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Calcium fluoride
Calcium fluoride.jpg
Fluorid vápenatý.PNG
CAS number 7789-75-5 Yes check.svgY
PubChem 24617
EC number 232-188-7
RTECS number EW1760000
ChemSpider ID 23019
Molecular formula CaF2
Molar mass 78.07 g mol−1
Appearance White crystalline solid (single crystals are transparent)
Density 3.18 g/cm3
Melting point

1418 °C, 1691 K, 2584 °F

Boiling point

2533 °C, 2806 K, 4591 °F

Solubility in water 0.0015 g/100 mL (18 °C)
0.0016 g/100 mL (20 °C)
Solubility product, Ksp 3.9 x 10-11 [1]
Solubility in acetone insoluble
Refractive index (nD) 1.4328
Crystal structure cubic crystal system, cF12[2]
Space group Fm3m, #225
Ca, 8, cubic
F, 4, tetrahedral
EU Index Not listed
Main hazards Reacts with conc. sulfuric acid to produce hydrofluoric acid
NFPA 704
NFPA 704.svg
Flash point Non-flammable
LD50 4250 mg/kg (oral, rat)
Related compounds
Other anions Calcium chloride
Calcium bromide
Calcium iodide
Other cations Magnesium fluoride
Strontium fluoride
Barium fluoride
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Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Calcium fluoride is the inorganic compound with the formula CaF2. This ionic compound of calcium and fluorine occurs naturally as the mineral fluorite (also called fluorspar). It is the source of most of the world's fluorine. This insoluble solid adopts a cubic structure wherein calcium is coordinated to eight fluoride anions and each F ion is surrounded by four Ca2+ ions.[3] Although the pure material is colourless, the mineral is often deeply coloured due to the presence of F-centers.



The mineral fluorspar is abundant, widespread, and mainly of interest as a precursor to HF. Thus, little motivation exists for the preparation of CaF2. High purity CaF2 produced by treating calcium carbonate with hydrofluoric acid:[4]

CaCO3 + 2 HF → CaF2 + CO2 + H2O

Source of HF

Naturally occurring CaF2 is the principal source of hydrogen fluoride, a commodity chemical used to produce a wide range of materials. Fluoride is liberated from the mineral by the action of concentrated sulfuric acid:

CaF2(s) + H2SO4CaSO4(solid) + 2 HF(g)

The resulting HF is converted into fluorine, fluorocarbons, and diverse fluoride materials. As of the late 1990s, five billion kilograms were mined annually.[5]


In the laboratory, calcium fluoride is commonly used as a window material for both infrared and ultraviolet wavelengths, since it is transparent in these regions (about 0.15 µm to 9 µm) and exhibits extremely low refractive index. Furthermore the material is attacked by few reagents. At wavelengths as short as 157 nm, a common wavelength used for semiconductor stepper manufacture for integrated circuit lithography, the refractive index of calcium fluoride shows some non-linearity at high power densities which has inhibited its use for this purpose. In the early years of the 21st century the stepper market for calcium fluoride collapsed and many large manufacturing facilities have been closed. Canon and other manufacturers have used synthetically grown crystals of calcium fluoride components in lenses to aid apochromatic design, and to reduce light dispersion. This use has largely been superseded by newer glasses and computer aided design. As an infrared optical material, calcium fluoride is widely available and was sometimes known by the Eastman Kodak trademarked name "Irtran-3," although this designation is obsolete.

Uranium-doped calcium fluoride was the second type of solid state laser invented, in the 1960s. Peter Sorokin and Mirek Stevenson at IBM's laboratories in Yorktown Heights (US) achieved lasing at 2.5 µm shortly after Maiman's ruby laser.

It is also used as a flux for melting and liquid processing of iron, steel and their composites. Its action is based on its similar melting point to iron, on its ability to dissolve oxides and on its ability to wet oxides and metals.


Fluorides are toxic to humans, however CaF2 is considered relatively harmless due to its extreme insolubility. The situation is analogous to BaSO4, where the toxicity normally associated with Ba2+ is offset by the very low solubility of its sulfate derivative.


  1. ^ Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN 0070494398
  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. ^ G. L. Miessler and D. A. Tarr “Inorganic Chemistry” 3rd Ed, Pearson/Prentice Hall publisher, ISBN 0-13-035471-6.
  4. ^ Aigueperse, Jean; Paul Mollard, Didier Devilliers, Marius Chemla, Robert Faron, Renée Romano, Jean Pierre Cuer (2005), "Fluorine Compounds, Inorganic", Ullmann's Encyclopedia of Industrial Chemistry, Weinheim: Wiley-VCH, doi:10.1002/14356007.a11_307  
  5. ^ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.

See also


Related materials

External links


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