Silicon (and Boron)

On a completely random note, who can make elemental silicon? It is shiny and hard to make, and those are plenty of reasons for a mad scientist.

Of course you can do

SiO2 + C -> Si + CO2

but that takes 3000 C! Can this be achieved by thermite, burning coal/coke, microwave ovens, lasers, etc...? Any suggestions besides expensive electric arc furnaces?

[Edited on 19-5-2004 by Cyrus]

Aluminum Alloys

I've made silicon from the reaction of Mg with SiO2, I dissolved the mass I ended up with in HCl, it exploded

The reaction makes appreciable quanities of magnesium silicide which will react with acids to produce silane. If you try this add your solid to water then slowly drip in acid with stirring.

A full preparation of crystalline silicium

According to Jander & Blasius, prep. inorg. chem. (which I am coincidentially scanning), Si is made the following way:

90 g of fine, well-dried sand, 100 g of Al grains (0.1-0.5 mm), and 120 g of sulphur are mixed well, and set off with the various methods described first (KNO3/Al, NaClO3/Al, Mg ribbon, and my fav. sparklers stuck right into the mix).
After the reaction, let it cool down, and remove the product. Polverise/make chunks with a hammer or similar, and add a large amount of water. Large amounts of H2S are evolved (so do it outside, it's a toxic gas), and get rid off Al(OH3)/Al2O3 by flowing water accross it (the term is 'abschlaemmen', for which I don't have a translation), to remove the less dense Al(OH3)/Al2O3 suspension.
Separate the Si spheres mechanically from Al2O3 etc (hammer once again), and treat this with hand-warm medium concentrated HCl for several days (replacing used HCl a few times).
At last, cook the insoluble remainder with concentrated HCl (37%), and decant.
Then, to remove any residual SiO2, treat the silicon crystals with 40% H2F2 for one hour (warm again).


Clean Boron is made the same way, although one always gets Al containing boron (AlB12).


In terms of feasibility - I guess the only problem is the H2F2 - but, ammonium hydrofluoride can be obtained relatively easily in some hardware stores. Dont remember what it's used for though.
I am sure though this can be used just as well.

For clean crystalline Silicon- thats the way to go!

Thanks a lot, I will definitely try some of these. However, my propane burner barely melts aluminum, maybe I will make a larger one. Does anyone have some temperatures for these reactions?

As for uses, this is Mad science, not practical science.
(No practical use whatsoever for me)

Boron

Although it's slightly or considerably off-topic - but yes, the preparation is done similar to the silicon prep. With B2O3 of course.
Amounts are
50 g water-free B2O3, 75 g of S, and 100 g of Al. The cleaning and processing is done as above, i.e. first heating in conc HCl, then H2F2.

Now I don't see why you had trouble making boric acid. I made quite a large batch some time ago. Make sure your solutions are saturated with borax, and only then add HCl. If you cool, B(OH)3 should precipated. This, when you melt it, forms a glass like mass, B2O3. You can make very fine threads with it, which sting (because they are so thin) if you handle it improperly.

If you want to discuss boron properly, have a look in the forum first for existing threads, if not, then create one. Refer to this one to bring some sense to it

Lestat - sometimes you make compounds/elements for the sheer joy of it. That's what amateur experimentalism is about - not to make things to eat/inject/inhale or blow up, but to make things because you can! Who could possibly deny the sense of achievement of having made silicone or boron??

Silicon (and Boron)

Hehe, Boron is off topic no longer!
-am I allowed to do that?

I did have a look at the other threads, but found them insufficient...
The reason I did not get any B(OH)3 is because I never knew you cooled the soln. I should have tried it.
Do you understand the reaction though?

Na2B4O7 + HCl -> B(OH)3 + ???

Tomorrow, I make Boron.

Chemoleo, I assume you meant making silicon, not silicone, though silicone is nice too.

[Edited on 19-5-2004 by Cyrus]

Cyrus

No

My pottery supplier sells it in pure form as a very fine dust for about 4 Euro / kg. Try a pottery supplier.

For small amounts of glass powder, could you take sand and several steel balls, and shake them around in a film canister?

Well, I tried making boron yesterday.

First I added conc. HCl to warm conc. borax soln., in about 1 to 1 proportions,
I think this is the right reaction.
Na2B4O7 + 2HCl + 5 H20 -> 4B(OH)3 + NaCl
I then cooled the solution to near 0 C, and a semitransparant precipitate formed,
but there was so LITTLE of it that I was disgusted, I used about ten scoops of borax, and got less than one scoop of B(OH). Borax is a decahydrate though, would that increase the volume enormously? I will do the reaction again with better measurements later.

Then I took a small amount and dried it in the microwave making about 0.5 ml
of the boric acid powder, which smelled strongly of HCl, and heated it to 600+ C in a piece of Al foil, it turned black, bubbled, and ate a hole in the Al foil. (which was probably because of the HCl) I dissolved the mass in HCl, and a few black specks remain today. I claim that these are boron
Now to see if they are green in a flame test... more to follow.

Yes, I know the Al foil was not sufficient for a great reaction, but I wanted to see how it dehydrated...

This time I used at least 1 ml of boric acid, which I heated in a copper endcap... ahem crucible, it dehydrated nicely, giving off BEAUTIFUL green light. I then finely ground up the B2O3 and mixed it with Al powder in a 1 to 2 ratio. Back it went
into the crucible, and I cooked it from all sides for a couple minutes, then took out the red-brown mass, crunched it up,
and put the junk in a watch glass.
How can I determine if it is boron?
Did the reaction you did have sulfur in it?
I don't have any on hand right now

Yes, it is an antiseptic, contact solutions have boric acid in them. It also is good at killing ants and termites.

