 | |||
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| Names | |||
|---|---|---|---|
| IUPAC name Silane | |||
| Systematic IUPAC name Silicane | |||
Other names
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| Identifiers | |||
3D model (JSmol) | |||
| ChEBI | |||
| ChemSpider |
| ||
| ECHA InfoCard | 100.029.331 | ||
| 273 | |||
| RTECS number |
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| UNII | |||
| UN number | 2203 | ||
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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] | |||
| 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 | ||
| GHS labelling: | |||
 | |||
| Danger | |||
| H220[5] | |||
| 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 theirstandard state (at 25 °C [77 °F], 100 kPa). | |||
Silane (Silicane) is aninorganic compound withchemical formulaSiH4. It is a colorless,pyrophoricgas with a sharp, repulsive,pungent smell, somewhat similar to that ofacetic acid.[6] Silane is of practical interest as a precursor to elementalsilicon. Silanes withalkyl groups are effective water repellents for mineral surfaces such as concrete and masonry. Silanes with bothorganic andinorganic attachments are used as coupling agents. They are commonly used to apply coatings to surfaces or as an adhesion promoter.[7]
Silane can be produced by several routes.[8] Typically, it arises from the reaction of hydrogen chloride withmagnesium silicide:
It is also prepared from metallurgical-grade silicon in a two-step process. First, silicon is treated withhydrogen chloride at about 300 °C to producetrichlorosilane, HSiCl3, along withhydrogen gas, according to thechemical equation
The trichlorosilane is then converted to a mixture of silane andsilicon tetrachloride:
Thisredistribution reaction requires a catalyst.
The most commonly used catalysts for this process aremetalhalides, particularlyaluminium 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 adisproportionation 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 inSiCl4 and the lowest formal oxidation state inSiH4, 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, andsilicon 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 ofsilicon tetrafluoride (SiF4) withsodium hydride (NaH) or reduction ofSiCl4 withlithium aluminium hydride (LiAlH4).
Another commercial production of silane involves reduction ofsilicon dioxide (SiO2) under Al andH2 gas in a mixture ofNaCl andaluminum chloride (AlCl3) at high pressures:[9]
In 1857, the German chemistsHeinrich Buff andFriedrich Woehler discovered silane among the products formed by the action ofhydrochloric acid on aluminum silicide, which they had previously prepared. They called the compoundsiliciuretted hydrogen.[10]
For classroom demonstrations, silane can be produced by heatingsand withmagnesium powder to producemagnesium silicide (Mg2Si), then pouring the mixture into hydrochloric acid. The magnesium silicide reacts with the acid to produce silane gas, whichburns on contact with air and produces tiny explosions.[11] This may be classified as aheterogeneous[clarification needed]acid–base chemical reaction, since the isolatedSi4− ion in theMg2Siantifluorite structure can serve as aBrønsted–Lowry base capable of accepting four protons. It can be written as
In general, the alkaline-earth metals form silicides with the followingstoichiometries:MII2Si,MIISi, andMIISi2. 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 includeSiH4 and/or higher molecules in the homologous seriesSinH2n+2, a polymeric silicon hydride, or asilicic acid. Hence,MIISi with their zigzag chains ofSi2− 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 ofsodium amalgam ondichlorosilane,SiH2Cl2, to yield monosilane along with some yellowpolymerized silicon hydride(SiH)x.[12]
Silane is thesiliconanalogue ofmethane. All fourSi−H bonds are equal and their length is 147.98pm.[13] 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 ispyrophoric — it undergoes spontaneous combustion in air, without the need for external ignition.[14] 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.[15][16] Above 420 °C (788 °F), silane decomposes into silicon andhydrogen; it can therefore be used in thechemical vapor deposition of silicon.
The Si–Hbond strength is around 384 kJ/mol, which is about 20% weaker than the H–H bond inH2. Consequently, compounds containing Si–H bonds are much more reactive than isH2. 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.[17][18]
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While diverse applications exist fororganosilanes, 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 300metric tons per year of silane were consumed in the late 1990s.[needs update][16] Low-costsolar photovoltaic module manufacturing has led to substantial consumption of silane for depositing hydrogenatedamorphous silicon (a-Si:H) on glass and other substrates like metal and plastic. Theplasma-enhanced chemical vapor deposition (PECVD) process is relatively inefficient at materials utilization with approximately 85% of the silane being wasted. To reduce the waste andecological footprint of a-Si:H-based solar cells further, several recycling efforts have been developed.[19][20]
A number of fatal industrial accidents produced by combustion and detonation of leaked silane in air have been reported.[21][22][23]
Silane is a pyrophoric gas (capable of autoignition at temperatures below 54 °C or 129 °F).[24]
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 ofSiO2(s) condensation increases the burning velocity due to thermal feedback.[25]
Diluted silane mixtures with inert gases such asnitrogen orargon 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.[26]
In Japan, in order to reduce the danger of silane for amorphous silicon solar cell manufacturing, several companies began to dilute silane withhydrogen gas. This resulted in a symbiotic benefit of making more stablesolar photovoltaic cells as it reduced theStaebler–Wronski effect.[citation needed]
Unlike methane, silane is slightly 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 formsilicic acid with resultant irritation.[27]
In regards to occupational exposure of silane to workers, the USNational Institute for Occupational Safety and Health has set arecommended exposure limit of 5 ppm (7 mg/m3) over an eight-hour time-weighted average.[28]
