Cadmium fluoride (CdF2) is a mostly water-insoluble source of cadmium used in oxygen-sensitive applications, such as the production of metallic alloys. In extremely low concentrations (ppm), this and other fluoride compounds are used in limited medical treatment protocols. Fluoride compounds also have significant uses in synthetic organic chemistry.[3] The standard enthalpy has been found to be -167.39 kcal. mole−1 and the Gibbs energy of formation has been found to be -155.4 kcal. mole−1, and the heat of sublimation was determined to be 76 kcal. mole−1.[4][5]
Cadmium fluoride has also been prepared by reacting fluorine with cadmium sulfide. This reaction happens very quickly and forms nearly pure fluoride at much lower temperatures than other reactions used.[7]
CdF2 can be transformed into an electronic conductor when doped with certain rare earth elements or yttrium and treated with cadmium vapor under high temperature conditions. This process creates blue crystals with varying absorption coefficients depending on the concentrations of the dopant. A proposed mechanism explains that the conductivity of these crystals can be explained by a reaction of Cd atoms with InterstitialF− ions. This creates more CdF2 molecules and releases electrons which are weakly bonded to trivalent dopant ions resulting in n-type conductivity and a hydrogenic donor level.[8]
Cadmium fluoride, like all cadmium compounds, is toxic and should be used with care.
Cadmium fluoride can cause potential health issues if it is not handled properly. It can cause irritation to the skin and the eyes, so gloves and protective eyewear are advised. The MSDS, or Material Safety Data Sheet, also includes warnings for ingestion and inhalation. Under acidic conditions, at high temperatures, and in moist environments, hydrogen fluoride and cadmium vapors may be released into the air. Inhalation may cause irritation of the respiratory system as well as congestion, fluorosis, and even pulmonary edema in extreme cases. Cadmium fluoride also has the same potential hazards caused by cadmium and fluoride.[9]
^Rudzitis, Edgars; Feder, Harold; Hubbard, Ward (November 1963). "Fluorine Bomb Calorimetry. VII. The Heat of Formation of Cadmium Difluoride". Journal of Physical Chemistry. 67 (11): 2388–2390. doi:10.1021/j100805a031.
^Besenbruch, G.; Kana'an, A. S.; Margrave, J. L. (March 3, 1965). "Knudson and Langmuir Measurements of the Sublimation Pressure of Cadmium (II) Fluoride". Journal of Physical Chemistry. 69 (9): 3174–3176. doi:10.1021/j100893a505.
^Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN0-07-049439-8
^Haendler, Helmut; Bernard, Walter (November 1951). "The Reaction of Fluorine with Cadmium and Some of its Binary Compounds. The Crystal Structure, Density and Melting Points of Cadmium Fluoride". Journal of the American Chemical Society. 73. doi:10.1021/ja01155a064.
^Weller, Paul (June 1, 1965). "Electrical and Optical Properties of Rare Earth Doped Cadmium Fluoride Single Crystals". Inorganic Chemistry. 4 (11): 1545–1551. doi:10.1021/ic50033a004.
