Difference between revisions of "Caesium fluoride"
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| Name = Caesium fluoride | | Name = Caesium fluoride | ||
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| Section2 = {{Chembox Properties | | Section2 = {{Chembox Properties | ||
| − | | Reference = <ref>{{RubberBible62nd|page=B-91}}.</ref> | + | | Reference = <ref name="CRC">{{RubberBible62nd|page=B-91}}.</ref> |
| Formula = CsF | | Formula = CsF | ||
| MolarMass = 151.90 g/mol | | MolarMass = 151.90 g/mol | ||
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| + | | Name = Caesium fluoride sesquihydrate | ||
| + | | IUPACName = Caesium fluoride—water (2/3) | ||
| + | | Section1 = {{Chembox Identifiers | ||
| + | | InChI=1/2Cs.2FH.3H2O/h;;2*1H;3*1H2/q2*+1;;;;;/p-2 | ||
| + | }} | ||
| + | | Section2 = {{Chembox Properties | ||
| + | | Reference = <ref name="CRC"/> | ||
| + | | Formula = CsF·1½H<sub>2</sub>O | ||
| + | | MolarMass = 178.93 g/mol | ||
| + | | MeltingPt = 703 ºC | ||
| + | | Solubility = 366.5 g/100 ml (18 ºC) | ||
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'''Caesium fluoride''' is an ionic compound usually found as a hygroscopic white solid. It is more soluble and more readily [[Dissociation (chemistry)|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''.<ref name="Friestad">{{citation | last1 = Friestad | first1 = G. K. | last2 = Branchaud | first2 = B. P. | title = Handbook of Reagents for Organic Synthesis: Acidic and Basic Reagents | editor1-last = Reich | editor1-first = H. J. | editor2-last = Rigby | editor2-first = J. H. | publisher = Wiley | location = New York | year = 1999 | pages = 99–103}}.</ref> It is therefore a useful, less [[hygroscopic]] alternative to [[tetrabutylammonium fluoride]] (TBAF) and [[TASF reagent|TAS-fluoride]] (TASF) when anhydrous "naked" [[fluoride]] [[ion]] is needed. Like all soluble fluorides, it is mildly [[base (chemistry)|basic]]. Contact with [[acid]] should be avoided, as this forms highly toxic/corrosive [[hydrofluoric acid]]. | '''Caesium fluoride''' is an ionic compound usually found as a hygroscopic white solid. It is more soluble and more readily [[Dissociation (chemistry)|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''.<ref name="Friestad">{{citation | last1 = Friestad | first1 = G. K. | last2 = Branchaud | first2 = B. P. | title = Handbook of Reagents for Organic Synthesis: Acidic and Basic Reagents | editor1-last = Reich | editor1-first = H. J. | editor2-last = Rigby | editor2-first = J. H. | publisher = Wiley | location = New York | year = 1999 | pages = 99–103}}.</ref> It is therefore a useful, less [[hygroscopic]] alternative to [[tetrabutylammonium fluoride]] (TBAF) and [[TASF reagent|TAS-fluoride]] (TASF) when anhydrous "naked" [[fluoride]] [[ion]] is needed. Like all soluble fluorides, it is mildly [[base (chemistry)|basic]]. Contact with [[acid]] should be avoided, as this forms highly toxic/corrosive [[hydrofluoric acid]]. | ||
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==External links== | ==External links== | ||
| + | *{{EHC|36|name=Fluorine and Fluorides}} | ||
| + | *{{EHC|227|name=Fluorides}} | ||
*{{NPI|id=44|name=Fluoride compounds}} | *{{NPI|id=44|name=Fluoride compounds}} | ||
Latest revision as of 05:48, 28 November 2009
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Caesium fluoride 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.[3] It is therefore a useful, less hygroscopic alternative to tetrabutylammonium 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.
Contents
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:[4]
- 2CsF + CaCl2 → 2CsCl + CaF2
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 silicon–fluorine 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 (vide infra). As with other soluble fluorides, CsF is moderately basic, because hydrofluoric acid is a weak acid. The low nucleophilicity of fluoride means it can be a useful base in organic chemistry.[4]
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.[4][1]
Preparation
Caesium fluoride may be prepared by the action of hydrofluoric acid on caesium hydroxide or caesium carbonate, followed by removal of water.
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.[5]
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,[1] 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.
Precautions
Like other soluble fluorides, CsF is moderately toxic.[7] Contact with acid should be avoided, as this forms highly toxic/corrosive hydrofluoric acid. The caesium ion (Cs+), or caesium chloride, is generally not considered toxic.[8]
References
- ↑ 1.0 1.1 1.2 1.3 CRC Handbook of Chemistry and Physics, 62nd ed.; Weast, Robert C., Ed.; CRC Press: Boca Raton, FL, 1981; p B-91. ISBN 0-8493-0462-8. Cite error: Invalid
<ref>tag; name "CRC" defined multiple times with different content - ↑ HSNO Chemical Classification Information Database, <http://www.ermanz.govt.nz/Chemicals/ChemicalDisplay.aspx?SubstanceID=16459> (accessed 25 August 2009), New Zealand Environmental Risk Management Authority.
- ↑ 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.
- ↑ 4.0 4.1 4.2 Greenwood, Norman N.; Earnshaw, A. Chemistry of the Elements; Pergamon: Oxford, 1984. ISBN 0-08-022057-6.
- ↑ 5.0 5.1 Fiorenza, Mariella; Mordini, Alessandro; Papaleo, Sandro; Pastorelli, Stefania; Ricci, Alfredo Fluoride ion induced reactions of organosilanes: the preparation of mono and dicarbonyl compounds from β-ketosilanes. Tetrahedron Lett. 1985, 26, 787–88. DOI: 10.1016/S0040-4039(00)89137-6.
- ↑ Evans, F. W.; Litt, M. H.; Weidler-Kubanek, A. M.; Avonda, F. P. Formation of adducts between fluorinated ketones and metal fluorides. J. Org. Chem. 1968, 33, 1837–39. DOI: 10.1021/jo01269a028.
- ↑ MSDS Listing for cesium fluoride, <http://www.hazard.com/msds/f2/bms/bmsqc.html> (accessed 7 September 2007), hazard.com; .
- ↑ MSDS Listing for cesium chloride, <http://www.jtbaker.com/msds/englishhtml/c1903.htm> (accessed 7 September 2007), jtbaker.com; .
External links
- IPCS Environmental Health Criteria 36: Fluorine and Fluorides
- IPCS Environmental Health Criteria 227: Fluorides
- Entry for "Fluoride compounds" on the Australian National Pollutant Inventory
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