Beryllium fluoride
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| Names | |
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| IUPAC name
Beryllium fluoride
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| Other names
Beryllium difluoride
Difluoroberyllane | |
| Identifiers | |
3D model (JSmol)
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| ChEBI | |
| ChemSpider | |
| ECHA InfoCard | 100.029.198 |
PubChem CID
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| RTECS number |
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CompTox Dashboard (EPA)
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| Properties | |
| BeF2 | |
| Molar mass | 47.01 g/mol hygroscopic |
| Appearance | colorless lumps |
| Density | 1.986 g/cm3 |
| Melting point | 554 °C (1,029 °F; 827 K) |
| Boiling point | 1,169 °C (2,136 °F; 1,442 K) |
| very soluble | |
| Solubility | sparingly soluble in alcohol |
| Structure | |
| Linear | |
| Thermochemistry | |
Heat capacity (C)
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1.102 J/K |
Std enthalpy of
formation (ΔfH⦵298) |
-21.84 kJ/g |
| Hazards | |
| Flash point | non-flammable |
| Lethal dose or concentration (LD, LC): | |
LD50 (median dose)
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98 mg/kg (oral, rat) |
| Related compounds | |
Other anions
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Beryllium chloride Beryllium bromide Beryllium iodide |
Other cations
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Magnesium fluoride Calcium fluoride Strontium fluoride Barium fluoride Radium fluoride |
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Beryllium fluoride is the inorganic compound with the formula BeF2. This white solid is the principal precursor for the manufacture of beryllium metal.
Structure and bonding
The structure of solid BeF2 resembles that of silicon dioxide. Be2+ centers are four coordinate and tetrahedral.[2] Solid BeF2 adopts a number of polymeric structures analogous to those adopted by SiO2 namely α-quartz, β-quartz, crystobalite and tridymite, and its reactions with fluorides are quite analogous to the reactions of SiO2 with oxides.[3] An analogy exists between BeF2 and AlF3: both adopt extended structures at mild temperature. BeF2 is considered to be highly covalent.
Gaseous BeF2 is found in the gas-phase above 1160 °C. Like the isoelectronic gases CO2 and SiO2, it is a linear molecule. The Be-F distance of 177 pm.[4] The difference between the ambient temperature structures of BeF2 (rock-like solid) and CO2 (gas) reflects the low tendency of alkali metals to form multiple bonds.
BeF2 reaches a vapor pressure of 10 Pa at 686 °C, 100 Pa at 767 °C, 1 kilopascal at 869 °C, 10 kPa at 999 °C, and 100 kPa at 1172 °C.[5]
Molten BeF2 resembles water in some ways, since it is a triatomic molecule with strong interactions via Be—F—Be bonds. As in water, the density of BeF2 decreases near its melting point. Liquid (molten) beryllium fluoride also has a fluctuating tetrahedral structure[6]
Production
The processing of beryllium ores generates impure Be(OH)2. This material reacts with ammonium bifluoride to give ammonium tetrafluoroberyllate:
- Be(OH)2 + 2 (NH4)HF2 → (NH4)2BeF4 + 2 H2O
Tetrafluoroberyllate is a robust ion, which allows its purification by precipitation of various impurities as their hydroxides. Heating purified (NH4)2BeF4 gives the desired product:
- (NH4)2BeF4 → 2 NH3 + 2 HF + BeF2
Applications
Reduction of BeF2 at 1300 °C with magnesium in a graphite crucible provides the most practical route to metallic beryllium:[4]
- BeF2 + Mg → Be + MgF2
The chloride is not a useful precursor because of its volatility.
Niche uses
Beryllium fluoride is used in biochemistry, particularly protein crystallography as a mimic of phosphate. Thus, ADP and beryllium fluoride together tend to bind to ATP sites and inhibit protein action, making it possible to crystallise proteins in the bound state.[7][8]
Beryllium fluoride forms a basic constituent of the preferred fluoride salt mixture used in liquid-fluoride nuclear reactors. Typically beryllium fluoride is mixed with lithium fluoride to form a base solvent (FLiBe), into which fluorides of uranium and thorium are introduced. Beryllium fluoride is exceptionally chemically stable and LiF/BeF2 mixtures (FLiBe) have low melting points (360 C - 459 C) and the best neutronic properties of fluoride salt combinations appropriate for reactor use. MSRE used two different mixtures in the two cooling circuits.
Safety
All beryllium compounds are highly toxic. Beryllium fluoride is very soluble in water and is thus absorbed easily; as mentioned above, it inhibits ATP uptake. The LD50 in mice is about 100 mg/kg by ingestion and 1.8 mg/kg by intravenous injection.
References
- ^ Lide, David R., ed. (2006). CRC Handbook of Chemistry and Physics (87th ed.). Boca Raton, FL: CRC Press. ISBN 0-8493-0487-3.
- ^ Wells A.F. (1984) Structural Inorganic Chemistry 5th edition Oxford Science Publications ISBN 0-19-855370-6
- ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
- ^ a b Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
- ^ http://www.physics.nyu.edu/kentlab/How_to/ChemicalInfo/VaporPressure/morepressure.pdf 6-63
- ^ Agarwal, M.; Chakravarty C (2007). "Waterlike Structural and Excess Entropy Anomalies in Liquid Beryllium Fluoride". J. Phys. Chem. B. 111 (46): 13294. doi:10.1021/jp0753272. PMID 17963376.
{{cite journal}}: CS1 maint: multiple names: authors list (link) - ^ Reiko Kagawa, Martin G Montgomery, Kerstin Braig, Andrew G W Leslie and John E Walker (2004). "The structure of bovine F1-ATPase inhibited by ADP and beryllium fluoride". The EMBO Journal. 23 (5): 2734–2744. doi:10.1038/sj.emboj.7600293. PMC 514953. PMID 15229653.
{{cite journal}}: CS1 maint: multiple names: authors list (link) - ^ Bigay J, Deterre P, Pfister C, Chabre M (1987). "Fluoride complexes of aluminium or beryllium act on G-proteins as reversibly bound analogues of the gamma phosphate of GTP". The EMBO Journal. 6 (10): 2907–2913. PMC 553725. PMID 2826123.
{{cite journal}}: CS1 maint: multiple names: authors list (link)