woelen
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Silanes, self-igniting gases
I did a really nice experiment and the only chemicals needed for the basic experiment are fine sand, magnesium powder and some dilute acid.
http://woelen.homescience.net/science/chem/exps/silane/index...
I liked the experiment very much, especially the self-igniting properties of the silane gas(mix).
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bfesser
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Very well done. Excellent pictures, as always.
This might be a nice one to demo in series with the preparation of silicon dioxide using sodium silicate and sodium bisulfite. And, of course, sodium silicate is easy enough to prepare.
[Edited on 5/3/09 by bfesser]
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Formatik
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Aluminium should also work in place of Mg, no? A bonus about this is that silane is not terribly toxic like other metalloid hydrides, e.g. AsH3, SbH3.
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kclo4
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Very Impressive! I forgot about this gas, It would be fun to collect a few mls of the gas in a testube, or a balloon in small amounts.
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UnintentionalChaos
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Quote: Originally posted by Formatik  | | Aluminium should also work in place of Mg, no? A bonus about this is that silane is not terribly toxic like other metalloid hydrides, e.g. AsH3, SbH3.
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Silicon-Aluminum alloys are not uncommon, and I'm fairly sure they don't burst into flames in acids.
I believe Tim (12AX7) has made some of this alloy (intended as a master alloy) by adding silica sand to molten aluminum, but I don't believe there
were any flammable results.
[Edited on 5-4-09 by UnintentionalChaos]
Department of Redundancy Department - Now with paperwork!
'In organic synthesis, we call decomposition products "crap", however this is not a IUPAC approved nomenclature.' -Nicodem
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kclo4
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Perhaps you are thinking boranes? IIRC they ignite in the air, and can be produced from Al and B2O3. I could be totally wrong about that though.
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Formatik
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The Al reaction might not work with SiO2, but apparently it works with silicates:
| Quote: | | Aluminium silicide is obtained when aluminium is melted with silicates. In this way a compound containing 70 per cent, silicon can be obtained. A
portion is dissolved in the aluminium, and another part seems to be mechanically mixed like graphite in pig iron. When aluminium silicide is treated
with hydrochloric acid, part of the silicon escapes as siliciuretted hydrogen, another passes into solution as silica, and a third part remains behind
as a black powder. |
| Quote: Originally posted by kclo4 | | Perhaps you are thinking boranes? IIRC they ignite in the air, and can be produced from Al and B2O3. I could be totally wrong about that though.
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You can get some impure aluminium borides that way:
| Quote: | | F. E. Weston and H. R. Ellis also examined the reduction of powdered B2O3 by aluminium powder (Trans. of the Faraday Society, 1907). The product was
Al2O3 and mixtures of boron with aluminium borides. |
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woelen
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This is a very good suggestion. This allows one to make REALLY finely divided SiO2 and then the product probably even is better. I''ll try the
reaction for making ultrafine SiO2, this seems easier than hitting sand with a hammer. I now used the finest available 'silver sand' from a potteries
supplier, but the material you can obtain from silicate most likely can be made into a finer powder.
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JohnWW
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The usual method for making B2H6 with a mixture of other boranes, containing 3-center bonds, is by firstly making Mg3B2 by reaction of B2O3 with
powdered Mg in sufficient quantity for the Mg to both remove all the oxygen as MgO and then form the boride, and then reacting the boride with an
aqueous acid. As with the silanes, they ignite spontaneously on contact with air; and may even react with nitrogen.
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woelen
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Isn't borane very water-sensitive (as opposed to silane)? I always had the impression that B2H6 reacts with water quickly, giving hydrogen and boric
acid. Only if the borane is stabilized (e.g. by adding hydride and breaking the 3 center 2 electron bonds and forming a non-electrondeficient
borohydride anion) it resists reaction with water. But in acids, such ions still are very easily hydrolysed and give hydrogen and boric acid.
BTW, I have experimented a little more with silane and it can be kept under water, provided it is not alkaline. You can keep it in an inverted test
tube, which is under tapwater with a few drops of acid added to the water.
A good way to do this is taking a syringe, with the Mg2Si/MgO mix put in the syringe, and then filling it completely with water. Then, all air and
most of the water must be pressed out, without crushing the pieces of Mg2Si/MgO. So, a little water is left (a layer of a few mm) in which the
Mg2Si/MgO particles reside. Then some hydrochloric acid is drawn into the syringe, assuring that no air is sucked in. Immediately, the tip of the
syringe must be blocked. The gas, which is formed pushes away the plunger. When the reaction stops, you can press the gas out of the syringe, into an
inverted test tube. Using this method I could collect several ml of gas in an inverted test tube. I think that the gas mix I obtained is a mix of SiH4
and H2 (and traces of higher silanes). When a bubble of this gas mix is bubbled out of the test tube, then it goes to the surface and there it gives a
fairly big flash of fire and a "BOOP" sound.
