Silane

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Silane (Silicane) is an inorganic compound with chemical formula Template:Chem2. It is a colorless, pyrophoric gas with a sharp, repulsive, pungent smell, somewhat similar to that of acetic acid.<ref>Template:Greenwood&Earnshaw2nd</ref> Silane is of practical interest as a precursor to elemental silicon. Silanes 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.<ref>Template:Cite journal</ref>

ProductionEdit

Commercial-scale routesEdit

Silane can be produced by several routes.<ref name=Ullmann>Template:Ullmann</ref> Typically, it arises from the reaction of hydrogen chloride with magnesium silicide:

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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

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The trichlorosilane is then converted to a mixture of silane and silicon tetrachloride:

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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 moleculeTemplate:Vague, even a polar covalent molecule, is ambiguous.Template:Citation needed The silicon atom could be rationalized as having the highest formal oxidation state and partial positive charge in Template:Chem2 and the lowest formal oxidation state in Template:Chem2, since Cl is far more electronegative than is H.Template: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 (Template:Chem2) with sodium hydride (NaH) or reduction of Template:Chem2 with lithium aluminium hydride (Template:Chem2).

Another commercial production of silane involves reduction of silicon dioxide (Template:Chem2) under Al and Template:Chem2 gas in a mixture of NaCl and aluminum chloride (Template:Chem2) at high pressures:<ref>Shriver and Atkins. Inorganic Chemistry (5th edition). W. H. Freeman and Company, New York, 2010, p. 358.</ref>

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Laboratory-scale routesEdit

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.<ref>Mellor, J. W. "A Comprehensive Treatise on Inorganic and Theoretical Chemistry", vol. VI, Longmans, Green and Co. (1947), p. 216.</ref>

For classroom demonstrations, silane can be produced by heating sand with magnesium powder to produce magnesium silicide (Template:Chem2), 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.<ref name="theodoregray">{{#invoke:citation/CS1|citation |CitationClass=web }}.</ref> This may be classified as a heterogeneous Template:Clarify acid–base chemical reaction, since the isolated Template:Chem2 ion in the Template:Chem2 antifluorite structure can serve as a Brønsted–Lowry base capable of accepting four protons. It can be written as

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In general, the alkaline-earth metals form silicides with the following stoichiometries: Template:Chem2, Template:Chem2, and Template:Chem2. 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 Template:Chem2 and/or higher molecules in the homologous series Template:Chem2, a polymeric silicon hydride, or a silicic acid. Hence, Template:Chem2 with their zigzag chains of Template:Chem2 anions (containing two lone pairs of electrons on each Si anion that can accept protons) yield the polymeric hydride Template:Chem2.

Yet another small-scale route for the production of silane is from the action of sodium amalgam on dichlorosilane, Template:Chem2, to yield monosilane along with some yellow polymerized silicon hydride Template:Chem2.<ref>Mellor, J. W. "A Comprehensive Treatise on Inorganic and Theoretical Chemistry", vol. VI. Longmans, Green and Co. (1947), pp. 970–971.</ref>

PropertiesEdit

Silane is the silicon analogue of methane. All four Template:Chem2 bonds are equal and their length is 147.98 pm.<ref name=cccbdb>Template:Cite journal</ref> 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.<ref>Template:Cite journal</ref> 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.<ref>Template:Cite journal</ref><ref name=timms/> Above Template:Convert, 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 Template:Chem2. Consequently, compounds containing Si–H bonds are much more reactive than is Template:Chem2. The strength of the Si–H bond is modestly affected by other substituents: the Si–H bond strengths are: Template:Chem2 419 kJ/mol, Template:Chem2 382 kJ/mol, and SiHMe₃ 398 kJ/mol.<ref>M. A. Brook "Silicon in Organic, Organometallic, and Polymer Chemistry" 2000, J. Wiley, New York. Template:ISBN.</ref><ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref>

ApplicationsEdit

File:Container 【 28T9 】 SINU 212002(1)---No,1 【 Pictures taken in Japan 】.jpg

Monosilane gas shipping containers in Japan.

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.Template:Update inline<ref name=timms>Template:Cite journal</ref> 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 the waste and ecological footprint of a-Si:H-based solar cells further, several recycling efforts have been developed.<ref>Briend P, Alban B, Chevrel H, Jahan D. American Air, Liquide Inc. (2009) "Method for Recycling Silane (SiH₄)". US20110011129, EP2252550A2 .</ref><ref>Template:Cite journal</ref>

Safety and precautionsEdit

A number of fatal industrial accidents produced by combustion and detonation of leaked silane in air have been reported.<ref>Template:Cite journal</ref><ref>Template:Cite journal</ref><ref>Template:Cite journal</ref>

Silane is a pyrophoric gas (capable of autoignition at temperatures below Template:Convert).<ref>Silane MSDS Template:Webarchive</ref>

Template:Chem2 Template:Spaces<math>\Delta H = -1517 \text{ kJ/mol } = -47.23 \text{ kJ/g}</math>

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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 Template:Chem2 condensation increases the burning velocity due to thermal feedback.<ref>Template:Cite journal</ref>

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.<ref>Template:Cite journal</ref>

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.Template: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 form silicic acid with resultant irritation.<ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref>

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.<ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref>