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 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
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]
Battery research
FeF2 has been investigated as a cathode material for both lithium-ion and fluoride-ion batteries. Unlike conventional metal oxides, which rely on an intercalation-based lithium storage mechanism, FeFX (x = 2, 3) operates via a complex conversion mechanism, resulting in higher energy density. Fluoride cathodes are stable up to 1000°C.[15] Stability not only enhances safety and lowers the risk of thermal runaway.[16]
FeFX exhibits distinctive phase evolution, intermediate phases, and morphological transformations during lithiation and delithiation.[17][18] A stable lattice of fluoride anions is maintained throughout charge and discharge cycles, consistent with high cycling reversibility.[19][20]
^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. ISSN0108-7681.
^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. ISSN0021-9568.
^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.
