Calcium fluoride

Last updated on 20 September 2017

Calcium fluoride is the inorganic compound with the formula CaF2. It is a white insoluble solid. It occurs as the mineral fluorite (also called fluorspar), which is often deeply coloured owing to impurities.

Calcium fluoride
Fluorite-unit-cell-3D-ionic.png
Calcium fluoride.jpg
Fluorid v%C3%A1penat%C3%BD.PNG
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.029.262
EC Number 232-188-7
RTECS number EW1760000
UNII
Properties
CaF2
Molar mass 78.07 g·mol−1
Appearance White crystalline solid (single crystals are transparent)
Density 3.18 g/cm3
Melting point 1,418 °C (2,584 °F; 1,691 K)
Boiling point 2,533 °C (4,591 °F; 2,806 K)
0.0015 g/100 mL (18 °C)
0.0016 g/100 mL (20 °C)
3.9 × 10−11 [1]
Solubility insoluble in acetone
slightly soluble in acid
-28.0·10−6 cm3/mol
1.4338
Structure
cubic crystal system, cF12[2]
Fm3m, #225
Ca, 8, cubic
F, 4, tetrahedral
Hazards
Main hazards Reacts with conc. sulfuric acid to produce hydrofluoric acid
Safety data sheet ICSC 1323
R-phrases (outdated) R20, R22, R36, R37, R38
S-phrases (outdated) S26, S36
NFPA 704
Flammability code 0: Will not burn. E.g., water Health code 0: Exposure under fire conditions would offer no hazard beyond that of ordinary combustible material. E.g., sodium chloride Reactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g., liquid nitrogen Special hazards (white): no codeNFPA 704 four-colored diamond
Flash point Non-flammable
Lethal dose or concentration (LD, LC):
>5000 mg/kg (oral, guinea pig)
4250 mg/kg (oral, rat)[3]
Related compounds
Other anions
Calcium chloride
Calcium bromide
Calcium iodide
Other cations
Beryllium fluoride
Magnesium fluoride
Strontium fluoride
Barium fluoride
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

Chemical structure

The compound crystallizes in a cubic motif called a fluorite structure.

Xtals combined 2 300ppi.png
Unit cell of CaF2, known as fluorite structure, from two equivalent perspectives. The second origin is often used when visualising point defects entered on the cation.[4]

Ca2+ centres are eight-coordinate, being centered in a "box" for eight F centres. Each F centre is coordinated to four Ca2+ centres.[5] Although perfectly packed crystalline samples are colorless, the mineral is often deeply colored due to the presence of F-centers. The same crystal structure is found in numerous ionic compounds with formula AB2, such as CeO2, cubic ZrO2, UO2, ThO2 and PuO2. A related structure is the antifluorite structure, where the anions and cations are swapped, such as Be2C

Preparation

The mineral fluorite is abundant, widespread, and mainly of interest as a precursor to HF. Thus, little motivation exists for the industrial production of CaF2. High purity CaF2 is produced by treating calcium carbonate with hydrofluoric acid:[6]

CaCO3 + 2 HF → CaF2 + CO2 + H2O

Applications

Naturally occurring CaF2 is the principal source of hydrogen fluoride, a commodity chemical used to produce a wide range of materials. Calcium fluoride in the fluorite state is of significant commercial importance as a fluoride source.[7] Hydrogen fluoride is liberated from the mineral by the action of concentrated sulfuric acid:[8]

CaF2 + H2SO4CaSO4(solid) + 2 HF

Niche uses

Calcium fluoride is used to manufacture optical components such as windows and lenses, used in thermal imaging systems, spectroscopy, telescopes and excimer lasers. It is transparent over a broad range from ultraviolet (UV) to infrared (IR) frequencies. Its low refractive index reduces the need for anti-reflection coatings. Its insolubility in water is convenient as well.

Safety

CaF2 is classified as "not dangerous", although reacting it with sulfuric acid produces toxic hydrofluoric acid. With regards to inhalation, the NIOSH-recommended concentration of fluorine-containing dusts is 2.5 mg/m3 in air.[6]

See also

Related materials

References

  1. ^ Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN 0-07-049439-8
  2. ^ X-ray Diffraction Investigations of CaF2 at High Pressure, L. Gerward, J. S. Olsen, S. Steenstrup, M. Malinowski, S. Åsbrink and A. Waskowska, Journal of Applied Crystallography (1992), 25, 578-581 doi:10.1107/S0021889892004096
  3. ^ "Fluorides (as F)". Immediately Dangerous to Life and Health. National Institute for Occupational Safety and Health (NIOSH).
  4. ^ Burr, P. A.; Cooper, M. W. D. (2017-09-15). "Importance of elastic finite-size effects: Neutral defects in ionic compounds". Physical Review B. 96 (9): 094107. doi:10.1103/PhysRevB.96.094107.
  5. ^ G. L. Miessler and D. A. Tarr “Inorganic Chemistry” 3rd Ed, Pearson/Prentice Hall publisher, ISBN 0-13-035471-6.
  6. ^ a b Aigueperse, Jean; Mollard, Paul; Devilliers, Didier; Chemla, Marius; Faron, Robert; Romano, René; Cuer, Jean Pierre (2000). "Fluorine Compounds, Inorganic". doi:10.1002/14356007.a11_307.
  7. ^ Aigueperse, Jean; Mollard, Paul; Devilliers, Didier; Chemla, Marius; Faron, Robert; Romano, Renée; Cuer, Jean Pierre (2005), "Fluorine Compounds, Inorganic", Ullmann's Encyclopedia of Industrial Chemistry, Weinheim: Wiley-VCH, p. 307, doi:10.1002/14356007.a11_307.
  8. ^ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.

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