Iron(II) fluoride
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3D model (JSmol)
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| ChemSpider | |
| ECHA InfoCard | 100.029.232 |
PubChem CID
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| UNII | |
CompTox Dashboard (EPA)
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| 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) |
Solubility product (Ksp)
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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
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Causes severe skin burns & eye damage; Hazardous decomposition products formed under fire conditions- Iron oxides[4] |
| GHS labelling: | |
  
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| NFPA 704 (fire diamond) | |
| Flash point | not applicable[4] |
| Related compounds | |
Other anions
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Iron(II) chloride Iron(II) bromide Iron(II) iodide Iron(II) oxide |
Other cations
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Manganese(II) fluoride Cobalt(II) fluoride |
Related compounds
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Iron(III) fluoride |
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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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
Anhydrous FeF2 adopts the TiO2 rutile 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
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
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]
Uses
FeF2 is used to catalyze some organic reactions.[14]
References
- ^ 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.
- ^ John Rumble (June 18, 2018). CRC Handbook of Chemistry and Physics (99 ed.). CRC Press. pp. 5–188. ISBN 978-1138561632.
- ^ Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN 0-07-049439-8
- ^ a b Sigma-Aldrich. "Material Safety Data Sheet". Sigma-Aldrich. Retrieved 5 April 2011.
- ^ a b Dale L. Perry (1995), "Handbook of Inorganic Compounds", page 167. CRC Press. ISBN 9780849386718
- ^ 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.
- ^ 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.
- ^ 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.
- ^ 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.
- ^ 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.
- ^ 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.
- ^ 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.
- ^ "SOLUBILITY PRODUCT CONSTANTS" (PDF). Archived from the original (PDF) on 2018-07-12. Retrieved 2016-11-07.
- ^ 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. ISBN 978-3527306732.