This is a good article. Click here for more information.

Caesium fluoride

From Wikipedia, the free encyclopedia
Jump to: navigation, search
Caesium fluoride
Caesium fluoride
Caesium fluoride
IUPAC name
Caesium fluoride
Other names
Cesium fluoride
13400-13-0 YesY
ChemSpider 24179 YesY
Jmol-3D images Image
RTECS number FK9650000
Molar mass 151.90 g/mol
Appearance white crystalline solid
Density 4.115 g/cm3
Melting point 682 °C (1,260 °F; 955 K)
Boiling point 1,251 °C (2,284 °F; 1,524 K)
367 g/100 ml (18 °C)
Crystal structure cubic, cF8
Space group Fm3m, No. 225
Dipole moment 7.9 D
MSDS External MSDS
EU Index ot listed
Flash point Non-flammable
Related compounds
Other anions
Caesium chloride
Caesium bromide
Caesium iodide
Caesium astatide
Other cations
Lithium fluoride
Sodium fluoride
Potassium fluoride
Rubidium fluoride
Except where noted otherwise, data is given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
 YesY verify (what isYesY/N?)
Infobox references

Caesium fluoride or cesium fluoride is an inorganic compound usually encountered as a hygroscopic white solid. It is used in organic synthesis as a source of the fluoride anion. Caesium has the lowest electronegativity of all non-radioactive elements and fluorine has the highest electronegativity of all elements. It therefore has the strongest ionic compound bond discovered to be non-radioactive.

Synthesis and properties[edit]

Caesium fluoride can be prepared by the reaction of caesium hydroxide (CsOH) with hydrofluoric acid (HF). The resulting salt can then be purified by recrystallization. The reaction is shown below:

CsOH(aq) + HF(aq) → CsF(aq) + H2O(l)

Another way to make caesium fluoride is to react caesium carbonate (Cs2CO3) with hydrofluoric acid. The resulting salt can then be purified by recrystallization. The reaction is shown below:

Cs2CO3(aq) + 2HF(aq) → 2CsF(aq) + H2O(l) + CO2(g)

In addition, elemental fluorine and caesium can be used to form caesium fluoride as well, but doing so is very impractical because of the expense.[1] While this is not a normal route of preparation, caesium metal reacts vigorously with all the halogens to form caesium halides. Thus, it burns with fluorine gas, F2, to form caesium fluoride, CsF according to the following reaction:

2Cs(s) + F2(g) → 2CsF(s)

CsF is more soluble 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.[2] CsF reaches a vapor pressure of 1 kilopascal at 825 °C, 10 kPa at 999 °C, and 100 kPa at 1249 °C.[3]


Caesium fluoride has the halite structure, which means that the Cs+ and F pack in a cubic closest packed array as do Na+ and Cl in sodium chloride.[4][5]

Applications in organic synthesis[edit]

Being highly dissociated it is a more reactive source of fluoride than related salts. CsF is less hygroscopic alternative to tetra-n-butylammonium fluoride (TBAF) and TAS-fluoride (TASF) when anhydrous "naked" fluoride ion is needed.

As a base[edit]

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.[4] CsF gives higher yields in Knoevenagel condensation reactions than KF or NaF.[6]

Formation of Cs-F bonds[edit]

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. It will convert electron-deficient aryl chlorides to aryl fluorides (halex reaction).[7]

Deprotection agent[edit]

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. 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. Solutions of caesium fluoride in THF or DMF attack a wide variety of organosilicon compounds to produce an organosilicon fluoride and a carbanion, which can then react with electrophiles,[5] for example:[6]

CsF desilylation.png

Desilylation is also useful for the removal of silyl protecting groups.[8]

Other uses[edit]

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


Like other soluble fluorides, CsF is moderately toxic.[9] Contact with acid should be avoided, as this forms highly toxic/corrosive hydrofluoric acid. The caesium ion (Cs+) and caesium chloride are generally not considered toxic.[10]


  1. ^ Reacting Fluorine with Caesium
  2. ^ Friestad, G. K.; Branchaud, B. P. (1999). Reich, H. J.; Rigby, J. H., ed. Handbook of Reagents for Organic Synthesis: Acidic and Basic Reagents. New York: Wiley. pp. 99–103. 
  3. ^ 6-63
  4. ^ a b Greenwood, N.N.; Earnshaw, A. Chemistry of the Elements, Pergamon Press, Oxford, UK, 1984.
  5. ^ a b Lide, D. R., ed. (2005). CRC Handbook of Chemistry and Physics (86th ed.). Boca Raton (FL): CRC Press. ISBN 0-8493-0486-5. 
  6. ^ a b Fiorenza, M; Mordini, A; Papaleo, S; Pastorelli, S; Ricci, A (1985). "Fluoride ion induced reactions of organosilanes: the preparation of mono and dicarbonyl compounds from β-ketosilanes". Tetrahedron Letters 26 (6): 787. doi:10.1016/S0040-4039(00)89137-6. 
  7. ^ F. W. Evans, M. H. Litt, A. M. Weidler-Kubanek, F. P. Avonda (1968). "Reactions Catalyzed by Potassium Fluoride. 111. The Knoevenagel Reaction". Journal of Organic Chemistry 33 (5): 1837–1839. doi:10.1021/jo01269a028. 
  8. ^ Adam P. Smith, Jaydeep J. S. Lamba, and Cassandra L. Fraser (2002). "Efficient Synthesis of Halomethyl-2,2'-bipyridines: 4,4'-Bis(chloromethyl)-2,2'-bipyridine". Org. Synth. 78: 82. ; Coll. Vol. 10, p. 107 
  9. ^ MSDS Listing for cesium fluoride. MSDS Date: April 27, 1993. Retrieved on September 7, 2007.
  10. ^ "MSDS Listing for cesium chloride." MSDS Date: January 16, 2006. Retrieved on September 7, 2007.

External links[edit]