Chemistry:Silane

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Short description: Chemical compound (SiH4)
Silane
Stereo structural formula of silane
Ball-and-stick model of silane
Spacefill model of silane
Names
IUPAC name
Silane
Systematic IUPAC name
Silicane
Other names
  • Monosilane
  • Silicon(IV) hydride
  • Silicon tetrahydride
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
273
RTECS number
  • VV1400000
UNII
UN number 2203
Properties
H4Si
Molar mass 32.117 g·mol−1
Appearance Colourless 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
204.61 J/mol·K
34.31 kJ/mol
56.91 kJ/mol
Hazards
Main hazards Extremely flammable, pyrophoric in air
Safety data sheet ICSC 0564
GHS pictograms GHS02: FlammableGHS04: Compressed Gas
GHS Signal word Danger
H220, H280
PP210Script error: No such module "Preview warning".Category:GHS errors, PP222Script error: No such module "Preview warning".Category:GHS errors, PP230Script error: No such module "Preview warning".Category:GHS errors, PP280Script error: No such module "Preview warning".Category:GHS errors, PP377Script error: No such module "Preview warning".Category:GHS errors, PP381Script error: No such module "Preview warning".Category:GHS errors, PP403Script error: No such module "Preview warning".Category:GHS errors, PP410+P403Script error: No such module "Preview warning".Category:GHS errors
NFPA 704 (fire diamond)
Flammability code 4: Will rapidly or completely vaporize at normal atmospheric pressure and temperature, or is readily dispersed in air and will burn readily. Flash point below 23 °C (73 °F). E.g. propaneHealth code 2: Intense or continued but not chronic exposure could cause temporary incapacitation or possible residual injury. E.g. chloroformReactivity code 3: Capable of detonation or explosive decomposition but requires a strong initiating source, must be heated under confinement before initiation, reacts explosively with water, or will detonate if severely shocked. E.g. hydrogen peroxideSpecial hazards (white): no codeNFPA 704 four-colored diamond
4
2
3
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]
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 their standard state (at 25 °C [77 °F], 100 kPa).
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Infobox references

Silane (Silicane) is an inorganic compound with chemical formula SiH
4. It is a colourless, pyrophoric, toxic gas with a sharp, repulsive, pungent smell, somewhat similar to that of acetic acid.[5] Silane is of practical interest as a precursor to elemental silicon. Silane 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. Silanes are commonly used to apply coatings to surfaces or as an adhesion promoter.[6]

Production

Commercial-scale routes

Silane can be produced by several routes.[7] Typically, it arises from the reaction of hydrogen chloride with magnesium silicide:

[math]\ce{ Mg2Si + 4 HCl -> 2 MgCl2 + SiH4 }[/math]

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

[math]\ce{ Si + 3 HCl -> HSiCl3 + H2 }[/math]

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

[math]\ce{ 4 HSiCl3 -> SiH4 + 3 SiCl4 }[/math]

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[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 in SiCl
4 and the lowest formal oxidation state in SiH
4, 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, 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:

[math]\ce{ Si + 2 H2 + 3 SiCl4 -> 4 SiHCl3 }[/math]
[math]\ce{ 2 SiHCl3 -> SiH2Cl2 + SiCl4 }[/math]
[math]\ce{ 2 SiH2Cl2 -> SiHCl3 + SiH3Cl }[/math]
[math]\ce{ 2 SiH3Cl -> SiH4 + SiH2Cl2 }[/math]

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 (SiF
4) with sodium hydride (NaH) or reduction of SiCl
4 with lithium aluminium hydride (LiAlH
4).

Another commercial production of silane involves reduction of silicon dioxide (SiO
2) under Al and H
2 gas in a mixture of NaCl and aluminum chloride (AlCl
3) at high pressures:[8]

[math]\ce{ 3 SiO2 + 6 H2 + 4 Al -> 3 SiH4 + 2 Al2O3 }[/math]

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.[9]

For classroom demonstrations, silane can be produced by heating sand with magnesium powder to produce magnesium silicide (Mg
2Si), 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.[10] This may be classified as a heterogeneous[clarification needed] acid–base chemical reaction, since the isolated Si4− ion in the Mg
2Si antifluorite structure can serve as a Brønsted–Lowry base capable of accepting four protons. It can be written as

[math]\ce{ 4 HCl + Mg2Si -> SiH4 + 2 MgCl2 }[/math]

In general, the alkaline-earth metals form silicides with the following stoichiometries: MII

2Si, MII
Si, and MII
Si
2. 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 SiH
4 and/or higher molecules in the homologous series Si
nH
2n+2, a polymeric silicon hydride, or a silicic acid. Hence, MII
Si 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 (SiH
2)
x.

