Silane
Chemical compound (SiH4) From Wikipedia, the free encyclopedia
Chemical compound (SiH4) From Wikipedia, the free encyclopedia
Silane (Silicane) is an inorganic compound with chemical formula SiH4. It is a colorless, pyrophoric, toxic gas with a sharp, repulsive, pungent smell, somewhat similar to that of acetic acid.[5] Silane is of practical interest as a precursor to elemental silicon. Silane with alkyl groups are effective water repellents for mineral surfaces such as concrete and masonry. Silanes with both organic and inorganic attachments are used as coupling agents. They are commonly used to apply coatings to surfaces or as an adhesion promoter.[6]
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| Names | |||
|---|---|---|---|
| IUPAC name
Silane | |||
| Systematic IUPAC name
Silicane | |||
Other names
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| Identifiers | |||
3D model (JSmol) |
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| ChEBI | |||
| ChemSpider | |||
| ECHA InfoCard | 100.029.331 | ||
| 273 | |||
PubChem CID |
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| RTECS number |
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| UNII | |||
| UN number | 2203 | ||
CompTox Dashboard (EPA) |
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| Properties | |||
| H4Si | |||
| Molar mass | 32.117 g·mol−1 | ||
| Appearance | Colorless gas | ||
| Odor | Repulsive[1] | ||
| Density | 1.313 g/L[2] | ||
| Melting point | −185 °C (−301.0 °F; 88.1 K)[2] | ||
| Boiling point | −111.9 °C (−169.4 °F; 161.2 K)[2] | ||
| Reacts slowly[2] | |||
| Vapor pressure | >1 atm (20 °C)[1] | ||
| Conjugate acid | Silanium (sometimes spelled silonium) | ||
| Structure | |||
| Tetrahedral r(Si-H) = 1.4798 Å[3] | |||
| 0 D | |||
| Thermochemistry[4] | |||
Heat capacity (C) |
42.81 J/mol·K | ||
Std molar entropy (S⦵298) |
204.61 J/mol·K | ||
Std enthalpy of formation (ΔfH⦵298) |
34.31 kJ/mol | ||
Gibbs free energy (ΔfG⦵) |
56.91 kJ/mol | ||
| Hazards | |||
| Occupational safety and health (OHS/OSH): | |||
Main hazards |
Extremely flammable, pyrophoric in air, toxic | ||
| GHS labelling: | |||
  | |||
| Danger | |||
| H220 | |||
| P210, P222, P230, P280, P377, P381, P403, P410+P403 | |||
| NFPA 704 (fire diamond) | |||
| Flash point | Not applicable, pyrophoric gas | ||
| ~ 18 °C (64 °F; 291 K) | |||
| Explosive limits | 1.37–100% | ||
| NIOSH (US health exposure limits): | |||
PEL (Permissible) |
None[1] | ||
REL (Recommended) |
TWA 5 ppm (7 mg/m3)[1] | ||
IDLH (Immediate danger) |
N.D.[1] | ||
| Safety data sheet (SDS) | ICSC 0564 | ||
| Related compounds | |||
Related tetrahydride compounds |
Methane Germane Stannane Plumbane | ||
Related compounds |
Phenylsilane Vinylsilane Disilane Trisilane | ||
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Silane can be produced by several routes.[7] Typically, it arises from the reaction of hydrogen chloride with magnesium silicide:
It is also prepared from metallurgical-grade silicon in a two-step process. First, silicon is treated with hydrogen chloride at about 300 °C to produce trichlorosilane, HSiCl3, along with hydrogen gas, according to the chemical equation
The trichlorosilane is then converted to a mixture of silane and silicon tetrachloride:
This redistribution reaction requires a catalyst.
The most commonly used catalysts for this process are metal halides, particularly aluminium chloride. This is referred to as a redistribution reaction, which is a double displacement involving the same central element. It may also be thought of as a disproportionation reaction, even though there is no change in the oxidation number for silicon (Si has a nominal oxidation number IV in all three species). However, the utility of the oxidation number concept for a covalent molecule[vague], even a polar covalent molecule, is ambiguous.[citation needed] The silicon atom could be rationalized as having the highest formal oxidation state and partial positive charge in SiCl4 and the lowest formal oxidation state in SiH4, since Cl is far more electronegative than is H.[citation needed]
An alternative industrial process for the preparation of very high-purity silane, suitable for use in the production of semiconductor-grade silicon, starts with metallurgical-grade silicon, hydrogen, and silicon tetrachloride and involves a complex series of redistribution reactions (producing byproducts that are recycled in the process) and distillations. The reactions are summarized below:
The silane produced by this route can be thermally decomposed to produce high-purity silicon and hydrogen in a single pass.
Still other industrial routes to silane involve reduction of silicon tetrafluoride (SiF4) with sodium hydride (NaH) or reduction of SiCl4 with lithium aluminium hydride (LiAlH4).