I like the green flame idea... All of my future pyro experiments are in grave danger of having B2O3 added to them. The color is so magnificent!

On a lesser note, perhaps candles could be made with wax impregnated with B2O3
for a neat effect.

How do you seperate B(OH)3 and NaCl?

Ah yes, that is what I did, but NaCl and B(OH)3 both precipitated!

The trick must be to find the point at which B(OH)3 but no NaCl will precipitate, but this seems rather wasteful as SO much B(OH)3 is left in solution. Arrg!

What if you heated up the NaCl + boric acid ppt to drive off water and form diboron trioxide, and found a solvent that would dissolve NaCl, but not hydrate the diboron trioxide?

Edit::: Another idea- since the salt forms large crystals, while the boric acid forms a fine powder ppt, I could stir the suspension/mixture and decant off the boric acid leaving salt behind. To get rid of any salt left in the boric acid, recrystalize. Or run the junk through a fine mesh to remove most NaCl...


[Edited on 25-5-2004 by Cyrus]

Wow, I am surprised you are having problems to get clean H3BO3. IIRC I made over 500 g of it, nice and clean?!? I used borax just like you, and dissolved in the SMALLEST possible amount of water (hot), let it cool, and filter off any crystals. Then add the correct amount of HCl, cool, and filter rapidly. Wash with icecold water to get rid off NaCl, and recrystallise if you are worried about purity. The reason your crystals glow green/yellow is of course it is mainly boric acid, but covered with NaCl, not because NaCl crystals have precipitated too. THe solubility is greater, plus the concentration of NaCl is twice as low as boric acid.
THis you can dehydrate at high temps, until it melts, to get B2O3.

For making the methyl/ethyl ester, just mix acid & alcohol + a bit of acid IIRC, and distill. That way I once made 200 ml of boric trimethylester, which, when poured onto the street, engulfed it in green emerald flame ... that was one of the few experiments my parentals actually liked

PS of course I meant silicon, or silicium - I didnt realise there was a difference in spelling

PS 2 Ah, another way to get relatively clean SiO2 powder is to burn silicone, under conditions where the heat can build up. THis produces a greyish white powder, where the grey stuff is mainly carbon residue.
I once made a thermite with this reaction product, and it worked fine (see exotic thermite thread)

[Edited on 25-5-2004 by chemoleo]

I heated up the solid/soln. mix until it all redissolved, added some water for good measure, and waited.
A soft white ppt. formed, and after cooling, I filtered this out. It was very watery though.

I took a small sample, and heated it in my copper crucible- it dissolved in the water, boiled, hardened, and then glowed green and yellow.

I then took a very small sample and heated it in a watch glass in the microwave. First it dissolved in its own water and started boiling, then it started bubbling like bubble gum. The sample is sticky, white, and rubbery- gum This is very different than the last "boric acid" I made, which steamed in the microwave, and dried into a fine white powder.

What have I made THIS time?

EDIT::: well, the substance behaves JUST like borax in the microwave-try it.
However, it gives off emerald green flames in the crucible, I think it is boric acid.

Yes, chemoleo, having a problem with this simple of a reaction is sad, but I was trying to maximize yields by making the soln. as conc. as possible and by waiting for several days before filtering-I wasn't sure about the ppt. rate of boric acid.
This caused the NaCl ppt. Now I have about 60 grams of boric acid drying, thanks for the help guys.

Could you please provide more info on the boric trimethylester synthesis...
Your description was a little vague.
[Edited on 26-5-2004 by Cyrus]

[Edited on 27-5-2004 by Cyrus]

I am pretty sure that the ppt. is some hyrate of H3BO3, this will mess up calculations, the powder weighed more than three times as much as the projected yield of H3BO3,and the "dry" powder dissolves in its own water at about 75C. Definitely hydrated, so not nice and clean . Above 100C it dehydrates to HBO2, then to H2B4O7, so you can't just boil the water off without risking some H3BO3 changing to HBO2. Will putting the boric acid powder in a dessicator with CaCl2 work?

Quote:

Originally posted by chemoleo For making the methyl/ethyl ester, just mix acid & alcohol + a bit of acid IIRC, and distill. That way I once made 200 ml of boric trimethylester, which, when poured onto the street, engulfed it in green emerald flame ... that was one of the few experiments my parentals actually liked [Edited on 25-5-2004 by chemoleo]

Could you give some more info on this please? I am not quite an organic chem. expert .

Quote:

Originally posted by Ostwald There are ways to "purify" sand... I'd bake it in a furnace/burn it directly in a bunsen burner to make sure anything even remotely organic is gone. And then I'd take a magnet after it's cool and remove anything that's magnetic. Some places have a lot of magnetite mixed up with sand (like the shores of Lake Superior!).

You would also need to treat it with HCl to dissolve out any grains of carbonate minerals, which would be derived utlimately from corals or the shells of molluscs. Magnetite or other ferromagnetic minerals (e.g. haematite, ilmenite, limonite) occur where lava from basaltic or ultrabasic volcanic eruptions has decomposed due to weathering, and the products washed onto beaches. Even with these treatments, if derived from granite the sand may still contain rutile, perovskite, and resistant silicate minerals like zircon, orthoclase, plagioclase, and rare-earth silicates.

John W.

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12AX7 - 8-3-2005 at 12:57

*Bumps old thread*

Today I reacted sand and (molten) aluminum with molten salt (as solvent), got a pretty good yield although I think there's still plenty of sand left in the slag. Depending on how much excess aluminum is used, it makes a hypereutectic (i.e., silicon content crystallizes before aluminum; OBTW aluminum does not for a silicide). I'm hoping I have 50% (intended for master alloy) here, I need to assay it.

