CAS number 7803-62-5 YesY
PubChem 23953
ChemSpider 22393 YesY
UN number 2203
ChEBI CHEBI:29389 YesY
RTECS number VV1400000
Gmelin Reference 273
Jmol-3D images Image 1
Molecular formula H4Si
Molar mass 32.12 g mol−1
Exact mass 32.008226661 g mol−1
Appearance Colourless gas
Density 1.342 g dm−3
Melting point

−185 °C, 88 K, -301 °F

Boiling point

−112 °C, 161 K, -170 °F

Molecular shape tetrahedral

r(Si-H) = 1.4798 angstroms

Dipole moment 0 D
Std enthalpy of
Standard molar
204.6 J mol−1 K−1
EU Index Not listed
Main hazards Extremely flammable, pyrophoric in air
NFPA 704
NFPA 704.svg
Flash point flammable gas
294 K (21 °C) (~70 °F)
Explosive limits 1.37–100%
U.S. Permissible
exposure limit (PEL)
5 ppm (ACGIH TLV)
Related compounds
Related monosilanes Phenylsilane


Related compounds Methane


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Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Silane (or monosilane) is a toxic, extremely flammable chemical compound with chemical formula SiH4. 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.[1]



Commercial-scale routes

Industrially, silane is produced from metallurgical grade silicon in a two-step process. In the first step, powdered silicon is reacted with hydrogen chloride at about 300 °C to produce trichlorosilane, HSiCl3, along with hydrogen gas, according to the chemical equation:

Si + 3 HCl → HSiCl3 + H2

The trichlorosilane is then boiled on a resinous bed containing a catalyst which promotes the formation of silane and silicon tetrachloride according to the chemical equation:

4 HSiCl3 → SiH4 + 3 SiCl4

The most commonly used catalysts for this process are metal halides, particularly aluminium chloride. This is referred to 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 SiCl4 and the lowest in SiH4 since Cl is far more electronegative than is H.

An alternative industrial 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:

Si + 2 H2 + 3 SiCl4 → 4 SiHCl3
2 SiHCl3 → SiH2Cl2 + SiCl4
2 SiH2Cl2 → SiHCl3 + SiH3Cl
2 SiH3Cl → SiH4 + SiH2Cl2

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 SiF4 with sodium hydride (NaH) or reduction of SiCl4 with lithium aluminum hydride (LiAlH4).

Laboratory-scale routes

For classroom demonstrations, silane can be produced by heating sand with magnesium powder to produce magnesium silicide (Mg2Si), then pouring the mixture into a 20% dilution in non-aqueous solution of hydrochloric acid. Caution: If silane contacts water, it will react violently. The magnesium silicide reacts with the acid to produce silane gas, which burns on contact with air and produces tiny explosions.[2] This may be classified as a heterogenous[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:

4 HCl + Mg2Si → SiH4 + 2 MgCl2

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.[3]


Silane is the silicon analogue of methane. Because of the greater electronegativity of hydrogen in comparison to silicon, in silane the hydrogen atoms have a partial negative charge and the silicon a positive charge. This Si-H bond polarity is the opposite of that observed in the C-H bonds of methane. However, the C-H bonds in methane are generally regarded as non-polar since carbon is only slightly more electronegative than hydrogen. At room temperature, silane is a gas, and is pyrophoric — it undergoes spontaneous combustion in air, without the need for external ignition.[4] 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.[5][6] Above 420 °C, silane decomposes into silicon and hydrogen; it can therefore be used in the chemical vapor deposition of silicon.

Silane has a repulsive smell.[7]

Silane has recently been shown to act as a superconductor under extremely high pressures (96 and 120 GPa), with a transition temperature of 17 K.[8]


Several industrial and medical applications exist for silane and functionalized silanes. For instance, silanes are used as coupling agents to adhere glass fibers to a polymer matrix, stabilizing the composite material. In other words, silane coats the glass fibers to create better adhesion to the polymer chain. They can also be used to couple a bio-inert layer on a titanium implant. Other applications include water repellents, masonry protection, control of graffiti,[9] applying polycrystalline silicon layers on silicon wafers when manufacturing semiconductors, and sealants. The semiconductor industry used about 300 metric tons per year of silane in the late 1990s.[6] More recently, a growth in low-cost solar panel manufacturing has led to substantial consumption of silane for depositing amorphous silicon on glass and other surfaces.

