Silicon tetrafluoride

From Wikipedia, the free encyclopedia
Jump to: navigation, search
Silicon tetrafluoride
Silicon tetrafluoride
Silicon tetrafluoride
Names
IUPAC names
Tetrafluorosilane
Silicon tetrafluoride
Other names
Silicon fluoride
Fluoro acid air
Identifiers
7783-61-1 YesY
Jmol-3D images Image
PubChem 24556
RTECS number VW2327000
UN number 1859
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
Main hazards toxic, corrosive
Safety data sheet ICSC 0576
EU Index Not listed
NFPA 704
Flammability code 0: Will not burn. E.g., water Health code 3: Short exposure could cause serious temporary or residual injury. E.g., chlorine gas Reactivity code 2: Undergoes violent chemical change at elevated temperatures and pressures, reacts violently with water, or may form explosive mixtures with water. E.g., phosphorus Special hazard W: Reacts with water in an unusual or dangerous manner. E.g., cesium, sodiumNFPA 704 four-colored diamond
Lethal dose or concentration (LD, LC):
69,220 mg/m3 (rat, 4 hr)[1]
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).
 YesY verify (what isYesY/N?)
Infobox references

Silicon tetrafluoride or Tetrafluorosilane is the chemical compound with the formula SiF4. This tetrahedral molecule is notable for having a remarkably narrow liquid range (its boiling point is only 4 °C above its melting point). It was first synthesized by John Davy in 1812.[2]

Preparation[edit]

SiF
4
is a by-product of the production of phosphate fertilizers, resulting from the attack of HF (derived from fluorapatite protonolysis) on silicates. 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 also be created by placing silicon dioxide in hydrofluoric acid using the following equation:

4HF + SiO2 → SiF4 + 2H2O

Uses[edit]

This volatile compound finds limited use in microelectronics and organic synthesis.[5]

Occurrence[edit]

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[edit]

  1. ^ "Fluorides (as F)". Immediately Dangerous to Life and Health. National Institute for Occupational Safety and Health (NIOSH). 
  2. ^ 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. 
  3. ^ Hoffman, C. J.; Gutowsky, H. S. "Silicon Tetrafluoride" Inorganic Syntheses McGraw-Hill: New York, Volume 4, pages 145-6, 1953.
  4. ^ Hoffman, C. J.; Gutowsky, H. S. "Germanium Tetrafluoride" Inorganic Syntheses McGraw-Hill: New York, Volume 4, pages 147-8, 1953.
  5. ^ Shimizu, M. "Silicon(IV) Fluoride" Encyclopedia of Reagents for Organic Synthesis, 2001 John Wiley & Sons. doi:10.1002/047084289X.rs011
  6. ^ 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.