In theory, such bars could be ground up and the aluminum dissolved with acid to obtain 99.9% or better purity silicon (it has a smidgeon of solubility for aluminum, so would be extremely heavily doped (i.e. uselessly so) P type silicon, for those solid state heads out there wondering).

Tim

P.S. Hi guys!

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BromicAcid - 5-5-2005 at 19:40

Being that my attempt at making magnesium boride was basically the thermite reaction between boric oxide and magnesium it ended up becoming strictly the product of boron that was the goal in the end.

The black liquid remaining after adding HCl to the mix was placed into a beaker and fifty milliliters of HCl was added along with a stir bar. The mixture was brought to the boiling point and held there for twenty minutes and the resulting mixture was filtered immediatley (see attached picture [Note, the areas that look white are acutally shiny]).

The solid recovered was put into the empty beaker and covered with 100 ml distilled water and again brought to a boil making a turbid black solution. This too was filtered hot, the resulting solid did not look as shiny as the solid obtained after the first filtration, perhaps boric acid clinging to the particles enhanced their shinyness. Overall the reaction went well, yield was ~40% which is decent considering this was not the initial intended goal. Very nice powdery substance

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Blackout - 6-5-2005 at 06:40

http://www.freewebs.com/akexperimental/chemexperiments.htm
go down where it's write "Magnesium and sand"

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Cyrus - 8-5-2005 at 18:17

I just finished a little test reaction.

2Mg + SiO2 -> Si + 2MgO

10.00 grams Mg in the form of turnings (about 4 g) and larger chunks, which I figured would melt as soon as the reaction got going.

12.36 g SiO2 in the form of fine sand.

I mixed the materials together and placed them on a thin layer of extra sand in a raku clay crucible (1 or so mm thick, already cracked a lot because of previous (ab)use. ) Extra sand (except for in the middle of the pile) was also placed on the top of the mix to prevent the Mg from burning away too much.

I tried to light it using a Mg strip, which burned but wouldn't light the rest of the pile. So I heated the top of the mix for a minute or so with my propane torch and stirred it around a bit. It then lit, started making a bit of a hissing noise, and produced a bright white flame a 4 or so inches long- it was glowing orange-yellow hot about 15 seconds afterwards, so I know there was a good reaction going on.

(I was viewing the reaction through a piece of shade 10 welding glass taped onto the front of my goggles- a nice setup because I could look at the stuff through the uncovered part of the goggles, and when the reaction started, just move my head down a bit to look through the shaded glass. I would recommend, of course, some eye protection for these types of reaction, but y'all know that already...)


After letting the stuff cool, I broke it up (the top part I had stirred around was a mostly loose whitish (MgO) powder, while the bottom half had fused into a crumbly black mass.)
I seperated the top and bottom parts; both were crushed and covered with water- no reaction, and then a bit of 12M HCl was added to both; the white stuff heated up (MgO reacting, probably) and few bubbles came off. The black stuff heated up, and started releasing lots of bubbles. Every now some exploded, producing nice little flashes of light. I'm pretty sure that's silane. Also larger bubbles were produced that didn't explode, which I think were H2.

Anyhow, long post, but unless all of my silicon dissapears as silane, this is a ridiculously easy way to make silicon.

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Cyrus - 10-5-2005 at 17:54

Huh. Well after a day or so of digesting it in conc. HCl, I appear to have sand.
I'm positive that it is mostly sand (the clear hard grains kinda give it away) but the sand is a bit brown and there's also fine black "silt" type stuff, which I'm hoping is Si. Any simple tests to see if this is silicon, besides dissolving it in Al, which SiO2 won't do AFAIK, and then dissolving the Al in HCl to get Si?

I should note that when I did a similar reaction using sand, Al powder, and S, the silicon formed little shiny spheres, and a shiny crystalline powder. I think if I use finer sizes of reactants (ie powders) it will heat up the reaction enough to fuse everything and get some nice crystalline Si. (Or explode) Hey, adding a bit of flux, perhaps NaCl, will cause the silica to melt at lower temperatures. Or, I could just use 400 mesh silica...

Perhaps I dissolved my silicon by using too much conc. HCl?

Cyrus

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12AX7 - 10-5-2005 at 23:44

It could be. The master alloy I dissolved with HCl was completely gone. But that could be 5% Fe content forming silicides, though I noticed no spontaneous explosions. Maybe try something weaker, like...uh...sodium bisulfate?

I tried reduction of ~80-150 grit off-white blasting sand with magnalium tonight. Absolutely NOTHING. I even tried lighting it with a good hot blend of freshly calcined Fe2O3 (I like the Al/Mg mix, it makes the slag and iron ball up nicely), nothing. Heated the charge to redness with torch, bupkiss.

Tim

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Cyrus - 11-5-2005 at 15:15

I'm not sure that sodium bisulfate will get all of the Mg and MgO, but I've never worked with it.

Try using Mg turnings too; I think what happened was that as I stirred the mix some of the Mg turnings or chunks caught on fire, causing the reaction to start.

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12AX7 - 11-5-2005 at 15:29

I tried that a long while ago to no avail as well.

I mean c'mon, I used THERMITE to try lighting it... and it didn't work. Someone gimme a break!

When I was stirring and heating it, I could heat stirred-up peaks to redness and they'd burn down, leaving the silica obliviously intact...

Tim

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Cyrus - 11-5-2005 at 15:43

You have been given an official break.

And to make things more confusing, here's this!

