Copper(II) fluoride

Copper(II) fluoride is an inorganic compound with the chemical formula CuF2. It is a white crystalline, hygroscopic solid with a rutile-type crystal structure, similar to other fluorides of chemical formulae MF2 (where M is a metal).

Copper(II) fluoride
Names
IUPAC name
Copper difluoride
Other names
Cupric fluoride; Copper fluoride; Copper (2+) Difluoride
Identifiers
3D model (JSmol)
ChemSpider
ECHA InfoCard 100.029.225
UNII
Properties
CuF2
Molar mass 101.543 g/mol (anhydrous)
137.573 g/mol (dihydrate)
Appearance White crystalline powder
When hydrated: Blue
Density 4.23 g/cm3 (anhydrous)
2.934 g/cm3 (dihydrate)[1]
Melting point 836 °C (1,537 °F; 1,109 K) (anhydrous)
130 °C (dihydrate, decomposes)
Boiling point 1,676 °C (3,049 °F; 1,949 K) (anhydrous)
Solubility in other solvents Hygroscopic
+1050.0·10−6 cm3/mol
Hazards
NIOSH (US health exposure limits):
PEL (Permissible)
 TWA 1 mg/m3 (as Cu)[2]
REL (Recommended)
 TWA 1 mg/m3 (as Cu)[2]
IDLH (Immediate danger)
 TWA 100 mg/m3 (as Cu)[2]
Related compounds
Other anions
 Copper(II) bromide
Copper(II) chloride
Other cations
 Silver(II) fluoride
Cobalt(II) fluoride
Related compounds
 Copper(I) fluoride
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Infobox references

Structure

Copper(II) fluoride has a monoclinic crystal structure and cannot achieve a higher-symmetry structure. It forms rectangular prisms with a parallelogram base.

Uses

Copper (II) fluoride can be used to make fluorinated aromatic hydrocarbons by reacting with aromatic hydrocarbons in an oxygen-containing atmosphere at temperatures above 450 °C (842 °F). This reaction is simpler than the Sandmeyer reaction, but is only effective in making compounds that can survive at the temperature used. A coupled reaction using oxygen and 2 HF regenerates the copper(II) fluoride, producing water.[3] This method has been proposed as a "greener" method of producing fluoroaromatics since it avoids producing toxic waste products such as NaF and NH4F salts.

Chemistry

Copper(II) fluoride can be synthesized from copper and fluorine at temperatures of 400 °C (752 °F). It occurs as a direct reaction.

Cu + F2 → CuF2

It loses fluorine in the molten stage at temperatures above 950 °C (1742 °F).

2CuF2 → 2CuF + F2
2CuF → CuF2 + Cu

The complex anions of CuF3−, CuF42− and CuF64− are formed if CuF2 is exposed to substances containing fluoride ions F−.

Solubility

Copper(II) fluoride is slightly soluble in water, but starts to decompose when it is in hot water, producing basic F− and Cu(OH) ions. [reference needed]

Toxicity

There is little specific information on the toxicity of Copper(II) fluoride. However, copper and fluoride can both be toxic individually when consumed.

Copper toxicity can affect the skin, eyes, and respiratory tract. Serious conditions include metal fume fever, and hemolysis of red blood cells. Copper can also cause damage to the liver and other major organs.

Fluoride is safe at low levels and is added to water in many countries to protect against tooth decay. At higher levels it can cause toxic effects ranging from nausea and vomiting to tremors, breathing problems, serious convulsions and even coma. Brain and kidney damage can result. Chronic exposure can cause losses in bone density, weight loss and anorexia.

Hazards

Experiments using copper(II) fluoride should be conducted in a fume hood because metal oxide fumes can occur. The combination of acids with copper(II) fluoride may lead to the production of hydrogen fluoride, which is highly toxic and corrosive.

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References
  1. Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN 0-07-049439-8
  2. NIOSH Pocket Guide to Chemical Hazards. "#0150". National Institute for Occupational Safety and Health (NIOSH).
  3. M. A. Subramanian; L. E. Manzer (2002). "A "Greener" Synthetic Route for Fluoroaromatics via Copper (II) Fluoride". Science. 297 (5587): 1665. doi:10.1126/science.1076397. PMID 12215637.
  • C. Billy; H. M. Haendler (1957). "The Crystal Structure of Copper(II) Fluoride". Journal of the American Chemical Society. 79 (5): 1049–51. doi:10.1021/ja01562a011.
  • P. C. de Mello; M. Hehenberg; S. Larson; M. Zerner (1980). "Studies of the electronic structure of copper fluorides and copper chlorides". Journal of the American Chemical Society. 102 (4): 1278–1288. doi:10.1021/ja00524a010.
  • H. M. Haendler; L. H. Towle; E. F. Bennett; W. L. Patterson (1954). "The Reaction of Fluorine with Copper and Some of its Compounds. Some Properties of Copper(II) Fluoride". Journal of the American Chemical Society. 76 (8): 2178–2179. doi:10.1021/ja01637a039.
  • T. C. Ehlert; J. S. Wang (1977). "Thermochemistry of the copper fluorides". Journal of Physical Chemistry. 81 (22): 2069–2073. doi:10.1021/j100537a005.
  • Dierks, S. "Copper Fluoride". http://www.espimetals.com/index.php/msds/537-copper-fluoride (accessed October 9).
  • Subramanian, M. A.; Manzer, L. E. (2002). "A 'Greener' Synthetic Route for Fluoroaromantics via Copper (II) Fluoride". Science. 297 (5587): 1665. doi:10.1126/science.1076397. PMID 12215637.
  • Olejniczak, I.; Wolak, J.; Barszcz, B.; Schlueter, J.; Manson, J. (2010). "CuF2 Structural Changes in Two-Dimensional Quantum Magnet (H2O)2(pyz) Under Pressure: Raman Study". AIP Conference Proceedings. 1267 (1): 597–598. doi:10.1063/1.3482697.
  • Kent, R. A.; Mcdonald, J. D.; Margrave, J. L. (1966). "Mass Spectrometric Studies at High Temperatures. IX. The Sublimation Pressure of Copper(II) Fluoride". Journal of Physical Chemistry. 70 (3): 874–877. doi:10.1021/j100875a042.
  • Shashkin, S. Y.; Goddard III, W. A. (1986). "Electron Correlation effects in ligand field parameters and other properties of copper II fluoride". Journal of Physical Chemistry. 90 (2): 250–255. doi:10.1021/j100274a010.
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