I now still have 15 ml or so of the gas, does anyone have any other interesting ideas of what to do with this? If there is a good idea, then I might
try to make a video of that reaction and add that to the website.
[Edited on 5-5-09 by woelen]
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Jor
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According to Brauer it reacts with HCl, although the book only mentions a simple equation. SiH3Cl is supposed to be formed.
It boils at 36,5C and is water reactive. It fumes in air.
Maybe it might be interesting to react the gas with anhydrous HCl gas (generated by NaCl+H2SO4), and if a liquid forms, you might be able to isolate
it and keep it. I wonder if it would react with glass.
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garage chemist
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Borane(s) is/are indeed water-sensitive, and borides and aqueous acids will not give any boranes, at least not acceptable yields.
Pure, dust-fine SiO2 can be bought from pottery/ceramics suppliers as "silica flour" or "quartz". This material isn't just sand, it's pegmatite
quartz, specially mined and finely ground.
It is usually supplied moist since the dust from the dry powder is dangerous (carcinogenic and causes silicosis).
This is without a doubt the most economic source of pure SiO2 for syntheses.
The SiO2 synthesis from unitednuclear using sodium silicate and an acid like sodium bisulfate is incorrect, it is missing the important step of
washing the product to get rid of the Na2SO4 byproduct.
Aluminum and SiO2 will not give a self-sustaining reaction since it is not exothermic enough.
With extra Al and added sulfur, however, this becomes a thermite-like reaction for the preparation of silicon.
Externally supplied heat could also be used to effect the complete reaction of Al with SiO2.
Al does not readily form silicides like magnesium does.
But the prepared Si could be used for interesting syntheses like that of SiCl4 (with chlorine gas in a heated tube) and silico-chloroform
(trichlorosilane) SiHCl3, with HCl gas in a heated tube, and also for melting together with Mg to give pure magnesium silicide.
SiH4 gas isn't that useful, I'm not aware of any interesting reactions. Elemental Si, on the other hand, is much more interesting.
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Formatik
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After a bit of digging, from Lehrbuch der anorganischen Chemie by Holleman,etc. I've found out that SiH4 mixed with SiHCl3 was first discovered from
F. Wöhler and H. Buff in 1857 by protolysis of aluminium silicide with HCl acid. Disilane (S2H6) mixed with SiH4 and higher silanes was done by H.
Moissan and S. Smiles in 1902 by protolysis of dimagnesium silicide. Wiki cites two references which state SiH4 gets its pyrophoricity from the higher
silanes and some other factors. So then, although aluminium silicide will get silane, it shouldn't give the ignition because no higher silanes form.
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Formatik
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Then again, the German wiki for silanes ("Silane") states Wöhler observed from his "silicon-containing aluminium" reacted with hydrochloric acid to
yield impure hydrogen, which gas reacted violently with air to explode. An actual Wöhler paper where he did this would answer my questions for sure.
But all references I've seen that mention this love to never cite it. Assuming it was described in a journal.
Edit: Found it. Lieb. Ann. 1857, 103, 218. They first talk about an electrolysis experiment yielding the gas and then they talk about how they tried after alot of
experimentation to prepare the silane chemically. The self-ignition they did observe was only from the electrolysis experiments where aluminium was
used as an anode.
They said they observed formation only in small amounts by solubilizing silicon-containing aluminium in dilute hydrochloric acid. Drying this H2 with
CaCl2, and then burning it left behind white silicic acid, and sometimes brownish Si. But they emphasize never were they able to obtain such a
significant concentration of silane that it would ignite itself. Even if they used an aluminium which was saturated with silicon, by melting it with
cryolite and alkali silicates.
[Edited on 10-6-2009 by Formatik]
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12AX7
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There are no silicides on the Al-Si phase diagram. Perhaps it is possible under certain circumstances (pressure, rapid cooling, catalysts?), but not
under equilibrium. (In the Fe-C phase diagram, cementite (Fe3C) does not officially appear because it's not an equilibrium product, however it is
commonly marked because it's so damn important, and does show up under regular conditions.)
The thermite reaction works fine, although without an accelerant it proceeds fairly slowly and yields fine silicon, probably coated with a hearty
layer of SiO2 or Al2O3 -- I wasn't ever able to dissolve any in molten aluminum, it just sat there on top. More fire can probably fix that (Si melts
around 1500C).
Tim
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