"There are two main industrial technologies to produce magnesium: the thermic and electrolytic method. The thermic process utilizes silicon to reduce magnesium oxide to produce magnesium"

Taken from
~
http://digitalcommons.hil.unb.ca/dissertations/AAIMQ35555/
~

Huh. I'm pretty sure that it wasn't just my Mg burning that I saw; the whole mass was glowing bright yellow/orange after the reaction was done.

edit, I just finished another little test reaction; 6.34g Mg (turnings only, and they were put in a little blender for a few minutes, so they were somewhat smaller and powdered.) 7.84 g SiO2 (400 mesh silica) these were mixed together and placed in a Cu end cap "crucible". A piece of Mg ribbon was placed into the center of the mix (where the ribbon met the mix I added an air opening to let the ribbon keep burning while it contacted the mix) and then lit (as if I just let it sit there. ).

The ribbon actually ignited the rest of the mix, which heated up to a nice yellow heat, and producing a little flame. The skin of the mix after heating was white, and the rest a dark black. The whole thing was lightly fused together (just enough to be solid and easily crunched apart) and some shiny black crystals were barely evident here and there.

The copper end cap had a black copper oxide coating, and the pan the reaction was done on was browned around the crucible, and the block of wood under the pan was blackened around the crucible. I don't think that was caused just by Mg burning on the surface. One way to be really sure would be to do the reaction in an inert atmosphere... but I have no He on hand or anything.


[Edited on 12-5-2005 by Cyrus]

[Edited on 12-5-2005 by Cyrus]

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chemoleo - 11-5-2005 at 17:00

Remember this post in the Exotic thermites thread?

Quote:

Silicon Dioxide thermite Today this was tried, in stoichiometric proportions, i.e. 54 g Al and 180 g SiO2, according to 3 SiO2 + 2 Al --> 3 Si + 2 Al2O3. 1. SiO2 as in quartz sand, >0.1 mm grains 2. SiO2 as from pottery supplies (fine powder), seemed a little wet though 3. SiO2 as from pottery supplies, fine too, and dry, but purity not known. Using 400 mesh Al, none of the thermites worked! Not even 5 sparkling candles tied together, or direct ignition with a Bunsen burner, or a NaClO3/Al mixture (which is very bright and hot). None of them. I am quite baffled by these results, in the light of a method in Jander&Blasius on the preparation of silicone using quartz sand. Then, I remember using the reaction product from burning silicone (which is SiO2 with impurities), and that did seem to work sluggishly (see above). I would have assumed that using the pure substrate would yield better results, but this... hmm. Any ideas? One thing I will do is to 1) precipitate SiO2 from waterglass, dry and 2) dry the existing pottery supplies SiO2, and try again. But I very much feel that this won't work - having made so many thermites to this point, this seems the most reluctant one

Lateron I found that 200 mesh Al and superdry SiO2 powder would still not ignite, not with a torch, nor with NaClO3/Al or sparklers.

S.C. Wack then posted this

Quote:

Sulfur. Schlessinger writes of 90 g sand, 100 g Al powder, and 120 g S in a crucible which is in sand.

And this I did! All of the SiO2/sand and Al was combined, and the adjusted amount of S added, and filled into a flowerpot.
This indeed was ignitable by sparklers, and would burn with intense white light for many minutes! Very pyrotechnic and pretty.

The resulting glowing mass was allowed to cool, and left overnight in the wet grass - and the next day this had completely disintegrated into grey squishy powder.... with lumps in them. Guess what they were - silicon !! Crystalline at that (see pic), one can see the crystal faces glittering here and there. It does not dissolve in HCl, and is shiny on the surface! Ideally I'd have subjected it to HF treatment but didnt deem it worth it.


Pretty eh?

I suspect the larger agglomerates of Si were possible becuase the reaction was fairly large, giving finely divided Si time to agglomerate at high white heat. So the smaller the reaction, the harder it will be to produce large chunks of Si!

One word of warning though - an enormous amount of Al2S3 is produced - which happily reacts with H2O forming tons of H2S. Unfortunately I left a small test amount in front of a window, which completely reeked out the room overnight by air/H2S being blown in!


[Edited on 12-5-2005 by chemoleo]

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12AX7 - 11-5-2005 at 17:10

Yippie Now, where to find sulfur...(I scoured Home Despot, Ace, Wal-Mart and the local gardening supply and NONE of them have the least interesting fertilizers or other products!!! Did the midwest suddenly become California!?)

Tim

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Cyrus - 11-5-2005 at 21:23

Cool, chemoleo. I did the exact same reaction, except about 1/10 of the size, so my silicon chunks weren't as large.

I'd say it is a perfect reaction, but the H2S problem bothers me a lot, even with my small reaction, neighbors a few houses away were outside trying to find the source of that smell.

Do you think that the S is an integral part of the reaction, or does it combine with Al to get the mix heated up? I ask because it might be possible to ignite the mix with a bit of Al/S/SiO2 placed on top of a larger portion of just Al and SiO2....

Tim- you can get S at pet/feed stores.

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12AX7 - 11-5-2005 at 22:36

I suspect it's just to heat it up. Suppose I'll go try it mixed with iron thermite and try making some ferrosilicon...not like I don't already have 5 pounds of the stuff.

Iron is an impurity in aluminum alloys though.. maybe copper would work better? I'd go try that too, if I hadn't burned my Cu2O already.
(Damnit, I hate this Cu-Si phase diagram I found... it's in atomic percent, so the intermetallic at 77% Cu is like all of half weight!)

Tim

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chemoleo - 12-5-2005 at 05:10

Cyrus, I don't think the reaction of SiO2/Al (both fine mesh) is self-sustaining if the original mix is at RT (i.e. not white heat). I even tried to ignite it with burning Al powder (in NaClO3), and it wouldnt go.
So yes, the S is there to generate the heat necessary.
I will at some point try a really big reaction, to see if I am able to isolate big chunks of Si. Plus I may add a flux agent to encourage the agglomeration of liquid Si. Because altogether the yield still sucked, compared to the amount of SiO2 used originally.

As to the H2S - well of course you can take the reaction product and use it for making Na2S, or reductions of various org./anorg compounds. You don't need to let it go to waste necessarily!

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12AX7 - 12-5-2005 at 08:33

Sounds like the trick is to make stoichiometric mixes of Al+S and Al+SiO2 + CaF2 (about 10% as flux, it is fluorite after all!) then blend different proportions of the two mixtures to see which will sustain. Or other things... lead oxide may be a candidate, as it does not form silicides.

Tim

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Update-ish

12AX7 - 24-7-2005 at 20:47

The other day I mixed a handful of 325 mesh flint with magnalium, and I'll be damned, it ignited. Solid state reaction progressed through the pile, similar to B2O3 burning. Am currently dissolving Al/Mg O's to release Si/SiO2, which will then be dissolved in aluminum metal with help from some liquid salt flux.

Tim

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Cyrus - 25-7-2005 at 15:56

Toldjya so!

It does react. But I never got enough heat going for the silicon to fuse into globules- the silicon or silicide or whatever what formed stayed intimately mixed with the rest of the reaction products and leftover sand. By the way, dissolving the products in HCl did produce silane as far as I know. (syntheticish fruity smelling explosive gas). I never figured out how to seperate the unknown amount of Si powder formed from the SiO2 easily... I hope your method works.

Cyrus

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A sulfur-free silicon thermite

blogfast25 - 11-3-2008 at 11:28

I'm sure many of you here have experimented with thermite reactions, including one of the harder reactions, that of silicon dioxide and aluminium to produce silicon metal(loid) and alumina.

Although the reaction 3 SiO₂ + 4 Al ---> 3 Si + 2 Al₂O₃ is thermodynamically favourable (ΔH < 0), this reaction does not propagate by itself, presumably because the heat of reaction isn't enough to overcome the activation energy it needs. Straight mixes of a silica source and aluminium powder therefore fizzle out or can't be ignited.

To make the reaction self-sustainable, the most used method by backyard scientists (and fellow travellers) is to add a booster mix of aluminium powder and sulfur, which reacts according to 2 Al + 3 S ---> Al₂S₃ with great development of heat (ΔH ≈ -5.3 kJ/g of stoichiometric mix). This heat provides the missing activation energy and makes the reduction of silica to silicon with Al self-sustaining.

A typical sulfur boosted silicon thermite mixture is silica/Al/S = 100/111/133 (9:10:12) but I've also successfully used mixes much lower in S, such as 100/72/21.

Apart from yielding a self-sustaining reaction, S-boosted silicon thermites have also other advantages:

• Quite easy to ignite, using Mg ribbon (e.g.).
• The resulting slag is a mix of alumina and aluminium sulfide. The much lower MP of the sulfide (around 1,100 C) causes the slag mix to be more fluid than pure alumina, which freezes at around 2,000 C. This greatly helps slag/metal separation, as the slag/metal mix remains liquid longer, allowing it to collect in the bottom of the crucible and the metal to coalesce out. And a mix of alumina and aluminium sulfide is also much softer than pure, fused alumina, making the slag easier to break up mechanically.
• The alumina/aluminium sulphide slag mix reacts readily with water through hydrolysis of the sulphide: Al₂S₃ + 6 H₂O ---> 2 Al(OH)₃ + 3 H₂S. This breaks up the slag into a (stinky) hydrated alumina slurry (or mud), giving easy access to the metal globules.


You can find an example of a 300 g S-boosted silicon thermite at this blog post of mine.

But that's the good news and there's some bad news too: the aluminium sulfide is so prone to hydrolysis, that even the newly fused slag positively reeks of H₂S, in plain English: rotten eggs. Needless to say, adding water or a mineral acid to it, seriously aggravates the problem. Not only does H₂S stink terribly, it's also toxic and it's perceptible even in trace amounts.

(Tip: if you're going to treat an alumina/aluminium slag mix with water, use bleach instead of pure water: the hypochlorite will convert much of the H₂S to elemental sulfur, which is even recoverable).

So much for the long-winded intro, I guess.

In a nutshell, I got so fed up with the smell of rotten eggs, I decided to try and replace the S-booster mix with a sulfur-free system. I chose to investigate a potassium chlorate/Al mix, which reacts according to 2 Al + KClO₃ ---> Al₂O₃ + KCl with an estimated ΔH ≈ 9.5 kJ/g (of stoichio mix). I had used such mixes before for lighting thermites.

Initial tests with a silica/Al/KClO₃ = 100/72/27 mix showed clearly that the reaction proceeded self-sustainingly and that Si metal was formed, in a hard, porous alumina matrix. I gradually stepped up the amount of booster mix to 100/84/57 and later 100/96/81, to find that progressively more of the slag ends up at the bottom of the crucible (I used mostly 20 g mini batches for the development work) because of the increasingly high peak temperatures during the reactions.

Much of this development can be followed here and on subsequent pages at the ABYMC forum (where I post as Gert from England).

The main problem remains slag/metal separation, both in situ and after the reaction products have cooled down: the pure, fused alumina freezes up quickly into a very, very hard mass. 32 w% HCl doesn't even begin to dent it and forget about mechanical separation: this stuff is HARD!

I then proceeded to test calcium fluoride (Fluorite, CaF₂ as a potential flux, at 20 w% added to the promising 100/96/81 formulation (this then became 100/96/81/55 - with 55 the CaF₂ . Although it made a world of difference in the sense that much larger globules of Si metal form, the slag remains extremely hard and insensitive to HCl. Another test at 40 w% CaF₂ showed that at that level the reaction was being slowed down, probably due to adding so much inert material, and hence slag metal separation deteriorated again due to lower peak temperatures.

In the mean time I've tried borax (without success) and a glassy flux formulation (designed for fluxing copper or bronze) without positive results either.

My main hope now lies with Cryolite (Na₃AlF₆ which has a lower melting point than CaF₂ (about 1,000 against about 1,400 C) and I'm waiting for a delivery of this stuff.

Other ideas that have been floated include reducing the radiative losses by covering the crucible partially, thereby reaching even higher peak temperatures and postponing the freeze.

Any useful ideas or experiences would be appreciated...

[Edited on 11-3-2008 by blogfast25]

[Edited on 11-3-2008 by blogfast25]

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ScienceGeek - 11-3-2008 at 12:50

Very interesting post. It seems you been busy lately
My experience is the same as yours. Hate the smell of rotten eggs in the morning after having done a Silicon Oxide thermite.
I think all your ideas are/was promising, and using Cryolite does not seem like a bad one either.
Unfortunately, my experience with Thermite is limited. Just wanted to say: Nice research!

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chloric1 - 11-3-2008 at 15:22

Glad someone tried the KClO3/Al booster. I had seen this in a scholar article about obtaining pure vanadium from calcium metavanadate. The mixture had to be heated to like 600 or 800°C before it could be ignited. Two suggestions stated using a booster of KClO3/Al or V2O5/Al. This is very simular to the most modern form of titanium isolation where titania is mixed with molten CaCl2 and electrolyzed to yield pure Ti via an active Ca intermediate. The elemental Ca intermediate is also the active reducer in vanadium and tungsten by aluminothermic reactions.

In leu of the calcium factor, why not try calcium metasilicate mixed with aluminum and a KClO3/Al booster additive. You might try additional Calcium silicate to obtain calcium aluminosilicate which is a major ingredient to concrete. It will still be hard but perhaps more brittle.

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chemoleo - 11-3-2008 at 17:02

I'm not sure what the advantage of this is, after all you end up with a monstrously hard and inert korundum slag from which it is very hard to seperate elemental Si... whilst the sulfur method just wastes a bit of cheap S and somewhat more Al (which considering our quantities) is also relatively cheap...

I notice you are into metal casting, what purpose did you have in mind for the Si?

[Edited on 12-3-2008 by chemoleo]

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not_important - 11-3-2008 at 17:48

A potential, meaning I don't know if it happens to a significant extent, problem with using fluorides is:

2CaF2 + 5 Si => 2 CaSi2 + SiF4 (g)
2CaF2 + 3 Si => 2 CaSi + SiF4 (g)

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blogfast25 - 12-3-2008 at 06:57

Quote:

Originally posted by chemoleo I'm not sure what the advantage of this is, after all you end up with a monstrously hard and inert korundum slag from which it is very hard to seperate elemental Si... whilst the sulfur method just wastes a bit of cheap S and somewhat more Al (which considering our quantities) is also relatively cheap... I notice you are into metal casting, what purpose did you have in mind for the Si? [Edited on 12-3-2008 by chemoleo]

chemoleo, the main purpose here is to get rid of the S, so as to avoid the horribly smelling H₂S. Those who've run S-boosted thermites repeatedly will probably concur that that gas is a problem, especially when running larger thermites.

Optimising the flux system should yield a considerably softer slag. That's the objective of the investigation...

The connection to metal casting was an investigation into the possibility of producing Si/Al master alloys by means of thermite reactions.

Chlorate-boosted thermites also have potential for other difficult thermites, notably TiO₂ or B₂O₃ and possibly others that are hard to light and/or sustain. Good burns have been obtained both with titania and boron trioxide. But in the case of titania, no metal was observed.

Quote:

Originally posted by not_important A potential, meaning I don't know if it happens to a significant extent, problem with using fluorides is: 2CaF2 + 5 Si => 2 CaSi2 + SiF4 (g) 2CaF2 + 3 Si => 2 CaSi + SiF4 (g)

No, this does not appear to be a problem. CaF₂ is too stable for reduction by Si or Al.

Quote:

Originally posted by chloric1 In leu of the calcium factor, why not try calcium metasilicate mixed with aluminum and a KClO3/Al booster additive. You might try additional Calcium silicate to obtain calcium aluminosilicate which is a major ingredient to concrete. It will still be hard but perhaps more brittle.

chloric, Al would reduce Ca silicate to Si and CaO, as it does CaSO₄ (to alumina, Al sulphide and CaO). CaO has an even higher MP than alumina!

[Edited on 12-3-2008 by blogfast25]

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chloric1 - 12-3-2008 at 13:44

Yes but CaO is soluble in dilute HCl and Si is not. Even if you still have aluminum oxide slag, I would think it would be easier to work with after leaching calcium out. Hot NaOH will probably dissolve the alumina or you could add the mix to hot concentrated sulfuric acid to get aluminum sulfate. I think you need magnesium or calcium to get titanium metal. Magnesium with titania and 10 or 15% potassium perchlorate/Mg mixed in would probably do it.

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blogfast25 - 13-3-2008 at 01:45

Quote:

Originally posted by chloric1 Yes but CaO is soluble in dilute HCl and Si is not. Even if you still have aluminum oxide slag, I would think it would be easier to work with after leaching calcium out. Hot NaOH will probably dissolve the alumina or you could add the mix to hot concentrated sulfuric acid to get aluminum sulfate. I think you need magnesium or calcium to get titanium metal. Magnesium with titania and 10 or 15% potassium perchlorate/Mg mixed in would probably do it.

chloric:

I did in fact run a test using CaSO₄ as a booster (AND chlorate):

3 CaSO₄ + 8 Al ---> 3 CaS + 4 Al₂O₃ + ΔH

CaS also reacts with HCl and I could smell the H₂S. But the slag wasn't appreciably easier to handle...

