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Caesium
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.

Contents

1 Synthesis and properties 2 Structure 3 Applications in organic synthesis

3.1 As a base 3.2 Formation of C-F bonds 3.3 Deprotection agent 3.4 Other uses

4 Precautions 5 References 6 External links

Synthesis and properties[edit]

Crystalline CsF chains grown inside double-wall carbon nanotubes.[5]

Caesium
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) + 2 HF(aq) → 2 CsF(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.[6] 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:

2 Cs(s) + F2(g) → 2 CsF(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.[7] CsF reaches a vapor pressure of 1 kilopascal at 825 °C, 10 kPa at 999 °C, and 100 kPa at 1249 °C.[8] CsF chains with a thickness as small as one or two atoms can be grown inside carbon nanotubes.[5] Structure[edit] Caesium
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.[3] 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.[9] CsF gives higher yields in Knoevenagel condensation
Knoevenagel condensation
reactions than KF or NaF.[10] Formation of C-F bonds[edit] Caesium
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).[11] Deprotection agent[edit] 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. 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
THF
or DMF attack a wide variety of organosilicon compounds to produce an organosilicon fluoride and a carbanion, which can then react with electrophiles, for example:[10]

Desilylation
Desilylation
is also useful for the removal of silyl protecting groups.[12] 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. Precautions[edit] Like other soluble fluorides, CsF is moderately toxic.[13] 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.[14] References[edit]

^ a b c d Haynes, William M., ed. (2011). CRC Handbook of Chemistry and Physics (92nd ed.). Boca Raton, FL: CRC Press. p. 4.57. ISBN 1439855110.  ^ Haynes, William M., ed. (2011). CRC Handbook of Chemistry and Physics (92nd ed.). Boca Raton, FL: CRC Press. p. 4.132. ISBN 1439855110.  ^ a b c d Davey, Wheeler P. (1923). "Precision Measurements of Crystals of the Alkali Halides". Physical Review. 21 (2): 143. doi:10.1103/PhysRev.21.143.  ^ a b c d Haynes, William M., ed. (2011). CRC Handbook of Chemistry and Physics (92nd ed.). Boca Raton, FL: CRC Press. p. 5.10. ISBN 1439855110.  ^ a b Senga, Ryosuke; Suenaga, Kazu (2015). "Single-atom electron energy loss spectroscopy of light elements". Nature Communications. 6: 7943. doi:10.1038/ncomms8943. PMC 4532884 . PMID 26228378.  (Supplementary information) ^ Reacting Fluorine
Fluorine
with Caesium. Youtube ^ Friestad, G. K.; Branchaud, B. P. (1999). Reich, H. J.; Rigby, J. H., eds. Handbook of Reagents for Organic Synthesis: Acidic and Basic Reagents. New York: Wiley. pp. 99–103. ISBN 978-0-471-97925-8.  ^ Lide, D. R., ed. (2005). "Vapor Pressure". CRC Handbook of Chemistry and Physics (PDF) (86th ed.). Boca Raton (FL): CRC Press. p. 6.63. ISBN 0-8493-0486-5.  ^ Greenwood, Norman N.; Earnshaw, Alan (1984). Chemistry of the Elements. Oxford: Pergamon Press. pp. 82–83. ISBN 0-08-022057-6.  ^ a b Fiorenza, M; Mordini, A; Papaleo, S; Pastorelli, S; Ricci, A (1985). " Fluoride
Fluoride
ion induced reactions of organosilanes: the preparation of mono and dicarbonyl compounds from β-ketosilanes". Tetrahedron Letters. 26 (6): 787–788. doi:10.1016/S0040-4039(00)89137-6.  ^ Evans, F. W.; Litt, M. H.; Weidler-Kubanek, A. M.; Avonda, F. P. (1968). "Formation of adducts between fluorinated ketones and metal fluorides". Journal of Organic Chemistry. 33 (5): 1837–1839. doi:10.1021/jo01269a028.  ^ Smith, Adam P.; Lamba, Jaydeep J. S.; Fraser, Cassandra L. (2002). "Efficient Synthesis of Halomethyl-2,2'-bipyridines: 4,4'-Bis(chloromethyl)-2,2'-bipyridine". Organic Syntheses. 78: 82. ; Collective Volume, 10, p. 107  ^ MSDS Listing for cesium fluoride. www.hazard.com. MSDS Date: April 27, 1993. Retrieved on September 7, 2007. ^ "MSDS Listing for cesium chloride." www.jtbaker.com. MSDS Date: January 16, 2006. Retrieved on September 7, 2007.

