Caesium fluoride
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| Names | |
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IUPAC name
Caesium fluoride
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| Other names
Cesium fluoride
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| Identifiers | |
13400-13-0 
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| 3D model (Jmol) | Interactive image |
| ChemSpider |
24179
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| ECHA InfoCard | 100.033.156 |
| PubChem | 25953 |
| RTECS number | FK9650000 |
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| Properties | |
| CsF | |
| 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) | |
| -44.5·10−6 cm3/mol | |
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Refractive index (nD)
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1.477 |
| Structure | |
| cubic, cF8 | |
| Fm3m, No. 225 | |
| Octahedral | |
| 7.9 D | |
| Thermochemistry | |
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Std enthalpy of
formation (ΔfH |
-555 kJ/mol |
| Hazards | |
| Safety data sheet | External MSDS |
| Flash point | Non-flammable |
| Related compounds | |
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Other anions
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Caesium chloride Caesium bromide Caesium iodide Caesium astatide |
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Other cations
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Lithium fluoride Sodium fluoride Potassium fluoride Rubidium fluoride Francium fluoride |
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Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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 (what is  ?) |
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| 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 highest electropositivity of all non-radioactive elements and fluorine has the highest electronegativity of all elements.
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:
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:
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. 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:
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. CsF reaches a vapor pressure of 1 kilopascal at 825 °C, 10 kPa at 999 °C, and 100 kPa at 1249 °C.
CsF chains with a thickness as small as one or two atoms can be grown inside carbon nanotubes.
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.
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