Caesium fluoride
| |
| |
| Names | |
|---|---|
| IUPAC name
Caesium fluoride
| |
| Other names
Cesium fluoride
| |
| Identifiers | |
3D model (JSmol)
|
|
| ChemSpider | |
| ECHA InfoCard | 100.033.156 
|
| EC Number |
|
PubChem CID
|
|
| RTECS number |
|
| UNII | |
CompTox Dashboard (EPA)
|
|
| Properties | |
| CsF | |
| Molar mass | 151.903 g/mol[1] |
| Appearance | white crystalline solid |
| Density | 4.64 g/cm3[1] |
| Melting point | 703 °C (1,297 °F; 976 K) [1] |
| Boiling point | 1,251 °C (2,284 °F; 1,524 K) |
| 573.0 g/100 mL (25 °C)[1] | |
| Solubility | Insoluble in acetone, diethyl ether, pyridine and ethanol 191 g/100 mL in methanol. |
Magnetic susceptibility (χ)
|
-44.5·10−6 cm3/mol[2] |
Refractive index (nD)
|
1.477 |
| Structure | |
| cubic, cF8 | |
Space group
|
Fm3m, No. 225[3] |
a = 0.6008 nm[3]
| |
Lattice volume (V)
|
0.2169 nm3[3] |
Formula units (Z)
|
4 |
| Octahedral | |
Dipole moment
|
7.9 D |
| Thermochemistry | |
Heat capacity (C)
|
51.1 J/mol·K[4] |
Std molar
entropy (S |
92.8 J/mol·K[4] |
Std enthalpy of
formation (ΔfH⦵298) |
-553.5 kJ/mol[4] |
Gibbs free energy (ΔfG˚)
|
-525.5 kJ/mol[4] |
| Hazards | |
| Main hazards | toxic |
| Safety data sheet | External MSDS |
| GHS pictograms |   
|
| GHS Signal word | Danger |
GHS hazard statements
|
H301, H311, H315, H318, H331, H361f |
| P201, P202, P260, P261, P264, P270, P271, P280, P281, P301+310, P301+330+331, P302+352, P303+361+353, P304+340, P305+351+338, P308+313, P310, P311, P312, P321, P322, P330, P332+313, P361, P362 | |
| NFPA 704 (fire diamond) | 
3
0
0 |
| Flash point | Non-flammable |
| Related compounds | |
Other anions
|
Caesium chloride Caesium bromide Caesium iodide |
Other cations
|
Lithium fluoride Sodium fluoride Potassium fluoride Rubidium fluoride Francium fluoride |
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa). | |
 (what is   ?)
| |
| Infobox references | |
Caesium fluoride or cesium fluoride is an inorganic compound with the formula CsF and it is a hygroscopic white salt. Caesium fluoride can be used in organic synthesis as a source of the fluoride anion. Caesium also has the highest electropositivity of all non-radioactive elements and fluorine has the highest electronegativity of all elements.
Synthesis and properties[]
Caesium fluoride can be prepared by the reaction of caesium hydroxide (CsOH) with hydrofluoric acid (HF) and the resulting salt can then be purified by recrystallization. The reaction is shown below:
- CsOH + HF → CsF + H2O
Using the same reaction, another way to create caesium fluoride is to treat caesium carbonate (Cs2CO3) with hydrofluoric acid and again, the resulting salt can then be purified by recrystallization. The reaction is shown below:
- Cs2CO3 + 2 HF → 2 CsF + H2O + CO2
CsF is more soluble than sodium fluoride or potassium fluoride in organic solvents. It is available in its anhydrous form, and if water has been absorbed, it is easy to dry by heating at 100 °C for two hours in vacuo.[6] CsF reaches a vapor pressure of 1 kilopascal at 825 °C, 10 kPa at 999 °C, and 100 kPa at 1249 °C.[7]
CsF chains with a thickness as small as one or two atoms can be grown inside carbon nanotubes.[5]
Structure[]
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[]
Being highly dissociated, CsF is a more reactive source of fluoride than related salts. CsF is an alternative to tetra-n-butylammonium fluoride (TBAF) and TAS-fluoride (TASF).
As a base[]
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.[8] CsF gives higher yields in Knoevenagel condensation reactions than KF or NaF.[9]
Formation of Cs-F bonds[]
Caesium fluoride serves as a source of fluoride in organofluorine chemistry. Similarly to potassium fluoride, CsF reacts with hexafluoroacetone to form a stable perfluoroalkoxide salt.[10] It will convert electron-deficient aryl chlorides to aryl fluorides (Halex process), although potassium fluoride is more commonly used.
Deprotection agent[]
Due to the strength of the Si–F bond, fluoride is useful for desilylation reactions, i.e. cleavage of Si-O bonds in organic synthesis.[11] CsF is commonly used for such reactions. 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, for example:[9]
Precautions[]
Like other soluble fluorides, CsF is moderately toxic.[12] 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.[13]
References[]
- ^ Jump up to: 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.
- ^ Jump up to: a b c d Davey, Wheeler P. (1923). "Precision Measurements of Crystals of the Alkali Halides". Physical Review. 21 (2): 143–161. Bibcode:1923PhRv...21..143D. doi:10.1103/PhysRev.21.143.
- ^ Jump up to: 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.
- ^ Jump up to: a b Senga, Ryosuke; Suenaga, Kazu (2015). "Single-atom electron energy loss spectroscopy of light elements". Nature Communications. 6: 7943. Bibcode:2015NatCo...6.7943S. doi:10.1038/ncomms8943. PMC 4532884. PMID 26228378. (Supplementary information)
- ^ 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" (PDF). CRC Handbook of Chemistry and Physics (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 978-0-08-022057-4.
- ^ Jump up to: 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): 787–788. 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 Archived 2012-02-09 at the Wayback Machine. www.hazard.com Archived 2013-06-05 at the Wayback Machine. MSDS Date: April 27, 1993. Retrieved on September 7, 2007.
- ^ "MSDS Listing for cesium chloride Archived 2012-03-13 at the Wayback Machine." www.jtbaker.com. MSDS Date: January 16, 2006. Retrieved on September 7, 2007.
| show Authority control |
|---|
- Fluorides
- Caesium compounds
- Metal halides
- Alkali metal fluorides
- Rock salt crystal structure