Silane is also used in supersonic combustion ramjets to initiate combustion in the compressed air stream. As it can burn using carbon dioxide as an oxidizer it is a candidate fuel for engines operating on Mars.[10] Since this reaction has some byproducts which are solid (silicon dioxide and carbon) it is applicable only to liquid-fuel rockets (with liquid carbon dioxide), ramjets, or other reaction engines.

Silane and similar compounds containing Si—H bonds are used as reducing agents in organic and organometallic chemistry.[11]

Safety and precautions

A number of fatal industrial accidents produced by detonation and combustion of leaked silane in air have been reported.[12][13][14] 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.[15] Unlike methane, silane is fairly toxic: the lethal concentration in air for rats (LC50) is 0.96% (9600 ppm) over a 4-hour exposure. In addition, contact with eyes may form silicic acid with resultant irritation.[16]

See also


  1. ^ J. W. Mellor, "A Comprehensive Treatise on Inorganic and Theoretical Chemistry," Vol VI, Longmans, Green and Co. (1947), p. 216.
  2. ^ Making Silicon from Sand, by Theodore Gray. Originally published in Popular Science magazine.
  3. ^ J. W. Mellor "A Comprehensive Treatise on Inorganic and Theoretical Chemistry, Vol. VI" Longmans, Green and Co. (1947) pp. 970–971.
  4. ^ Emeléus, H. J. and Stewart, K. (1935). "The oxidation of the silicon hydrides". Journal of the Chemical Society: 1182–1189. doi:10.1039/JR9350001182. 
  5. ^ Koda, S. (1992). "Kinetic Aspects of Oxidation and Combustion of Silane and Related Compounds". Progress in Energy and Combustion Science 18 (6): 513–528. doi:10.1016/0360-1285(92)90037-2. 
  6. ^ a b Timms, P. L. (1999). "The chemistry of volatile waste from silicon wafer processing". Journal of the Chemical Society – Dalton Transactions (6): 815–822. doi:10.1039/a806743k. 
  7. ^ CFC Startec properties of Silane
  8. ^ M. I. Eremets, I. A. Trojan, S. A. Medvedev, J. S. Tse, Y. Yao (2008). "Superconductivity in Hydrogen Dominant Materials: Silane". Science 319 (5869): 1506–1509. Bibcode 2008Sci...319.1506E. doi:10.1126/science.1153282. PMID 18339933. 
  9. ^ Graffiti protection systems
  10. ^ Zubrin, Robert (1996). The Case for Mars. NY: Touchstone. p. 203. ISBN 0-684-83550-9. 
  11. ^ Reductions of organic compounds using silanes
  12. ^ Chen, J. R. (2002). "Characteristics of fire and explosion in semiconductor fabrication processes". Process Safety Progress 21 (1): 19–25. doi:10.1002/prs.680210106. 
  13. ^ Chen, J. R.; Tsai, H. Y.; Chen, S. K.; Pan, H. R.; Hu, S. C.; Shen, C. C.; Kuan, C. M.; Lee, Y. C.; and Wu, C. C. (2006). "Analysis of a silane explosion in a photovoltaic fabrication plant". Process Safety Progress 25 (3): 237–244. doi:10.1002/prs.10136. 
  14. ^ Chang, Y. Y.; Peng, D. J.; Wu, H. C.; Tsaur, C. C.; Shen, C. C.; Tsai, H. Y.; and Chen, J. R. (2007). "Revisiting of a silane explosion in a photovoltaic fabrication plant". Process Safety Progress 26 (2): 155–158. doi:10.1002/prs.10194. 
  15. ^ Kondo, S.; Tokuhashi, K.; Nagai, H.; Iwasaka, M.; and Kaise, M. (1995). "Spontaneous Ignition Limits of Silane and Phosphine". Combustion and Flame 101 (1–2): 170–174. doi:10.1016/0010-2180(94)00175-R. 
  16. ^ See MSDS for silane.

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