Silicon tetrafluoride
Silicon tetrafluoride
Silicon tetrafluoride
Names
IUPAC names
Tetrafluorosilane
Silicon tetrafluoride
Other names
Silicon fluoride
Fluoro acid air
Identifiers
3D model (JSmol)
ECHA InfoCard 100.029.104
RTECS number
  • VW2327000
UNII
UN number 1859
  • F[Si](F)(F)F
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
 toxic, corrosive
NFPA 704 (fire diamond)
NFPA 704 four-colored diamond
3
0
2
W
Lethal dose or concentration (LD, LC):
69.220 mg/m3 (rat, 4 hr)[3]
Safety data sheet (SDS) ICSC 0576
Related compounds
Other anions
 Silicon tetrachloride
Silicon tetrabromide
Silicon tetraiodide
Other cations
 Carbon tetrafluoride
Germanium tetrafluoride
Tin tetrafluoride
Lead tetrafluoride
Related compounds
 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

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]

Preparation

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

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

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

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

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

References

  1. 1 2 Silicon Compounds, Silicon Halides. Collins, W.: Kirk-Othmer Encyclopedia of Chemical Technology; John Wiley & Sons, Inc, 2001.
  2. 1 2 "SAFETY DATA SHEET: Silicon Tetrafluoride" (PDF). Airgas. April 9, 2018.
  3. ↑ "Fluorides (as F)". Immediately Dangerous to Life or Health Concentrations (IDLH). National Institute for Occupational Safety and Health (NIOSH).
  4. ↑ Greenwood & Earnshaw 1997, p. 328.
  5. ↑ 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.
  6. 1 2 3 "Hazardous Substance Fact Sheet" (PDF). New Jersey Department of Health and Senior services. November 2001.
  7. 1 2 Hoffman, C. J.; Gutowsky, H. S. (1953). Silicon Tetrafluoride. Inorganic Syntheses. Vol. 4. pp. 147–8. doi:10.1002/9780470132357.ch48.
  8. 1 2 Us Granted A345458, Keith, C. Hansen & L. Yaws, Carl, "Patent Silicon tetrafluoride generation", published January 3, 1982, issued 1982
  9. ↑ Shimizu, M. "Silicon(IV) Fluoride" Encyclopedia of Reagents for Organic Synthesis, 2001 John Wiley & Sons. doi:10.1002/047084289X.rs011
  10. ↑ 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.
  11. ↑ 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.
  12. ↑ CA 2741023A1, Anatoli, V. Pushko & Tozzoli, Silvio, "Method for the production of polycrystalline silicon", issued 2008
  13. ↑ "The Ethyl Corporation Process: Silane and Fluidised Bed Reactor". August 11, 2015.
  14. ↑ 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.
  15. ↑ 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
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