Summary
Silicon tetrafluoride or tetrafluorosilane is a chemical compound with the formula SiF4. This colorless gas is notable for having a narrow liquid range: its boiling point is only 4 °C above its melting point. It was first prepared in 1771 by Carl Wilhelm Scheele by dissolving silica in hydrofluoric acid.[4], later synthesized by John Davy in 1812.[5] It is a tetrahedral molecule and is corrosive.[6]
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| Names | |
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| IUPAC names
Tetrafluorosilane
Silicon tetrafluoride | |
| Other names
Silicon fluoride
Fluoro acid air | |
| Identifiers | |
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3D model (JSmol)
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| ECHA InfoCard | 100.029.104 |
PubChem CID
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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 | −95.0 °C (−139.0 °F; 178.2 K)[1][2] |
| Boiling point | −90.3 °C (−130.5 °F; 182.8 K)[1] |
| 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)[3] |
| 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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Infobox references
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Preparation edit
SiF
4 is a by-product of the production of phosphate fertilizers wet process production, resulting from the attack of HF (derived from fluorapatite protonolysis) on silicates, which are present as impurities in the phosphate rocks.[7] The hydrofluoric acid and silicon dioxide (SiO2) react to produce hexafluorosilicic acid:[7]
- 6 HF + SiO2 → H2SiF6 + 2 H2O
In the laboratory, the compound is prepared by heating barium hexafluorosilicate (Ba[SiF6]) above 300 °C (572 °F) whereupon the solid releases volatile SiF
4, leaving a residue of BaF
2.
- Ba[SiF6] + 400°C → BaF2 + SiF4
Alternatively, sodium hexafluorosilicate (Na2[SiF6]) may also be thermally decomposed at 400 °C (752 °F)—600 °C (1,112 °F) (optionally in inert nitrogen gas atmosphere) [8]: 8
- Na2[SiF6] + 400°C → 2NaF + SiF4
Uses edit
This volatile compound finds limited use in microelectronics and organic synthesis.[9]
It's also used in production of fluorosilicic acid (see above).[6]
Staying in the 1980s, as part of the Low-Cost Solar Array Project by Jet Propulsion Laboratory,[10] it was investigated as a potentially cheap feedstock for polycrystalline silicon production in fluidized bed reactors.[11] Few methods using it for the said production process were patented.[8][12]
The Ethyl Corporation process edit
In 80s the Ethyl Corporation came up with a process that uses hexafluorosilicic acid and sodium aluminium hydride (NaAlH4) (or other alkali metal hydride) to produce silane (SiH4).[13]
Occurrence edit
Volcanic plumes contain significant amounts of silicon tetrafluoride. Production can reach several tonnes per day.[14] Some amounts are also emitted from spontaneous coal fires.[15] The silicon tetrafluoride is partly hydrolysed and forms hexafluorosilicic acid.
Safety edit
In 2001 it was listed by New Jersey authorities as a hazardous substance that is corrosive and may severely irritate or even burn skin and eyes.[6] It is fatal if inhaled.[2]
See also edit
References edit
- ^ a b Silicon Compounds, Silicon Halides. Collins, W.: Kirk-Othmer Encyclopedia of Chemical Technology; John Wiley & Sons, Inc, 2001.
- ^ a b "SAFETY DATA SHEET: Silicon Tetrafluoride" (PDF). Airgas. April 9, 2018.
- ^ "Fluorides (as F)". Immediately Dangerous to Life or Health Concentrations (IDLH). National Institute for Occupational Safety and Health (NIOSH).
- ^ Greenwood & Earnshaw 1997, p. 328.
- ^ 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.
- ^ a b c "Hazardous Substance Fact Sheet" (PDF). New Jersey Department of Health and Senior services. November 2001.
- ^ a b Hoffman, C. J.; Gutowsky, H. S. (1953). "Germanium(IV) Fluoride". Inorganic Syntheses. Vol. 4. pp. 147–8. doi:10.1002/9780470132357.ch48. ISBN 978-0-470-13163-3.
- ^ a b Us Granted A345458, Keith, C. Hansen & L. Yaws, Carl, "Patent Silicon tetrafluoride generation", published January 3, 1982, issued 1982
- ^ Shimizu, M. "Silicon(IV) Fluoride" Encyclopedia of Reagents for Organic Synthesis, 2001 John Wiley & Sons. doi:10.1002/047084289X.rs011
- ^ Callaghan, William T. (1981). "Low-Cost Solar Array Project Progress and Plans". In Palz, W. (ed.). Photovoltaic Solar Energy Conference. Dordrecht: Springer Netherlands. pp. 279–286. doi:10.1007/978-94-009-8423-3_40. ISBN 978-94-009-8423-3.
- ^ Acharya, H. N.; Datta, S. K.; Banerjee, H. D.; Basu, S. (1982-09-01). "Low-temperature preparation of polycrystalline silicon from silicon tetrachloride". Materials Letters. 1 (2): 64–66. doi:10.1016/0167-577X(82)90008-8. ISSN 0167-577X.
- ^ CA 2741023A1, Anatoli, V. Pushko & Tozzoli, Silvio, "Method for the production of polycrystalline silicon", issued 2008
- ^ "The Ethyl Corporation Process: Silane and Fluidised Bed Reactor". August 11, 2015.
- ^ 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 (3): 249–256. Bibcode:2002EP&S...54..249M. doi:10.1186/BF03353024. S2CID 55173591.
- ^ Kruszewski, Ł., Fabiańska, M.J., Ciesielczuk, J., Segit, T., Orłowski, R., Motyliński, R., Moszumańska, I., Kusy, D. 2018 – First multi-tool exploration of a gas-condensate-pyrolysate system from the environment of burning coal mine heaps: An in situ FTIR and laboratory GC and PXRD study based on Upper Silesian materials. Science of the Total Environment, 640-641, 1044-1071; DOI: 10.1016/j.scitotenv.2018.05.319
- Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.