Iron(II) fluoride – Wikipedia

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Iron(II) fluoride
Identifiers
ChemSpider
ECHA InfoCard 100.029.232
UNII
  • InChI=1S/2FH.Fe/h2*1H;/q;;+2/p-2 Y
    Key: FZGIHSNZYGFUGM-UHFFFAOYSA-L Y
  • InChI=1/2FH.Fe/h2*1H;/q;;+2/p-2

    Key: FZGIHSNZYGFUGM-NUQVWONBAX

Properties
FeF2
Molar mass 93.84 g/mol (anhydrous)
165.902 g/mol (tetrahydrate)
Appearance colorless transparent crystals[1]
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)[3]
Boiling point 1,100 °C (2,010 °F; 1,370 K) (anhydrous)
2.36×10−6[2]
Solubility insoluble in ethanol, ether;
dissolves in HF
+9500.0·10−6 cm3/mol
Structure
Rutile (tetragonal), tP6
P42/mnm, No. 136
Hazards
Occupational safety and health (OHS/OSH):

Main hazards

Causes severe skin burns & eye damage;
Hazardous decomposition products formed under fire conditions- Iron oxides[4]
GHS labelling:
NFPA 704 (fire diamond)
Flash point not applicable[4]
Related compounds
Iron(II) chloride
Iron(II) bromide
Iron(II) iodide
Iron(II) oxide
Manganese(II) fluoride
Cobalt(II) fluoride

Related compounds

Iron(III) fluoride
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

Chemical compound

Iron(II) fluoride or ferrous fluoride is an inorganic compound with the molecular formula FeF2. It forms a tetrahydrate FeF2·4H2O that is often referred to by the same names. The anhydrous and hydrated forms are white crystalline solids.[1][5]

Structure and bonding[edit]

Anhydrous FeF2 adopts the TiO2rutile structure. As such, the iron cations are octahedral and fluoride anions are trigonal planar.[6][7]

The tetrahydrate can exist in two structures, or polymorphs. One form is rhombohedral and the other is hexagonal, the former having a disorder.[1]

Like most fluoride compounds, the anhydrous and hydrated forms of iron(II) fluoride feature high spin metal center. Low temperature neutron diffraction studies show that the FeF2 is antiferromagnetic.[8]Heat capacity measurements reveal an event at 78.3 K corresponding to ordering of antiferromagnetic state.[9]

Selected physical properties[edit]

FeF2 sublimes between 958 and 1178 K. Using Torsion and Knudsen methods, the heat of sublimation was experimentally determined and averaged to be 271 ± 2 kJ mole−1.[10]

The following reaction is proposed in order to calculate the atomization energy for Fe+:[11]

FeF2 + e → Fe+ + F2 (or 2F) + 2e

Synthesis and reactions[edit]

The anhydrous salt can be prepared by reaction of ferrous chloride with anhydrous hydrogen fluoride.[12] It is slightly soluble in water (with solubility product Ksp = 2.36×10−6 at 25 °C)[13] as well as dilute hydrofluoric acid, giving a pale green solution.[1] It is insoluble in organic solvents.[5]

The tetrahydrate can be prepared by dissolving iron in warm hydrated hydrofluoric acid and precipitating the result by addition of ethanol.[1] It oxidizes in moist air to give, inter alia, a hydrate of iron(III) fluoride, (FeF3)2·9H2O.[1]

FeF2 is used to catalyze some organic reactions.[14]

References[edit]

  1. ^ a b c d e f Penfold, B. R.; Taylor, M. R. (1960). “The crystal structure of a disordered form of iron(II) fluoride tetrahydrate”. Acta Crystallographica. 13 (11): 953–956. doi:10.1107/S0365110X60002302.
  2. ^ John Rumble (June 18, 2018). CRC Handbook of Chemistry and Physics (99 ed.). CRC Press. pp. 5–188. ISBN 978-1138561632.
  3. ^ Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN 0-07-049439-8
  4. ^ a b Sigma-Aldrich. “Material Safety Data Sheet”. Sigma-Aldrich. Retrieved 5 April 2011.
  5. ^ a b Dale L. Perry (1995), “Handbook of Inorganic Compounds“, page 167. CRC Press. ISBN 9780849386718
  6. ^ Stout, J.; Stanley A. Reed (1954). “The Crystal Structure of MnF2, FeF2, CoF2, NiF2 and ZnF2“. J. Am. Chem. Soc. 76 (21): 5279–5281. doi:10.1021/ja01650a005.
  7. ^ M.J.M., de Almeida; M.M.R., Costa; J.A., Paixão (1989-12-01). “Charge density of FeF2”. Acta Crystallographica Section B. 45 (6): 549–555. doi:10.1107/S0108768189007664. ISSN 0108-7681.
  8. ^ Erickson, R. (June 1953). “Neutron Diffraction Studies of Antiferromagnetism in Manganous Fluoride and Some Isomorphous Compounds”. Physical Review. 90 (5): 779–785. Bibcode:1953PhRv…90..779E. doi:10.1103/PhysRev.90.779.
  9. ^ Stout, J.; Edward Catalano (December 1953). “Thermal Anomalies Associated with the Antiferromagnetic Ordering of FeF2, CoF3, and NiF2“. Physical Review. 92 (6): 1575. Bibcode:1953PhRv…92.1575S. doi:10.1103/PhysRev.92.1575.
  10. ^ Bardi, Gianpiero; Brunetti, Bruno; Piacente, Vincenzo (1996-01-01). “Vapor Pressure and Standard Enthalpies of Sublimation of Iron Difluoride, Iron Dichloride, and Iron Dibromide”. Journal of Chemical & Engineering Data. 41 (1): 14–20. doi:10.1021/je950115w. ISSN 0021-9568.
  11. ^ 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.
  12. ^ W. Kwasnik “Iron(II) Fluoride” in Handbook of Preparative Inorganic Chemistry, 2nd Ed. Edited by G. Brauer, Academic Press, 1963, NY. Vol. 1. p. 266.
  13. ^ “SOLUBILITY PRODUCT CONSTANTS” (PDF). Archived from the original (PDF) on 2018-07-12. Retrieved 2016-11-07.
  14. ^ Wildermuth, Egon; Stark, Hans; Friedrich, Gabriele; Ebenhöch, Franz Ludwig; Kühborth, Brigitte; Silver, Jack; Rituper, Rafael (2000). “Iron Compounds”. Ullmann’s Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a14_591.

External links[edit]