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Caesium fluoride

Caesium fluoride
IUPAC name Caesium fluoride
Other names Cesium fluoride
CAS number 13400-13-0
RTECS number FK9650000
Molecular formula CsF
Molar mass 151.904 g/mol
Appearance white crystalline solid
Density 4.115 g/cm³, solid
Melting point

682 °C (955 K)

Boiling point

1251 °C (1524 K)

Solubility in water 367 g/100 ml (18 °C)
Crystal structure cubic
Dipole moment 7.9 D
MSDS External MSDS
Main hazards toxic, forms HF with acid
R-phrases 23/24/25, 34
S-phrases 26, 36/37, 39, 45
Related Compounds
Other anions caesium chloride
caesium bromide
caesium iodide
Other cations sodium fluoride
potassium fluoride
rubidium fluoride
Related compounds tetra-n-butylammonium fluoride
Except where noted otherwise, data are given for
materials in their standard state
(at 25 °C, 100 kPa)

Infobox disclaimer and references

Caesium fluoride (cesium fluoride in North America), is an ionic compound usually found as a hygroscopic white solid. It is more soluble and more readily dissociated than sodium fluoride or potassium fluoride. It is available in anhydrous form, and if water has been absorbed it is easy to dry by heating at 100 °C for two hours in vacuo.[1] It is therefore a useful, less hygroscopic alternative to tetra-n-butylammonium fluoride (TBAF) and TAS-fluoride (TASF) when anhydrous "naked" fluoride ion is needed. Like all soluble fluorides, it is mildly basic. Contact with acid should be avoided, as this forms highly toxic/corrosive hydrofluoric acid.


Chemical properties

Caesium fluoride reacts usually as a source of fluoride ion, F-. It therefore undergoes all of the usual reactions associated with soluble fluorides such as potassium fluoride, for example:[2]

2 CsF (aq) + CaCl2 (aq) → 2 CsCl (aq) + CaF2 (s)

Being highly dissociated it is quite reactive as a fluoride source under anhydrous conditions too, and it will react with electron-deficient aryl chlorides to form aryl fluorides (halex reaction). Due to the strength of the SiF bond, fluoride ion is useful for desilylation reactions (removal of Si groups) in organic chemistry; caesium fluoride is an excellent source of anhydrous fluoride for such reactions (vida infra). As with other soluble fluorides, CsF is moderately basic, because HF is a weak acid. The low nucleophilicity of fluoride means it can be a useful base in organic chemistry.[2]

Crystal structure

Caesium fluoride has the inverse halite structure, because caesium ions are larger than fluoride ions; in the lithium, sodium, potassium, and rubidium halides, the cation is smaller than the anion. The caesium ions form a cubic closest packed array with fluoride ions in the octahedral holes.[2][3]


Caesium fluoride may be prepared by the action of hydrofluoric acid on caesium hydroxide or caesium carbonate, followed by removal of water.


Caesium fluoride is a useful base in organic chemistry, due the fact that fluoride ion is largely unreactive as a nucleophile. It is reported that CsF gives higher yields in Knoevenagel condensation reactions than KF or NaF.[4]

Removal of silicon groups (desilylation) is a major application for CsF in the laboratory, as its anhydrous nature allows clean formation of water-sensitive intermediates. Caesium fluoride in THF or DMF can attack a wide variety of organosilicon compounds to produce an organosilicon fluoride and a carbanion, which can then react with electrophiles,[3] for example:[4]

Desilylation is also useful for the removal of silyl protecting groups.

Caesium fluoride is also a popular source of fluoride in organofluorine chemistry. For example, CsF reacts with hexafluoroacetone to form a caesium perfluoroalkoxide salt which is stable up to 60 °C, unlike the corresponding sodium or potassium salt.[5]

Single crystals of the salt are transparent into the deep infrared. For this reason it is often used as the windows of cells used for infrared spectroscopy.


Like other soluble fluorides, CsF is moderately toxic.[6] Contact with acid should be avoided, as this forms highly toxic/corrosive hydrofluoric acid. Cesium ion (Cs+), or Cesium chloride, is generally not considered toxic.[7]


  1. ^ Friestad, G. K.; Branchaud, B. P. in: Handbook of Reagents for Organic Synthesis: Acidic and Basic Reagents, (Reich, H. J.; Rigby, J. H. eds.), Wiley, New York, 1999. pp. 99-103.
  2. ^ a b c Greenwood, N.N.; Earnshaw, A. Chemistry of the Elements, Pergamon Press, Oxford, UK, 1984.
  3. ^ a b Handbook of Chemistry and Physics, 71st edition, CRC Press, Ann Arbor, Michigan, 1990.
  4. ^ a b Fiorenza, M.; Mordini, A.; Papaleo, S.; Pastorelli, S.; Ricci, A. 1985. Tetrahedron Letters. 26, 787.
  5. ^ F. W. Evans, M. H. Litt, A. M. Weidler-Kubanek, F. P. Avonda "Reactions Catalyzed by Potassium Fluoride. 111. The Knoevenagel Reaction." 1968. Journal of Organic Chemistry. 33, 1837-1839. Retrieved on September 7, 2007.
  6. ^ "MSDS Listing for cesium flouride." MSDS Date: April 27, 1993. Retrieved on September 7, 2007.
  7. ^ "MSDS Listing for cesium chloride." MSDS Date: January 16, 2006. Retrieved on September 7, 2007.
This article is licensed under the GNU Free Documentation License. It uses material from the Wikipedia article "Caesium_fluoride". A list of authors is available in Wikipedia.
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