Another commercial production of silane involves reduction of silicon dioxide (SiO2) under Al and H2 gas in a mixture of NaCl and aluminum chloride (AlCl3) at high pressures:[8]
In 1857, the German chemists Heinrich Buff and Friedrich Woehler discovered silane among the products formed by the action of hydrochloric acid on aluminum silicide, which they had previously prepared. They called the compound siliciuretted hydrogen.[9]
For classroom demonstrations, silane can be produced by heating sand with magnesium powder to produce magnesium silicide (Mg2Si), then pouring the mixture into hydrochloric acid. The magnesium silicide reacts with the acid to produce silane gas, which burns on contact with air and produces tiny explosions.[10] This may be classified as a heterogeneous[clarification needed] acid–base chemical reaction, since the isolated Si4− ion in the Mg2Si antifluorite structure can serve as a Brønsted–Lowry base capable of accepting four protons. It can be written as
In general, the alkaline-earth metals form silicides with the following stoichiometries: MII2Si, MIISi, and MIISi2. In all cases, these substances react with Brønsted–Lowry acids to produce some type of hydride of silicon that is dependent on the Si anion connectivity in the silicide. The possible products include SiH4 and/or higher molecules in the homologous series SinH2n+2, a polymeric silicon hydride, or a silicic acid. Hence, MIISi with their zigzag chains of Si2− anions (containing two lone pairs of electrons on each Si anion that can accept protons) yield the polymeric hydride (SiH2)x.
Yet another small-scale route for the production of silane is from the action of sodium amalgam on dichlorosilane, SiH2Cl2, to yield monosilane along with some yellow polymerized silicon hydride (SiH)x.[11]
Silane is the silicon analogue of methane. All four Si−H bonds are equal and their length is 147.98 pm.[12] Because of the greater electronegativity of hydrogen in comparison to silicon, this Si–H bond polarity is the opposite of that in the C–H bonds of methane. One consequence of this reversed polarity is the greater tendency of silane to form complexes with transition metals. A second consequence is that silane is pyrophoric — it undergoes spontaneous combustion in air, without the need for external ignition.[13] However, the difficulties in explaining the available (often contradictory) combustion data are ascribed to the fact that silane itself is stable and that the natural formation of larger silanes during production, as well as the sensitivity of combustion to impurities such as moisture and to the catalytic effects of container surfaces causes its pyrophoricity.[14][15] Above 420 °C (788 °F), silane decomposes into silicon and hydrogen; it can therefore be used in the chemical vapor deposition of silicon.
The Si–H bond strength is around 384 kJ/mol, which is about 20% weaker than the H–H bond in H2. Consequently, compounds containing Si–H bonds are much more reactive than is H2. The strength of the Si–H bond is modestly affected by other substituents: the Si–H bond strengths are: SiHF3 419 kJ/mol, SiHCl3 382 kJ/mol, and SiHMe3 398 kJ/mol.[16][17]
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While diverse applications exist for organosilanes, silane itself has one dominant application, as a precursor to elemental silicon, particularly in the semiconductor industry. The higher silanes, such as di- and trisilane, are only of academic interest. About 300 metric tons per year of silane were consumed in the late 1990s.[needs update][15] Low-cost solar photovoltaic module manufacturing has led to substantial consumption of silane for depositing hydrogenated amorphous silicon (a-Si:H) on glass and other substrates like metal and plastic. The plasma-enhanced chemical vapor deposition (PECVD) process is relatively inefficient at materials utilization with approximately 85% of the silane being wasted. To reduce that waste and the ecological footprint of a-Si:H-based solar cells further several recycling efforts have been developed.[18][19]
A number of fatal industrial accidents produced by combustion and detonation of leaked silane in air have been reported.[20][21][22]
Due to weak bonds and hydrogen, silane is a pyrophoric gas (capable of autoignition at temperatures below 54 °C or 129 °F).[23]
For lean mixtures a two-stage reaction process has been proposed, which consists of a silane consumption process and a hydrogen oxidation process. The heat of SiO2(s) condensation increases the burning velocity due to thermal feedback.[24]
Diluted silane mixtures with inert gases such as nitrogen or argon are even more likely to ignite when leaked into open air, compared to pure silane: even a 1% mixture of silane in pure nitrogen easily ignites when exposed to air.[25]
In Japan, in order to reduce the danger of silane for amorphous silicon solar cell manufacturing, several companies began to dilute silane with hydrogen gas. This resulted in a symbiotic benefit of making more stable solar photovoltaic cells as it reduced the Staebler–Wronski effect.[citation needed]
Unlike methane, silane is fairly toxic: the lethal concentration in air for rats (LC50) is 0.96% (9,600 ppm) over a 4-hour exposure. In addition, contact with eyes may form silicic acid with resultant irritation.[26]
In regards to occupational exposure of silane to workers, the US National Institute for Occupational Safety and Health has set a recommended exposure limit of 5 ppm (7 mg/m3) over an eight-hour time-weighted average.[27]
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