[8] 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.

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:[9] 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.

[11] 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.

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.

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.

[needs update][16] 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.

[19][20] A number of fatal industrial accidents produced by combustion and detonation of leaked silane in air have been reported.

[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.