When was silane discovered?

Chemical compound (SiH4)

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This article is about the compound with chemical formula SiH. For the broader classes of compounds, see Silanes

Chemical compound

Silane (Silicane) is an inorganic compound with chemical formula SiH4. It is a colorless, 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. They 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:

Mg2Si + 4 HCl &#; 2 MgCl2 + SiH4

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

Si + 3 HCl &#; HSiCl3 + H2

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

4 HSiCl3 &#; SiH4 + 3 SiCl4

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 SiCl4 and the lowest formal oxidation state in SiH4, 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:

1. Si + 2 H2 + 3 SiCl4 &#; 4 SiHCl3

2. 2 SiHCl3 &#; SiH2Cl2 + SiCl4

3. 2 SiH2Cl2 &#; SiHCl3 + SiH3Cl

4. 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 silicon tetrafluoride (SiF4) with sodium hydride (NaH) or reduction of SiCl4 with lithium aluminium hydride (LiAlH4).

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:[8]

3 SiO2 + 6 H2 + 4 Al &#; 3 SiH4 + 2 Al2O3

Laboratory-scale routes

In , 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 (Mg2Si), 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 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.[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 (788 °F), 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

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 s.[needs update][15] 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.[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]

SiH4 + 2 O2 &#; SiO2 + 2 H2O

     

Δ H = &#;  kJ/mol  = &#; 47.23  kJ/g {\displaystyle \Delta H=-{\text{ kJ/mol }}=-47.23{\text{ kJ/g}}}

SiH4 + O2 &#; SiO2 + 2 H2

SiH4 + O2 &#; SiH2O + H2O

2 SiH4 + O2 &#; 2 SiH2O + 2 H2

SiH2O + O2 &#; SiO2 + H2O

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 SiO2(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

Cited sources

What is Silane

Silane is an inorganic compound with chemical formula, SiH4, making it a group 14 hydride. It is a colourless, pyrophoric gas with a sharp, repulsive smell, somewhat similar to that of acetic acid. Silane is of practical interest as a precursor to elemental silicon.

&#;Silanes&#; refers to many compounds with four substituents on silicon, including an organosilicon compound. Examples include trichlorosilane (SiHCl3), tetramethylsilane (Si(CH3)4), and tetraethoxysilane (Si(OC2H5)4).

Production

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Commercial-scale routes

See also: Direct process

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

Mg2Si + 4 HCl &#; 2 MgCl2 + SiH4

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:

Si + 3 HCl &#; HSiCl3 + H2

The trichlorosilane is then converted to a mixture of silane and silicon tetrachloride. This redistribution reaction requires a catalyst:

4 HSiCl3 &#; SiH4 + 3 SiCl4

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 SiCl4 and the lowest formal oxidation state in SiH4 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:

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).

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:

3 SiO2 + 6 H2 + 4 Al &#; 3 SiH4 + 2 Al2O3

Laboratory-scale routes

In , 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 (Mg2Si), 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[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.

Properties

Silane is the silicon analogue of methane. 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. 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. 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 in SiHF3, SiHCl3, and SiHMe3 are respectively 419, 382, and 398 kJ/mol

Applications

Several industrial and medical applications exist for silane and functionalized silanes. For instance, silanes are used as coupling agents to adhere fibers such as glass fibers and carbon fibers to certain polymer matrices, stabilizing the composite material. In other words, silane coats the glass fibers to create better adhesion to the polymer matrix. 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, 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 s. More recently, a growth in 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.

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.

Silane and similar compounds containing Si&#;H bonds are used as reducing agents in organic and organometallic chemistry.

Silane methacrylates are used in dentistry as part of tooth-colored, composite filling material. Silane methacrylates act as a coupling agent between the hard, silicate-based, ceramic filler and the organic, resin-based oligomer matrix.

Safety and precautions

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

If a leaking stream of silane is obstructed or confined, energy release due to combustion is more concentrated leading to increasing reaction speed and burning velocity &#; up to gas phase detonation and potentially severe damages. Silane may autoignite at under 54 °C (130 °F).

SiH4(g) + 2O2(g) &#; SiO2(s) + 2H2O(g) with ΔH = &#; kJ/mol = &#;47.23 kJ/g

Hazardous byproducts of combustion

SiH4(g) + O2(g) &#; SiO2(s) + 2H2(g)

SiH4(g) + O2(g) &#; SiH2O(s) + H2O(g)

2SiH4(g) + O2(g) &#; 2SiH2O(s) + 2H2(g)

SiH2O(s) + O2(g) &#; SiO2(s) + H2O(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 SiO2 (s) condensation increases the burning velocity due to thermal feedback.

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.

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 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.

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

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