Hot NaOH would dissolve the slag, given time and temperature, unfortunately it also dissolves Si, because the metalloid is amphoteric (like Al).

Hot 70 - 80% sulfuric acid would probably also dissolve the slag (again, given time and temp.) but it's hard to get hold of.

Regarding TiO₂, thermodynamically the reduction with Al is possible (but that doesn't mean it works in practice). Several Internet reports report a burning TiO₂/Al mixture, I've managed to get it to burn with chlorate booster and also with magnalium (50/50 Mg/Al) without booster. A work in progress, I'd say.

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chief - 13-5-2010 at 06:45

Quote: Originally posted by 12AX7

Sounds like the trick is to make stoichiometric mixes of Al+S and Al+SiO2 + CaF2 (about 10% as flux, it is fluorite after all!) then blend different proportions of the two mixtures to see which will sustain. Or other things... lead oxide may be a candidate, as it does not form silicides. Tim

I did this 2 days ago, but not with sulfur ...
==> Instead I used stoechiometric MnO2-Al-thermite for the exothermicity and ignition ...
==> and SiO2-Al-thermite for the Si ...

50-50-mixtures of the both thermites do burn, and quite hot ...

Results still are in the furnace, will have a look these days ... (abused one of my old furnaces, for thermal insulation and thereby slower cooling ... )

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Random - 8-7-2010 at 13:18

There is a video on youtube about extracting silicon from chips.

http://www.youtube.com/user/Naravoslovje#p/u/24/-H4rt8q1XtY

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blogfast25 - 8-7-2010 at 13:30

The method Chemoleo describes here:

http://www.sciencemadness.org/talk/viewthread.php?tid=2030#p...

... I've used many times. It's probably the easiest way for a backyard chemist to prepare some relatively pure silicon.

Boron is prepared in much the same way, although one obtains AlB12. There's a detailed procedure in Brauer's 'Inorganic Preparative Chemistry' (see library).

I'm not sure B2O3 would actually colour a fire work green: the spectrum of B contains a strong green line but B2O3 is very stable and wouldn't dissociate much in fire work conditions. AlB12 may be better suited for that purpose. Worth trying!

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a_bab - 8-7-2010 at 23:24

Just stumbled across this reaction while studying the silicone carbide (which itself has a fascinating manufacturing process, plus it looks great ). It is claimed that silicon carbide is decomposed to silicon by the silicon dioxide:

SiC+SiO2--->2Si + CO2


It looks too good to be true. Even if electric arc is needed it still looks nice giving the "clean" reaction and the "OTC-bility of the reagents". It may need less then 1000 degrees C.

Does anyone else heard about this? All I could find was about thin silicon films that could be formed on the surface of SiO2 crystals using this reaction.

[Edited on 9-7-2010 by a_bab]

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blogfast25 - 9-7-2010 at 07:00

Quote: Originally posted by a_bab

It looks too good to be true. Even if electric arc is needed it still looks nice giving the "clean" reaction and the "OTC-bility of the reagents". It may need less then 1000 degrees C. [Edited on 9-7-2010 by a_bab]

Sure does. Where did you stumble on it?

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a_bab - 9-7-2010 at 08:44

An old russian chemistry book. It was just a note but it attracted my attention.

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The WiZard is In - 9-7-2010 at 10:53

Quote: Originally posted by chemoleo

Remember this post in the Exotic thermites thread? Quote: Silicon Dioxide thermite Today this was tried, in stoichiometric proportions, i.e. 54 g Al and 180 g SiO2, according to 3 SiO2 + 2 Al --> 3 Si + 2 Al2O3. [snip] Lateron I found that 200 mesh Al and superdry SiO2 powder would still not ignite, not with a torch, nor with NaClO3/Al or sparklers. S.C. Wack then posted this Quote: Sulfur. Schlessinger writes of 90 g sand, 100 g Al powder, and 120 g S in a crucible which is in sand. And this I did! All of the SiO2/sand and Al was combined, and the adjusted amount of S added, and filled into a flowerpot. This indeed was ignitable by sparklers, and would burn with intense white light for many minutes! Very pyrotechnic and pretty. The resulting glowing mass was allowed to cool, and left overnight in the wet grass - and the next day this had completely disintegrated into grey squishy powder.... with lumps in them. Guess what they were - silicon !! Crystalline at that (see pic), one can see the crystal faces glittering here and there. It does not dissolve in HCl, and is shiny on the surface! Ideally I'd have subjected it to HF treatment but didnt deem it worth it. Pretty eh? I suspect the larger agglomerates of Si were possible becuase the reaction was fairly large, giving finely divided Si time to agglomerate at high white heat. So the smaller the reaction, the harder it will be to produce large chunks of Si! One word of warning though - an enormous amount of Al2S3 is produced - which happily reacts with H2O forming tons of H2S. Unfortunately I left a small test amount in front of a window, which completely reeked out the room overnight by air/H2S being blown in! [Edited on 12-5-2005 by chemoleo]

Mellor sez 6:149

K.A. Kühne used the thermite process with a mixture of 360 parts
silica, and 400 parts of aluminium ; he also used a mixture of
silica , 36 parts ; aluminium 40 parts ; and sulphur , 50 parts
in a fireclay crucible. Holleman and Sliper used a modification of
this process. Watts recommended using cryolite or felspar as a
flux. Gröppel, a basic aluminium silicate.

