12AX7 - 15-3-2005 at 17:05
I thought Si was completely insoluble in everything but HF... and yet, I cut two slabs off a piece of high-Si aluminum alloy I made, one I dissolved
with HCl and the other is almost done in room-temp lye solution.
The HCl was uncontrolled, so it boiled a lot as it dissolved. But that shouldn't react with Si, should it?
At any rate, the lye has black stuff in the bottom (various silicon crystals broken free of the aluminum matrix), yet the HCl solution is completely
clear. What th..?
Additionally, isn't Si soluble in NaOH? Shouldn't Si and/or the oxide make water glass?
Tim
neutrino - 15-3-2005 at 17:16
Your alloy was probably aluminum with aluminum silicide in it. Adding an acid would do the following:
Al<sub>4</sub>Si<sub>3</sub> + 12HCl -> 3SiH<sub>4(g)</sub> + 4AlCl<sub>3</sub>
Did the gas escaping from the reaction spontaneously ignite? This would be a sure sign of large quantities of silane (it’s pyrophoric).
[Edited on 16-3-2005 by neutrino]
12AX7 - 15-3-2005 at 17:52
http://web.met.kth.se/dct/pd/element/Al-Si.html
Aluminum does not form silicides. There is likely large iron contamination (5-10%), probably present as Al13Fe4, Fe2Si or a tertiary compound, but I
didn't notice any pyrophoria so there couldn't be much silane produced.
Is tetrachlorosilane soluble in water? Hrm, it hydrolyzes, doesn't it?
Tim
neutrino - 15-3-2005 at 19:40
You are right, aluminum doesn’t have any stable silicides. Is there anything else in the aluminum that might be reacting? I know that some
commercial forms of the stuff have a small amount of magnesium added…although this would only be a small amount and probably wouldn’t account for
all of the silicon reacting.
The HCl might be reacting with the silicon directly. Reacting HCl gas with silicon at high temperatures will give you various siliane chlorides but
silicon isn’t listed as reacting with HCl at room temperature. Possibly, this is one of those reactions that only progresses a small distance into a
piece of silicon, thus being labeled as not occurring at all (like oxygen attacking aluminum at room temperature). Coupled with the large surface area
that this procedure (eating away the surrounding aluminum with an acid) would give you, this would easily explain the reaction.
12AX7 - 15-3-2005 at 20:45
The aluminum feedstock is almost pure (99% or so). You can take a 1" square bar of this, hold one end, bash the other against the sidewalk and
just watch it deform effortlessly. I'm sure it dissolved quite a bit of iron from the crucible, which ought to form silicides.
The acid solution was slightly green, while the alkaline solution was always clear or white.
Hmm, a quick Google says SiCl4 hydrolyzes, but at least one page says it's soluble. But the reaction product is SiO2, which is insoluble in HCl,
so if it ever did form it would hydrolyze and precipitate. It must've boiled off, I guess.
Si strikes me as a reactive [semi-]metal. It's just a little less reactive than aluminum after all. It seems to me it would prefer to
dissociate water or acid molecules, being stopped by self-passivation (like aluminum is in most situations). But the high surface area, combined with
the oxide-free aluminum-silicon interfaces probably dissolved it.
On the other hand, I've done this to other samples, with cool, dilute acid, and it made a wonderfully black structure of silicon crystals like it
should. Hum...
[Edit] Oh, and I neutralized both solutions. (Kinda fun watching the equilibrium teeter back and forth as I add acid or base to get the most
precipitate. 
) Seems to have a suprising yield of Al(OH)3, considering I only
started with a 12 gram chunk of metal.
I'll test both ppts for silica, and of course weigh the silicon that didn't dissolve in the lye to find the percentage of the alloy.
Tim
[Edited on 16-3-2005 by 12AX7]
Theoretic - 17-3-2005 at 13:49
Aluminium silicides exist. Just type it in google and you'll see.
neutrino - 17-3-2005 at 14:05
If you look at the pages about aluminum silicon alloys, you will note that aluminum silicide is mentioned only as a metastable chemical which only
exists while the alloy is molten.
12AX7: Is it possible that you have impurities in your silicon? The behavior you describe seems the opposite of what is should be. Also, do aluminum
and silicon form a true alloy? If so, you certainly wouldn’t get a mass of silicon crystals.
12AX7 - 17-3-2005 at 18:44
Of all the articles on aluminum metallurgy, I've never read anything about a silicide but evidently the term does appear... but mostly in
reference to semiconductors. My quick search didn't find anything terribly useful so I'll take your word for it.
The ingots cooled relatively slowly (as slow as a 1 sq. in. cross section will in a steel angle iron) so they are most likely pretty much equilibrium.
I'm not sure what you mean by "true alloy", at any rate, these questions are answered on the phase diagram.
Tim
JohnWW - 20-3-2005 at 14:34
The Al-Si phase diagram is given, for example, on page 400 of Elements Of Material;s Science & Engineering, by L.H. Van Vlack (Addison-Wesley,
corrected 5th edition 1987), as adapted from the Metals Handbook (American Society for Metals). A less clear version is given at http://www2.umist.ac.uk/material/research/intmic/phase/alsid... .
It shows an eutectic point at 577ºC, at 12.6% silicon by weight, corresponding to 7.22 moles Al to 1 of Si. The silicide Al4Si3 would correspond to
56.156% Al, but there is absolutely no indication of any definite intermetallic-compound phase (which would be denoted by a solid phase boundary) at
that composition.