External links[edit]

Wikimedia Commons has media related to Caesium
Caesium
fluoride.

National Pollutant Inventory- Fluoride
Fluoride
and compounds fact sheet

Authority control

GND: 4249758-9

v t e

Caesium
Caesium
compounds

CsBr CsCl CsClO4 Cs2CrO4 CsF CsH CsI CsLiB6O10 CsNO3 CsOH Cs2CO3 CsHCO3 Cs2SO4 CsC2H3O2 Cs2O Cs2TiO3 Cs2WO4 CsAu

v t e

Fluorine
Fluorine
compounds

Binary compounds

AcF3 AgF AgF2 Ag2F AlF3 AmF3 AmF4 AsF3 AsF5 AuF3 BF3 B2F4 BaF2 BeF2 BiF3 BiF5 BrF3 BrF5 CF4 C2F6 CaF2 CdF2 CeF3 ClF3 ClF5 CoF2 CoF3 CrF2 CrF3 CrF4 CrF5 CrF6 CsF CuF DyF3 ErF3 EuF2 EuF3 DF HF FI KF LiF NaF RbF TlF FeF2 FeF3 HgF2 Hg2F2 KrF2 LaF3 LuF3 MgF2 MnF2 MoF4 MoF5 MoF6 NiF2 F2O F2O4 PbF2 PbF4 PdF2 SnF2 SrF2 XeF2 ZnF2 ZrF2 GaF3 GdF3 HoF3 InF3 MnF3 F3N NbF4 NbF5 NdF3 F3P PrF3 PuF3 SbF3 ScF3 SmF3 TbF3 TiF3 TiF4 TlF3 TmF3 UF3 VF3 YF3 YbF3 GeF4 HfF4 F4N2 PuF4 SF4 SeF4 SiF4 SnF4 TeF4 ThF4 UF4 VF4 WF4 XeF4 ZrF4 IF5 PF5 SbF5 TaF5 UF5 VF5 WF5 OsF6 PtF6 ReF6 SF6 SeF6 TeF6 UF6 WF6 XeF6 IF7 ReF7 IrF3 IrF6 OsF4 OsF5 PmF3 PrF4 PuF6 ReF4 ReF5 RhF6 RuF3 RuF4 RuF5 RuF6 TcF5 TcF6 VF2 YbF2

Other

AgBF4 AgPF6 Cs2AlF5 K3AlF6 Na3AlF6 KAsF6 LiAsF6 NaAsF6 HBF4 KBF4 LiBF4 NaBF4 RbBF4 Ba(BF4)2 Ni(BF4)2 Pb(BF4)2 Sn(BF4)2 BaClF BaSiF6 BaGeF6 BrOF3 BrO2F CBrF3 CBr2F2 CBr3F CClF3 CCl2F2 CCl3F CFN CF2O CF3I CHF3 CH2F2 CH3F C2Cl3F3 C2H3F C6H5F C7H5F3 C15F33N ClFO2 CrFO4 CrF2O2 CsBF4 NH4F FNO FNO2 FNO3 KHF2 NaHF2 ThOF2 NH5F2 (NH4)2SiF6 F2OS F3OP PSF3 HPF6 HSbF6 NH4PF6 KPF6 KSbF6 LiPF6 NaPF6 NaSbF6 Na2SiF6 Na2TiF6 Na2ZrF6 TlPF6 IOF3 K2NbF7 K2TaF7 IO3F

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