Yet another small-scale route for the production of silane is from the action of sodium amalgam on dichlorosilane, SiH
2Cl
2, to yield monosilane along with some yellow polymerized silicon hydride (SiH)
x.[11]

Properties

Silane is the silicon analogue of methane. All four Si–H bonds are equal and their length is 147.98 pm.[12] 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.[13] 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.[14][15] Above 420 °C, 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 H2. Consequently, compounds containing Si–H bonds are much more reactive than is H2. 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.[16][17]

Applications

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. [needs update] [15] Low-cost solar photovoltaic module manufacturing has led to substantial consumption of silane for depositing (PECVD) hydrogenated amorphous silicon (a-Si:H) on glass and other substrates like metal and plastic. The 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.[18][19]

Safety and precautions

A number of fatal industrial accidents produced by combustion and detonation of leaked silane in air have been reported.[20][21][22]

Due to weak bonds and hydrogen, silane is a pyrophoric gas (capable of autoignition at temperatures below 54 °C or 129 °F).[23]

[math]\ce{ SiH4 + 2 O2 -> SiO2 + 2 H2O }[/math]     [math]\displaystyle{ \Delta H = -1517 \text{ kJ/mol } = -47.23 \text{ kJ/g} }[/math]
[math]\ce{ SiH4 + O2 -> SiO2 + 2 H2 }[/math]
[math]\ce{ SiH4 + O2 -> SiH2O + H2O }[/math]
[math]\ce{ 2 SiH4 + O2 -> 2 SiH2O + 2H2 }[/math]
[math]\ce{ SiH2O + O2 -> SiO2 + H2O }[/math]

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 SiO
2(s) condensation increases the burning velocity due to thermal feedback.[24]

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.[25]

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[citation needed].

Unlike methane, silane is fairly 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.[26]

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.[27]

See also

References

  1. ↑ 1.0 1.1 1.2 1.3 1.4 NIOSH Pocket Guide to Chemical Hazards. "#0556". National Institute for Occupational Safety and Health (NIOSH). https://www.cdc.gov/niosh/npg/npgd0556.html. 
  2. ↑ 2.0 2.1 2.2 2.3 Haynes, p. 4.87
  3. ↑ Haynes, p. 9.29
  4. ↑ Haynes, p. 5.14
  5. ↑ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8. 
  6. ↑ London, Gábor; Carroll, Gregory T.; Feringa, Ben L. (2013). "Silanization of quartz, silicon and mica surfaces with light-driven molecular motors: construction of surface-bound photo-active nanolayers" (in en). Organic & Biomolecular Chemistry 11 (21): 3477–3483. doi:10.1039/c3ob40276b. ISSN 1477-0520. PMID 23592007. http://xlink.rsc.org/?DOI=c3ob40276b. 
  7. ↑ Simmler, W.. "Ullmann's Encyclopedia of Industrial Chemistry". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a24_001. 
  8. ↑ Shriver and Atkins. Inorganic Chemistry (5th edition). W. H. Freeman and Company, New York, 2010, p. 358.
  9. ↑ Mellor, J. W. "A Comprehensive Treatise on Inorganic and Theoretical Chemistry", vol. VI, Longmans, Green and Co. (1947), p. 216.
  10. ↑ "Making Silicon from Sand". Popular Science. https://theodoregray.com/PeriodicTable/PopularScience/2005/10/1/index.html. .
  11. ↑ Mellor, J. W. "A Comprehensive Treatise on Inorganic and Theoretical Chemistry", vol. VI. Longmans, Green and Co. (1947), pp. 970–971.
  12. ↑ "Maintenance". NIST. 17 October 2019. https://cccbdb.nist.gov/exp2x.asp?casno=7803625&charge=0. 
  13. ↑ Emeléus, H. J.; Stewart, K. (1935). "The oxidation of the silicon hydrides". Journal of the Chemical Society: 1182–1189. doi:10.1039/JR9350001182. 
  14. ↑ 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. 
  15. ↑ 15.0 15.1 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. 
  16. ↑ M. A. Brook "Silicon in Organic, Organometallic, and Polymer Chemistry" 2000, J. Wiley, New York. ISBN:0-471-19658-4.
  17. ↑ "Standard Bond Energies". Michigan State University Organic Chemistry. https://www2.chemistry.msu.edu/faculty/reusch/virttxtjml/questions/bndenrgy.htm. 
  18. ↑ Briend P, Alban B, Chevrel H, Jahan D. American Air, Liquide Inc. (2009) "Method for Recycling Silane (SiH4)". US20110011129 , EP2252550A2 .
  19. ↑ Kreiger, M.A.; Shonnard, D.R.; Pearce, J.M. (2013). "Life cycle analysis of silane recycling in amorphous silicon-based solar photovoltaic manufacturing". Resources, Conservation and Recycling 70: 44–49. doi:10.1016/j.resconrec.2012.10.002. https://www.academia.edu/2310926. 
  20. ↑ 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. 
  21. ↑ Chen, J. R.; Tsai, H. Y.; Chen, S. K.; Pan, H. R.; Hu, S. C.; Shen, C. C.; Kuan, C. M.; Lee, Y. C. et al. (2006). "Analysis of a silane explosion in a photovoltaic fabrication plant". Process Safety Progress 25 (3): 237–244. doi:10.1002/prs.10136. 
  22. ↑ Chang, Y. Y.; Peng, D. J.; Wu, H. C.; Tsaur, C. C.; Shen, C. C.; Tsai, H. Y.; 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. 
  23. ↑ Silane MSDS
  24. ↑ V.I Babushok (1998). "Numerical Study of Low and High Temperature Silane Combustion". The Combustion Institute 27 (2): 2431–2439. doi:10.1016/S0082-0784(98)80095-7. 
  25. ↑ Kondo, S.; Tokuhashi, K.; Nagai, H.; Iwasaka, M.; 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. 
  26. ↑ "MSDS for silane". http://www.vngas.com/pdf/g97.pdf. 
  27. ↑ "Silicon tetrahydride". NIOSH Pocket Guide to Chemical Hazards. Centers for Disease Control and Prevention. April 4, 2011. https://www.cdc.gov/niosh/npg/npgd0556.html. 

Cited sources

External links



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