Iron(II) fluoride explained

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]

Structure and bonding

Anhydrous FeF₂ adopts the TiO₂ rutile structure. As such, the iron cations are octahedral and fluoride anions are trigonal planar.^{[2] [3]}

The tetrahydrate can exist in two structures, or polymorphs. One form is rhombohedral and the other is hexagonal, the former having a disorder.^[4]

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.^[5] Heat capacity measurements reveal an event at 78.3 K corresponding to ordering of antiferromagnetic state.^[6]

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.^[7]

The following reaction is proposed in order to calculate the atomization energy for Fe⁺:^[8]

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.^[9] It is slightly soluble in water (with solubility product K_sp = 2.36×10⁻⁶ at 25 °C)^[10] as well as dilute hydrofluoric acid, giving a pale green solution.^[4] It is insoluble in organic solvents.

The tetrahydrate can be prepared by dissolving iron in warm hydrated hydrofluoric acid and precipitating the result by addition of ethanol.^[4] It oxidizes in moist air to give, inter alia, a hydrate of iron(III) fluoride, (FeF₃)₂·9H₂O.

Uses

FeF₂ is used to catalyze some organic reactions.

External links

• National Pollutant Inventory - Fluoride and compounds fact sheet

Notes and References

1. Dale L. Perry (1995), "Handbook of Inorganic Compounds", page 167. CRC Press.
2. Stout. J.. Stanley A. Reed. The Crystal Structure of MnF₂, FeF₂, CoF₂, NiF₂ and ZnF₂. J. Am. Chem. Soc.. 1954. 76 . 21. 5279–5281. 10.1021/ja01650a005.
3. M.J.M.. de Almeida. M.M.R.. Costa. J.A.. Paixão. 1989-12-01. Charge density of FeF2. Acta Crystallographica Section B. en. 45. 6. 549–555. 10.1107/S0108768189007664. 0108-7681.
4. 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 . 10.1107/S0365110X60002302 . free .
5. Erickson. R. . Neutron Diffraction Studies of Antiferromagnetism in Manganous Fluoride and Some Isomorphous Compounds. Physical Review. June 1953. 90. 5. 779–785. 10.1103/PhysRev.90.779. 1953PhRv...90..779E .
6. Stout . J. . Edward Catalano . Thermal Anomalies Associated with the Antiferromagnetic Ordering of FeF₂, CoF₃, and NiF₂. Physical Review. December 1953. 92. 6. 1575. 10.1103/PhysRev.92.1575. 1953PhRv...92.1575S .
7. 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. 10.1021/je950115w. 0021-9568.
8. Kent. Richard. 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. 10.1021/ja00949a016. John L. Margrave.
9. 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.
10. Web site: SOLUBILITY PRODUCT CONSTANTS. 2016-11-07. 2018-07-12. https://web.archive.org/web/20180712140117/http://www4.ncsu.edu/~franzen/public_html/CH201/data/Solubility_Product_Constants.pdf. dead.