Cobalt(II) fluoride

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Cobalt(II) fluoride
Fluorid kobaltnatý.PNG
Cobalt(II)-fluoride-unit-cell-3D-balls.png
Identifiers
CAS number 10026-17-2 YesY
PubChem 24820
ChemSpider 23205 YesY
EC number 233-061-9
RTECS number GG0770000
Jmol-3D images Image 1
Properties
Molecular 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
Structure
Crystal structure tetragonal (a,hydrous)
orthorhombic (tetrahydrate)
Hazards
NFPA 704
Flammability code 0: Will not burn. E.g., water Health code 3: Short exposure could cause serious temporary or residual injury. E.g., chlorine gas Reactivity code 2: Undergoes violent chemical change at elevated temperatures and pressures, reacts violently with water, or may form explosive mixtures with water. E.g., phosphorus Special hazards (white): no codeNFPA 704 four-colored diamond
LD50 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 noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
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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 (TN=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 CoF2·2H2 O and tri-hydrate CoF2·3H2O forms of the compound. The hydrate will also decompose with heat.

Preparation[edit]

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.[4]

Uses[edit]

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. The compound may be used in dental care[citation needed], since fluoride is also used in dental care.

Analysis[edit]

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[edit]

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[edit]

  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", Russian Journal of Applied Chemistry (Consultants Bureau), ISSN 1070-4272 
  3. ^ Ashcroft/Mermin: Solid State Physics (Tab. 33.2)
  4. ^ J.C. Bailar (1973), Comprehensive Inorganic Chemistry, Pergoamon 

External links[edit]