 | |
 | |
| Names | |
|---|---|
| IUPAC names Tetrafluorosilane Silicon tetrafluoride | |
| Other names Silicon fluoride Fluoro acid air | |
| Identifiers | |
3D model (JSmol) | |
| ChemSpider | |
| ECHA InfoCard | 100.029.104 |
| EC Number |
|
| RTECS number |
|
| UNII | |
| 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 | −95.0 °C (−139.0 °F; 178.2 K)[2][3] |
| Boiling point | −90.3 °C (−130.5 °F; 182.8 K)[2] |
| Critical point (T,P) | −14.15 °C (6.5 °F; 259.0 K), 36.71 standard atmospheres (3,719.6 kPa; 539.5 psi)[1] |
| decomposes | |
| Structure | |
| tetrahedral | |
| 0D | |
| Hazards | |
| Occupational safety and health (OHS/OSH): | |
Main hazards | toxic, corrosive |
| GHS labelling:[1] | |
   | |
| Danger | |
| H280,H300,H310,H314,H330,H331 | |
| P260,P261,P262,P264,P270,P271,P280,P284,P301+P316,P301+P330+P331,P302+P352,P302+P361+P354,P304+P340,P305+P354+P338,P316,P320,P321,P330,P361+P364,P363,P403+P233,P405,P410+P403,P501 | |
| NFPA 704 (fire diamond) | |
| Lethal dose or concentration (LD, LC): | |
LCLo (lowest published) | 69.220 mg/m3 (rat, 4 hr)[4] |
| Safety data sheet (SDS) | ICSC 0576 |
| Related compounds | |
Otheranions | Silicon tetrachloride Silicon tetrabromide Silicon tetraiodide |
Othercations | Carbon tetrafluoride Germanium tetrafluoride Tin tetrafluoride Lead tetrafluoride |
Related compounds | Hexafluorosilicic acid |
Except where otherwise noted, data are given for materials in theirstandard state (at 25 °C [77 °F], 100 kPa). | |
Silicon tetrafluoride ortetrafluorosilane is achemical compound with the formulaSiF4. 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 byCarl Wilhelm Scheele by dissolving silica inhydrofluoric acid,[5] and later synthesized byJohn Davy in 1812.[6] It is a tetrahedral molecule and is corrosive.[7]
Volcanic plumes contain significant amounts of silicon tetrafluoride. Production can reach several tonnes per day.[8] Some amounts are also emitted from spontaneous coal fires.[9] The silicon tetrafluoride is partly hydrolysed and formshexafluorosilicic acid.
SiF
4 is a by-product of the production ofphosphatefertilizerswet process production, resulting from the attack ofHF (derived fromfluorapatite protonolysis) onsilicates, which are present as impurities in thephosphate rocks.[10] Thehydrofluoric acid and silicon dioxide (SiO2) react to producehexafluorosilicic acid:[10]
In the laboratory, the compound is prepared by heatingbarium hexafluorosilicate (Ba[SiF6]) above 300 °C (572 °F) whereupon the solid releases volatileSiF
4, leaving a residue ofBaF
2.
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)[11]: 8
Thisvolatile compound finds limited use in microelectronics andorganic synthesis.[12]
It is also used in production offluorosilicic acid (see above).[7]
Staying in the 1980s, as part of the Low-Cost Solar Array Project byJet Propulsion Laboratory,[13] it was investigated as a potentially cheap feedstock forpolycrystalline silicon production influidized bed reactors.[14] Few methods using it for the said production process were patented.[11][15]
In 80s theEthyl Corporation came up with a process that uses hexafluorosilicic acid and sodiumaluminium hydride (NaAlH4) (or otheralkali metal hydride) to producesilane (SiH4).[16]
In 2001 it was listed byNew Jersey authorities as a hazardous substance that is corrosive and may severely irritate or even burn skin and eyes.[7] It is fatal if inhaled.[3]
{{cite book}}:ISBN / Date incompatibility (help)