Iron(II) fluoride

In this article, we will delve into the fascinating world of Iron(II) fluoride. We will explore its origins, its impact on current society and possible future developments related to this topic. From its inception to the present, Iron(II) fluoride has aroused great interest in various areas, attracting both experts and people interested in better understanding its relevance. Throughout these pages, we will analyze its many facets and how it has influenced people's lives. Without a doubt, Iron(II) fluoride is a topic that will not leave anyone indifferent, and we are sure that this article will be of great interest to all those who wish to learn more about it.

Iron(II) fluoride
Identifiers
3D model (JSmol)
ChemSpider
ECHA InfoCard 100.029.232 Edit this at Wikidata
UNII
  • InChI=1S/2FH.Fe/h2*1H;/q;;+2/p-2 checkY
    Key: FZGIHSNZYGFUGM-UHFFFAOYSA-L checkY
  • 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:
GHS05: CorrosiveGHS08: Health hazardGHS09: Environmental hazard
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 4: Very short exposure could cause death or major residual injury. E.g. VX gasFlammability 0: Will not burn. E.g. waterInstability 2: Undergoes violent chemical change at elevated temperatures and pressures, reacts violently with water, or may form explosive mixtures with water. E.g. white phosphorusSpecial hazard COR: Corrosive; strong acid or base. E.g. sulfuric acid, potassium hydroxide
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 , 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]

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]

References

  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. ISBN 978-3-527-30673-2.
  15. ^ Nikitin, M. I.; Chilingarov, N. S.; Alikhanyan, A. S. (2021-01-01). "Thermal Stability of Mixed Fluorides of 3d Elements". Russian Journal of Inorganic Chemistry. 66 (1): 89–95. doi:10.1134/S0036023621010058. ISSN 1531-8613.
  16. ^ Viswanathan, Venkatasubramanian; Epstein, Alan H.; Chiang, Yet-Ming; Takeuchi, Esther; Bradley, Marty; Langford, John; Winter, Michael (March 2022). "Author Correction: The challenges and opportunities of battery-powered flight". Nature. 603 (7903): E30. doi:10.1038/s41586-022-04612-5. ISSN 1476-4687. PMID 35293393.
  17. ^ Badway, F.; Cosandey, F.; Pereira, N.; Amatucci, G. G. (2003). "Carbon Metal Fluoride Nanocomposites". Journal of the Electrochemical Society. 150 (10): A1318. doi:10.1149/1.1602454.
  18. ^ Olbrich, Lorenz F.; Xiao, Albert W.; Pasta, Mauro (2021-12-01). "Conversion-type fluoride cathodes: Current state of the art". Current Opinion in Electrochemistry. 30: 100779. doi:10.1016/j.coelec.2021.100779. ISSN 2451-9103.
  19. ^ Xiao, Albert W.; Lee, Hyeon Jeong; Capone, Isaac; Robertson, Alex; Wi, Tae-Ung; Fawdon, Jack; Wheeler, Samuel; Lee, Hyun-Wook; Grobert, Nicole; Pasta, Mauro (June 2020). "Understanding the conversion mechanism and performance of monodisperse FeF2 nanocrystal cathodes". Nature Materials. 19 (6): 644–654. doi:10.1038/s41563-020-0621-z. ISSN 1476-4660. PMID 32094491.
  20. ^ Olbrich, Lorenz F.; Xiao, Albert W.; Schart, Maximilian; Ihli, Johannes; Matthews, Guillaume; Sanghadasa, Mohan; Pasta, Mauro (February 2024). "Iron fluoride-lithium metal batteries in bis(fluorosulfonyl)imide-based ionic liquid electrolytes". Cell Reports Physical Science. 5 (2): 101787. doi:10.1016/j.xcrp.2024.101787.