Iron(II) fluoride

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]

Iron(II) fluoride

Identifiers
CAS Number	7789-28-8 anhydrous Y 13940-89-1 (tetrahydrate) N
3D model (JSmol)	Interactive image
ChemSpider	74215 Y
ECHA InfoCard	100.029.232
PubChem CID	522690
UNII	NP4W87HLVO Y
CompTox Dashboard (EPA)	DTXSID50894779
InChI 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
SMILES [Fe+2].[F-].[F-]
Properties
Chemical formula	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)[2]
Boiling point	1,100 °C (2,010 °F; 1,370 K) (anhydrous)
Solubility product (Ksp)	2.36×10−6[3]
Solubility	insoluble in ethanol, ether; dissolves in HF
Magnetic susceptibility (χ)	+9500.0·10−6 cm3/mol
Structure
Crystal structure	Rutile (tetragonal), tP6
Space group	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:
Pictograms
NFPA 704 (fire diamond)	4 0 2 COR
Flash point	not applicable[4]
Related compounds
Other anions	Iron(II) chloride Iron(II) bromide Iron(II) iodide Iron(II) oxide
Other cations	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). N verify (what is YN ?) Infobox references

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 K_sp = 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

1. 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. Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN 0-07-049439-8
3. John Rumble (June 18, 2018). CRC Handbook of Chemistry and Physics (99 ed.). CRC Press. pp. 5–188. ISBN 978-1138561632.
4. Sigma-Aldrich. "Material Safety Data Sheet". Sigma-Aldrich. Retrieved 5 April 2011.
5. 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. ISBN 978-3527306732.

External links

• National Pollutant Inventory - Fluoride and compounds fact sheet