–125. ISBN 818356223X. 
  3. ^ a b c Gary Wulfsberg (1991). Principles of descriptive inorganic chemistry. University Science Books. pp. 142–143. ISBN 0935702660. 
  4. ^ a b c Egon Wiberg; Nils Wiberg; Arnold Frederick Holleman (2001). Inorganic chemistry. Academic Press. pp. 1457–1458. ISBN 0123526515. 
  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 0849386799. 
  6. ^ John Daintith, ed (2004). Oxford dictionary of chemistry (5th ed.). Oxford University Press. p. 235. ISBN 0198609183. 
  7. ^ Kenneth Malcolm Mackay; Rosemary Ann Mackay; W. Henderson (2002). Introduction to modern inorganic chemistry (6th ed.). CRC Press. pp. 334–335. ISBN 0748764208. 
  8. ^ Amit Arora (2005). Text Book Of Inorganic Chemistry. Discovery Publishing House. pp. 691–692. ISBN 818356013X. 
  9. ^ Karlis Svanks (June 1976). "Oxidation of Ammonia in Water by Ferrates(VI) and (IV)" (PDF). Water Resources Center, Ohio State University. p. 3. https://kb.osu.edu/dspace/bitstream/1811/36335/1/OH_WRC_444.pdf. Retrieved 2010-05-04. 
  10. ^ Stanley E. Manahan (2005). Environmental chemistry (8th ed.). CRC Press. p. 234. ISBN 1566706335. 

See also