Silicon tetrafluoride
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| Names | |
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| IUPAC names
Tetrafluorosilane
Silicon tetrafluoride | |
| Other names
Silicon fluoride
Fluoro acid air | |
| Identifiers | |
3D model (JSmol)
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| ECHA InfoCard | 100.029.104 |
PubChem CID
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| RTECS number |
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| UN number | 1859 |
CompTox Dashboard (EPA)
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| Properties | |
| SiF4 | |
| Molar mass | 104.0791 g/mol |
| Appearance | colourless gas, fumes in moist air |
| Density | 1.66 g/cm3, solid (−95 °C) 4.69 g/L (gas) |
| Melting point | −90 °C (−130 °F; 183 K) |
| Boiling point | −86 °C (−123 °F; 187 K) |
| decomposes | |
| Structure | |
| tetrahedral | |
| 0 D | |
| Hazards | |
| Occupational safety and health (OHS/OSH): | |
Main hazards
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toxic, corrosive |
| NFPA 704 (fire diamond) | |
| Lethal dose or concentration (LD, LC): | |
LCLo (lowest published)
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69,220 mg/m3 (rat, 4 hr)[1] |
| Safety data sheet (SDS) | ICSC 0576 |
| Related compounds | |
Other anions
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Silicon tetrachloride Silicon tetrabromide Silicon tetraiodide |
Other cations
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Carbon tetrafluoride Germanium tetrafluoride Tin tetrafluoride Lead tetrafluoride |
Related compounds
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Hexafluorosilicic acid |
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Silicon tetrafluoride or tetrafluorosilane is the chemical compound with the formula SiF4. This colorless compound is notable for having a narrow liquid range: its boiling point is only 4 °C above its melting point. It was first synthesized by John Davy in 1812.[2] It is a tetrahedral molecule.
Preparation
SiF
4 is a by-product of the production of phosphate fertilizers, resulting from the attack of HF (derived from fluorapatite protonolysis) on silicates, which are present as impurities in the phosphate rock. In the laboratory, the compound is prepared by heating BaSiF
6 above 300 °C, whereupon the solid releases volatile SiF
4, leaving a residue of BaF
2. The required BaSiF
6 is prepared by treating aqueous hexafluorosilicic acid with barium chloride.[3] The corresponding GeF
4 is prepared analogously, except that the thermal "cracking" requires 700 °C.[4]
SiF
4 can in principle also be generated by the reaction of silicon dioxide and hydrofluoric acid, but this process tends to give hexafluorosilicic acid:
- 6 HF + SiO2 → H2SiF6 + 2 H2O
Uses
This volatile compound finds limited use in microelectronics and organic synthesis.[5]
Occurrence
Volcanic plumes contain significant amounts of silicon tetrafluoride. Production can reach several tonnes per day.[6] The silicon tetrafluoride is partly hydrolysed and forms hexafluorosilicic acid.
References
- ^ "Fluorides (as F)". Immediately Dangerous to Life or Health Concentrations (IDLH). National Institute for Occupational Safety and Health (NIOSH).
- ^ John Davy (1812). "An Account of Some Experiments on Different Combinations of Fluoric Acid". Philosophical Transactions of the Royal Society of London. 102: 352–369. doi:10.1098/rstl.1812.0020. ISSN 0261-0523. JSTOR 107324.
- ^ "Silicon Tetrafluoride". Inorganic Syntheses. 4: 145–6. 1953. doi:10.1002/9780470132357.ch47.
{{cite journal}}: Unknown parameter|authors=ignored (help) - ^ "Silicon Tetrafluoride". Inorganic Syntheses. 4: 147–8. 1953. doi:10.1002/9780470132357.ch48.
{{cite journal}}: Unknown parameter|authors=ignored (help) - ^ Shimizu, M. "Silicon(IV) Fluoride" Encyclopedia of Reagents for Organic Synthesis, 2001 John Wiley & Sons. doi:10.1002/047084289X.rs011
- ^ T. Mori; M. Sato; Y. Shimoike; K. Notsu (2002). "High SiF4/HF ratio detected in Satsuma-Iwojima volcano's plume by remote FT-IR observation" (PDF). Earth Planets Space. 54: 249–256.
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