Iron(II) chloride
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| Names | |
|---|---|
| IUPAC names
Iron(II) chloride Iron dichloride | |
| Other names
Ferrous chloride, Rokühnite | |
| Identifiers | |
7758-94-3  16399-77-2 (dihydrate)  13478-10-9 (tetrahydrate) | |
| 3D model (Jmol) | Interactive image |
| ChEBI | CHEBI:30812 |
| ChemSpider | 22866 |
| ECHA InfoCard | 100.028.949 |
| EC Number | 231-843-4 |
| PubChem | 24458 |
| RTECS number | NO5400000 |
| UNII | S3Y25PHP1W |
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| Properties | |
| FeCl2 | |
| Molar mass | 126.751 g/mol (anhydrous) 198.8102 g/mol (tetrahydrate) |
| Appearance | tan solid (anhydrous) pale green solid (di-tetrahydrate) |
| Density | 3.16 g/cm3 (anhydrous) 2.39 g/cm3 (dihydrate) 1.93 g/cm3 (tetrahydrate) |
| Melting point | 677 °C (1,251 °F; 950 K) (anhydrous) 120 °C (dihydrate) 105 °C (tetrahydrate) [1] |
| Boiling point | 1,023 °C (1,873 °F; 1,296 K) (anhydrous) |
| 64.4 g/100 mL (10 °C), 68.5 g/100mL (20 °C), 105.7 g/100 mL (100 °C) | |
| Solubility in THF | soluble |
| Solubility in ethanol | 100 g/100 mL (value should be double checked , experiment shows merely soluble(anhydrous)) |
| log P | -0.15 |
| Structure | |
| Monoclinic | |
| octahedral at Fe | |
| Pharmacology | |
| B03AA05 (WHO) | |
| Hazards | |
| NFPA 704 | |
| US health exposure limits (NIOSH): | |
| REL (Recommended) |
TWA 1 mg/m3[2] |
| Related compounds | |
| Other anions |
Iron(II) fluoride Iron(II) bromide Iron(II) iodide |
| Other cations |
Cobalt(II) chloride Manganese(II) chloride Copper(II) chloride |
| Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa). | |
| verify (what is ?) | |
| Infobox references | |
Iron(II) chloride, also known as ferrous chloride, is the chemical compound of formula FeCl2. It is a paramagnetic solid with a high melting point. The compound is white, but typical samples are often off-white. FeCl2 crystallizes from water as the greenish tetrahydrate, which is the form that is most commonly encountered in commerce and the laboratory. There is also a dihydrate. The compound is also soluble in water; aqueous solutions of FeCl2 are highly transparent and pale green in color.
Production
Hydrated forms of ferrous chloride are generated by treatment of wastes from steel production with hydrochloric acid. Such solutions are designated "spent acid," especially when the hydrochloric acid is not completely consumed:
The spent acid requires treatment before its disposal. It is also byproduct from titanium production, since some titanium ores contain iron.[3]
Laboratory preparation
The dihydrate crystallizes from concentrated hydrochloric acid.[4]
Anhydrous FeCl2
Ferrous chloride is conveniently prepared by addition of iron powder to a solution of methanol and concentrated hydrochloric acid under an inert atmosphere. This reaction gives the methanol solvate, which upon heating in a vacuum at about 160 °C gives anhydrous FeCl2.[5] FeBr2 and FeI2 can be prepared analogously.
- Fe + 2 HCl → FeCl2 + H2
An alternative synthesis of entails the reduction of FeCl3 with chlorobenzene:[6]
- 2 FeCl3 + C6H5Cl → 2 FeCl2 + C6H4Cl2 + HCl
FeCl2 is solubility in tetrahydrofuran (THF). In one of two classic syntheses of ferrocene, Wilkinson generated FeCl2 by heating FeCl3 with iron powder in THF.[7] Ferric chloride decomposes to ferrous chloride at high temperatures.
Reactions
FeCl2 forms complexes with many ligands. It reacts with two molar equivalents of [(C2H5)4N]Cl to give the salt [(C2H5)4N]2[FeCl4]. Related compounds that can be prepared similarly include the [MnCl4]2−, [MnBr4]2−, [MnI4]2−, [FeBr4]2−, [CoCl4]2−, [CoBr4]2−, [NiCl4]2−, and [CuCl4]2− salts.[8]
Applications
Ferrous chloride has a variety of niche applications, but the related compounds ferrous sulfate and ferric chloride enjoy more applications. Aside from use in the laboratory synthesis of iron complexes, ferrous chloride serves as a reducing flocculating agent in wastewater treatment, especially for wastes containing chromate. It is the precursor to hydrated iron(III) oxides that are magnetic pigments.[3] Ferrous chloride is employed as a reducing agent in many organic synthesis reactions.
References
- ↑ Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN 0-07-049439-8
- ↑ "NIOSH Pocket Guide to Chemical Hazards #0346". National Institute for Occupational Safety and Health (NIOSH).
- 1 2 Egon Wildermuth, Hans Stark, Gabriele Friedrich, Franz Ludwig Ebenhöch, Brigitte Kühborth, Jack Silver, Rafael Rituper “Iron Compounds” in Ullmann’s Encyclopedia of Industrial Chemistry Wiley-VCH, Wienheim, 2005.
- ↑ K. H.. Gayer; L. Woontner (1957). "Iron(II) Chloride 2-Hydrate". Inorg. Synth. 5: 179–181. doi:10.1002/9780470132364.ch48.
- ↑ G. Winter; Thompson, D. W.; Loehe, J. R. (1973). "Iron(II) Halides". Inorg. Synth. 14: 99–104. doi:10.1002/9780470132456.ch20.
- ↑ P. Kovacic and N. O. Brace (1960). "Iron(II) Chloride". Inorg. Synth. 6: 172. doi:10.1002/9780470132371.ch54.
- ↑ G. Wilkinson (1963). "Ferrocene". Org. Synth.; Coll. Vol., 4, p. 473
- ↑ N. S. Gill, F. B. Taylor (1967). "Tetrahalo Complexes of Dipositive Metals in the First Transition Series". Inorg. Synth. 9: 136–142. doi:10.1002/9780470132401.ch37.