Silane

Production

Commercial-scale routes

Silane can be produced by several routes. 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, even a polar covalent molecule, is ambiguous. The silicon atom could be rationalized as having the highest formal oxidation state and partial positive charge in and the lowest formal oxidation state in , since Cl is far more electronegative than is H.

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 () with sodium hydride (NaH) or reduction of with lithium aluminium hydride ().

Another commercial production of silane involves reduction of silicon dioxide () under Al and gas in a mixture of NaCl and aluminum chloride () at high pressures:

:

Laboratory-scale routes

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.

For classroom demonstrations, silane can be produced by heating sand with magnesium powder to produce magnesium silicide (), 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. This may be classified as a heterogeneous acid–base chemical reaction, since the isolated ion in the 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: , , and . 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 and/or higher molecules in the homologous series , a polymeric silicon hydride, or a silicic acid. Hence, with their zigzag chains of anions (containing two lone pairs of electrons on each Si anion that can accept protons) yield the polymeric hydride .

Yet another small-scale route for the production of silane is from the action of sodium amalgam on dichlorosilane, , to yield monosilane along with some yellow polymerized silicon hydride .

Properties

Silane is the silicon analogue of methane. All four bonds are equal and their length is 147.98 pm. 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.{{ cite journal |name-list-style=amp |title = The oxidation of the silicon hydrides cite journal

The Si–H bond strength is around 384 kJ/mol, which is about 20% weaker than the H–H bond in . Consequently, compounds containing Si–H bonds are much more reactive than is . The strength of the Si–H bond is modestly affected by other substituents: the Si–H bond strengths are: 419 kJ/mol, 382 kJ/mol, and SiHMe3 398 kJ/mol.

Applications

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.{{ cite journal

Safety and precautions

A number of fatal industrial accidents produced by combustion and detonation of leaked silane in air have been reported.{{ cite journal cite journal cite journal

Silane is a pyrophoric gas (capable of autoignition at temperatures below 54 °C). :\Delta H = -1517 \text{ kJ/mol } = -47.23 \text{ kJ/g} : : : : 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 condensation increases the burning velocity due to thermal feedback.{{ cite journal

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.{{ cite journal

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.

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.

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.

References

Cited sources

References

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