Iron(II) fluoride

Iron(II) fluoride
Identifiers
7789-28-8 
13940-89-1 (tetrahydrate) 
ChemSpider 74215 Yes
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PubChem 522690
Properties
FeF2
Molar mass 93.84 g/mol (anhydrous)
165.902 g/mol (tetrahydrate)
Appearance red-violet transparent crystal
Density 4.09 g/cm3 (anhydrous)
2.20 g/cm3 (tetrahydrate)
Melting point 970 °C (1,780 °F; 1,240 K) (anhydrous)
100 °C (tetrahydrate)[1]
Boiling point 1,100 °C (2,010 °F; 1,370 K) (anhydrous)
165 g/100 mL
Solubility insoluble in ethanol, ether;
dissolves in HF
Structure
Crystal structure Rutile (tetragonal), tP6
Space group P42/mnm, No. 136
Hazards
Main hazards Causes severe skin burns & eye damage;
Hazardous decomposition products formed under fire conditions- Iron oxides[2]
Flash point not applicable[2]
Related compounds
Other anions
 iron(II) oxide, iron(II) chloride
Other cations
 manganese(II) fluoride, cobalt(II) fluoride
Related compounds
 iron(III) fluoride
Except where noted otherwise, data is given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
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Infobox references

Iron(II) fluoride (also ferrous fluoride) is the inorganic compound with formula FeF2. It is a green high-melting solid. The name iron(II) fluoride also applies to hydrates FeF2·4H2O.

Structure and physical properties

Like many difluorides, FeF2 adopts the TiO2 rutile structure wherein the iron centers are octahedrally and the fluoride ions three coordinate.[3][4] The Fe-F bond lengths are 2.03 and 2.10 Å.

Low temperature neutron diffraction studies show that the material is antiferromagnetic.[5] Heat capacity measurements reveal an event at 78.3 K corresponding to ordering of antiferromagnetic state.[6]

The vapor species were identified between 965 and 1149 K. Using mass spectrometry the heat of sublimation was experimentally determined and averaged to be 75.56 ± 0.23 kcal. mole−1.[7] The following reaction was proposed in order to calculate the atomization energy for Fe+:[7]

FeF2 + e → Fe+ + F2 (or 2F) + 2e

Synthesis and reactions

The anhydrous salt can be prepared by combining the elements.[4]

Recrystallization of the anhydrous solid from water[4] yields the colorless tetrahydrate, FeF2·4H2O, (CAS Number 13940-89-1). The latter can exists in two structures, or polymorphs. One form is rhombohedral and one is hexagonal, the former having a disordered[8] The tetrahydrate oxidizes in moist air to a hydrate of iron(III) fluoride, 2FeF3·9H2O.[8] The hydrate is poorly soluble in water, with a solubility product, Ksp, of 2.36 × 10−6.[9]

External links

References

  1. ↑ Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN 0-07-049439-8
  2. ↑ 2.0 2.1 Sigma-Aldrich. "Material Safety Data Sheet". Sigma-Aldrich. Retrieved 5 April 2011.
  3. ↑ Stout, J.; Stanley A. Reed (1954). "The Crystal Structure of MnF2, FeF2, CoF2, NiF2 and ZnF2". J. Am. Chem. Soc. 76: 5279–5281. doi:10.1021/ja01650a005.
  4. ↑ 4.0 4.1 4.2 Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0080379419.
  5. ↑ Erickson, R. (June 1953). "Neutron Diffraction Studies of Antiferromagnetism in Manganous Fluoride and Some Isomorphous Compounds". Physical Review 90 (5): 779–785. doi:10.1103/PhysRev.90.779.
  6. ↑ Stout, J.; Edward Catalano (December 1953). "Thermal Anomalies Associated with the Antiferromagnetic Ordering of FeF2, CoF3, and NiF2". Physical Review 92 (6): 1575–1575. doi:10.1103/PhysRev.92.1575.
  7. ↑ 7.0 7.1 Kent, Richard; John L. Margrave (November 1965). "Mass Spectrometric Studies at High Temperatures. VIII. The Sublimation Pressure of Iron(II) Fluoride". Journal of the American Chemical Society 87 (21): 4754–4756. doi:10.1021/ja00949a016.
  8. ↑ 8.0 8.1 Penfold, B. R.; Taylor, M. R. (1960). "The crystal structure of a disordered form of iron(II) fluoride tetrahydrate". Acta Crystallographica 13: 953–956. doi:10.1107/S0365110X60002302.
  9. ↑ Ksp solubility constant for common salts. Solubility of things site. Accessed on 2011-01-16.