Caesium fluoride
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| Caesium fluoride | |
|---|---|
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| General | |
| Systematic name | Caesium fluoride |
| Other names | Cesium fluoride |
| Molecular formula | CsF |
| Molar mass | 151.904 g/mol |
| Appearance | white crystalline solid |
| CAS number | [13400-13-0] |
| Properties | |
| Density and phase | 4.115 g/cm³, solid |
| Solubility in water | 367 g/100 ml (18 °C) |
| In methanol | soluble |
| In dioxane, pyridine | insoluble |
| Melting point | 682 °C (955 K) |
| Boiling point | 1251 °C (1524 K) |
| Basicity (pKb) | ? |
| Structure | |
| Coordination geometry |
? |
| Crystal structure | cubic |
| Dipole moment | 7.9 D |
| Hazards | |
| MSDS | External MSDS |
| Main hazards | toxic, forms HF with acid |
| NFPA 704 |    estimated |
| R/S statement | R: 23/24/25, 34 S: 26, 36/37, 39, 45 |
| RTECS number | FK9650000 |
| Supplementary data page | |
| Structure and properties |
n, εr, etc. |
| Thermodynamic data | Phase behaviour Solid, liquid, gas |
| Spectral data | UV, IR, NMR, MS |
| Related compounds | |
| Other anions | caesium chloride caesium bromide caesium iodide |
| Other cations | sodium fluoride potassium fluoride rubidium fluoride |
| Related compounds | tetra-n-butyl- ammonium fluoride |
| Except where noted otherwise, data are given for materials in their standard state (at 25°C, 100 kPa) Infobox disclaimer and references |
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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. CsF is commercially available – on a lab scale it costs around $50 per 100g (Synquest), cheaper than RbF. 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[3]. 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.
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[edit] 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 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 Si-F 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 (see uses below). As with other soluble fluorides, CsF is moderately basic, due to the fact that HF is a weak acid. The low nucleophilicity of fluoride means it can be a useful base in organic chemistry (see uses below).
[edit] Preparation
Caesium fluoride may be prepared by the action of hydrofluoric acid on caesium hydroxide or caesium carbonate, followed by removal of water.
[edit] Uses
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[5]: 
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[6].
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.
[edit] Precautions
Like soluble fluorides, CsF is moderately toxic, see Elf Atochem MSDS sheet. Contact with acid should be avoided, as this forms highly toxic/corrosive hydrofluoric acid. Caesium ion (Cs+) per se (for example as CsCl) is not generally considered toxic.
[edit] References
- N. N. Greenwood, A. Earnshaw, Chemistry of the Elements, Pergamon Press, Oxford, UK, 1984.
- Handbook of Chemistry and Physics, 71st edition, CRC Press, Ann Arbor, Michigan, 1990.
- G. K. Friestad, B. P. Branchaud, in: Handbook of Reagents for Organic Synthesis: Acidic and Basic Reagents, (H. J. Reich, J. H. Rigby, eds.), pp99-103, Wiley, New York, 1999.
- L. Rand, J. V. Swisher, C. J. Cronin Journal of Organic Chemistry 27, 3505 (1962).
- M. Fiorenza, A. Mordini, S. Papaleo, S. Pastorelli, A. Ricci Tetrahedron Letters 26, 787 (1985).
- F. W. Evans, M. H. Litt, A. M. Weidler-Kubanek, F. P. Avonda Journal of Organic Chemistry 33, 1837-1839 (1968).
