Cobalt(II) fluoride

Cobalt(II) fluoride
Names
	IUPAC name Cobalt(II) fluoride
	Other names cobalt difluoride
Identifiers
CAS Number	10026-17-2 ;
3D model (JSmol)	Interactive image;
ChemSpider	23205 ;
ECHA InfoCard	100.030.044
EC Number	233-061-9;
PubChem CID	24820;
RTECS number	GG0770000;
UNII	9KI67810UR ;
CompTox Dashboard (EPA)	DTXSID6064907 ;
	InChI InChI=1S/Co.2FH/h;21H/q+2;;/p-2 Key: YCYBZKSMUPTWEE-UHFFFAOYSA-L ; InChI=1/Co.2FH/h;21H/q+2;;/p-2 Key: YCYBZKSMUPTWEE-NUQVWONBAL;
	SMILES F[Co]F;
Properties
Chemical formula	CoF2
Molar mass	96.93 g/mol
Appearance	Red crystalline solid
Density	4.46 g/cm3 (anhydrous) ; 2.22 g/cm3 (tetrahydrate)
Melting point	1,217 °C (2,223 °F; 1,490 K)
Boiling point	1,400 °C (2,550 °F; 1,670 K)
Solubility in water	1.4 g/100 mL (25 °C)
Solubility	soluble in HF ; insoluble in alcohol, ether, benzene
Magnetic susceptibility (χ)	+9490.0·10−6 cm3/mol
Structure
Crystal structure	tetragonal (anhydrous) ; orthorhombic (tetrahydrate)
Hazards
NFPA 704 (fire diamond)	3 0 0
	Lethal dose or concentration (LD, LC):
LD50 (median dose)	oral (rat): 150 mg/kg
Related compounds
Other anions	cobalt(II) oxide, cobalt(II) chloride
Other cations	iron(II) fluoride, nickel(II) fluoride
Related compounds	cobalt trifluoride
	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

Cobalt(II) fluoride is a chemical compound with the formula (CoF₂). It is a pink crystalline solid compound^[1]^[2] which is antiferromagnetic at low temperatures (T_N=37.7 K)^[3] The formula is given for both the red tetragonal crystal, (CoF₂), and the tetrahydrate red orthogonal crystal, (CoF₂·4H₂O). CoF₂ is used in oxygen-sensitive fields, namely metal production. In low concentrations, it has public health uses. CoF₂ is sparingly soluble in water. The compound can be dissolved in warm mineral acid, and will decompose in boiling water. Yet the hydrate is water-soluble, especially the di-hydrate CoF₂·2H₂O and tri-hydrate CoF₂·3H₂O forms of the compound. The hydrate will also decompose with heat.

Like some other metal difluorides, CoF₂ crystallizes in the rutile structure, which features octahedral Co centers and planar fluorides.^[4]

Preparation

Cobalt(II) fluoride can be prepared from anhydrous cobalt(II) chloride or cobalt(II) oxide in a stream of hydrogen fluoride:

CoCl₂ + 2HF → CoF₂ + 2HCl

CoO + 2HF → CoF₂ + H₂O

It is produced in the reaction of cobalt (III) fluoride with water.

The tetrahydrate cobalt(II) fluoride is formed by dissolving cobalt(II) in hydrofluoric acid. The anhydrous fluoride can be extracted from this by dehydration. Other synthesis can occur at higher temperatures. It has been shown that at 500 °C fluorine will combine with cobalt producing a mixture of CoF₂ and CoF₃.^[5]

Uses

Cobalt(II) fluoride can be used as a catalyst to alloy metals. It is also used for optical deposition, of which it tremendously improves optical quality. Cobalt(II) fluoride is available in most volumes in an ultra high purity composition. High purity compositions improve optical qualities and its usefulness as a standard.

Analysis

To analyze this compound, Cobalt (II) fluoride can be dissolved in nitric acid. The solution is then diluted with water until appropriate concentration for AA or ICP spectrophotometry for the cobalt. A small amount of salt can be dissolved in cold water and analyzed for fluoride ion by a fluoride ion-selective electrode or ion chromatography.

Chemical properties

CoF₂ is a weak Lewis acid. Cobalt(II) complexes are usually octahedral or tetrahedral. As a 19-electron species it is a good reducing agent, fairly oxidizable into an 18-electron compound. Cobalt(II) fluoride can be reduced by hydrogen at a 300 °C.

References

^ Pradyot Patnaik (2002), Handbook of Inorganic Chemicals, McGraw-Hill Professional, ISBN 978-0-07-049439-8
^ Pashkevich, D. S.; Radchenko S. M.; Mukhortov, D. A., "Article title Heat Exchange between Cobalt(II) Fluoride Powder and the Wall of Rotating Cylinder" (PDF), Russian Journal of Applied Chemistry, Consultants Bureau, ISSN 1070-4272, archived from the original (PDF) on 2004-09-29, retrieved 2007-03-07
^ Ashcroft/Mermin: Solid State Physics (Tab. 33.2)
^ Stout, J. W.; Reed, Stanley A. (1954). "The Crystal Structure of MnF₂, FeF₂, CoF₂, NiF₂ and ZnF₂". J. Am. Chem. Soc. 76 (21): 5279–5281. doi:10.1021/ja01650a005.
^ J.C. Bailar (1973), Comprehensive Inorganic Chemistry, Pergoamon

External links

[1] Pradyot Patnaik (2002), Handbook of Inorganic Chemicals, McGraw-Hill Professional, ISBN 978-0-07-049439-8

[2] Pashkevich, D. S.; Radchenko S. M.; Mukhortov, D. A., "Article title Heat Exchange between Cobalt(II) Fluoride Powder and the Wall of Rotating Cylinder" (PDF), Russian Journal of Applied Chemistry, Consultants Bureau, ISSN 1070-4272, archived from the original (PDF) on 2004-09-29, retrieved 2007-03-07

[3] Ashcroft/Mermin: Solid State Physics (Tab. 33.2)

[4] Stout, J. W.; Reed, Stanley A. (1954). "The Crystal Structure of MnF₂, FeF₂, CoF₂, NiF₂ and ZnF₂". J. Am. Chem. Soc. 76 (21): 5279–5281. doi:10.1021/ja01650a005.

[5] J.C. Bailar (1973), Comprehensive Inorganic Chemistry, Pergoamon

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