K.A. Kühne, German patent D.R.P. 147871, 1902 ; 179403, 1905.

Holleman and Sliper, Rec, Trav. Chim. Pays-Bas, 23. 381, 1904.

Watts, An Investigation of the Borides and Silicides, Madison, WI,
1906

Gröppel ?

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blogfast25 - 9-7-2010 at 12:59

This is indeed one of the old ways of producing technical grade Si. Good enough for alloying purposes I'd imagine: most impurities (apart from some Al) will simply surface as dross...

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The WiZard is In - 9-7-2010 at 16:25

Quote: Originally posted by Cyrus

I just finished a little test reaction. 2Mg + SiO2 -> Si + 2MgO 10.00 grams Mg in the form of turnings (about 4 g) and larger chunks, which I figured would melt as soon as the reaction got going. [snip]

This from Watt's Dictionary of Chemistry. 1902. If found it
in my print edition and then checked Google.com/books
from which this was cut.

Preparation. — Amorphous silicon.— 1. An intimate mixture of 10
g. Mg powder and 40 g. thoroughly dry sand is placed in a
testtube, of fairly thick glass, c. 2-3 cms. diameter and c. 15 cms.
long; the tube is heated throughout by a large flame, and then
the lower part is very strongly heated, when reduction quickly
occurs. If the tube is gradually moved downwards so that one part
is strongly heated after another, the whole of the SiO, is reduced
in a few minutes. The contents of the tube are shaken out,
pulverised, and treated with HCl Aq (1:2), the solid being added to
the acid in small successive quantities (to prevent explosion from
evolution of SiH4) (Gattermann, B. 22, 186).

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condennnsa - 10-7-2010 at 00:53

how do you guys make your SiO2 powders for these experiments? I thought of using expanded perlite, since it's 70-75% SiO2 (wikipedia) and it's really soft, easily pulverised with a coffee grinder. would it work?

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The WiZard is In - 10-7-2010 at 06:38

Quote: Originally posted by condennnsa

how do you guys make your SiO2 powders for these experiments? I thought of using expanded perlite, since it's 70-75% SiO2 (wikipedia) and it's really soft, easily pulverised with a coffee grinder. would it work?

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I wonder if Cab-O-Sil (fumed silica), a really-rally fine powder
wouldn't work. Excepting you would have to compact it.

White quartz sand reduced to a powder in a ball mill
or some such perhaps ....

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a_bab - 10-7-2010 at 13:14

What about the very OTC silicagel? Easy tu pulverize, although I woudn't stay while doing it.

Also, some large electrical fuses (>100 amps) contain pure white fine silicon dioxide.

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blogfast25 - 11-7-2010 at 05:20

Cab-O-Sil ('fumed' or 'pyrogenic' silica) works: slightly wet it, then dry it: it compacts no end, while staying very fine.

There's also pottery silica or sand. Grind down your sand in a granite mortar and pestle of maybe a ball grinder... Even ground glass works but it's not pure SiO2, of course...

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condennnsa - 21-8-2010 at 04:35

What about dissolving an aluminum alloy of high Si content in hydrochloric acid? Al, Mg, Fe should go in solution , while silicon is not attacked by HCl as far as I know... ?

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12332123 - 21-8-2010 at 05:07

@condennnsa

While pure silicon may not be attacked by HCl, I very much suspect that it would form silane if reacted in the form of an alloy.

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blogfast25 - 21-8-2010 at 05:31

If the alloy containes silicide(s) (it ain't necessarily so) then you get: metal silicide + acid ---> Metal + silane

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aliced25 - 29-8-2010 at 04:29

I was just thinking, in the eternal quest for the inorganic products derived from boron and silicon (in my case the trihalides/tetrahalides), it is possible in the case of silicon to proceed from ferrosilicon with chlorine (250-350'C) to get SiCl4, so what would happen if aluminium diboride was treated similarly (ie. hot tube with chlorine)? Would that give a mixture of AlCl3 and BCl3? If so, purification shouldn't be that hard, the BP of BCl3 (12.6'C - Wikipedia) is hell lower than that of the by-product AlCl3 (MP:192.4'C - Wikipedia).

Has anyone got any suggestions on why this would not work? It just seems that if we can reach the boron/silicon halides via a simple aluminothermic reaction, followed by chlorination, then that would bring them into the "obtanium" range.

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blogfast25 - 29-8-2010 at 07:38

Quote: Originally posted by aliced25

I was just thinking, in the eternal quest for the inorganic products derived from boron and silicon (in my case the trihalides/tetrahalides), it is possible in the case of silicon to proceed from ferrosilicon with chlorine (250-350'C) to get SiCl4, so what would happen if aluminium diboride was treated similarly (ie. hot tube with chlorine)? Would that give a mixture of AlCl3 and BCl3? If so, purification shouldn't be that hard, the BP of BCl3 (12.6'C - Wikipedia) is hell lower than that of the by-product AlCl3 (MP:192.4'C - Wikipedia). Has anyone got any suggestions on why this would not work? It just seems that if we can reach the boron/silicon halides via a simple aluminothermic reaction, followed by chlorination, then that would bring them into the "obtanium" range.

Possibly, but note that AlB2 is one of those 'strange' electron deficient compounds, anything can happen when messing with those.

Wiki doesn't mention chlorination of AlB2 as a route to BCl3, but it does mention reaction of AlB2 and HCl to AlCl3 and borane...

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Boric acid

Neil - 15-4-2011 at 12:34

Bit of a bump but USP Boric acid is available from most pharmacies saving you the trouble of washing the Na out of borax. It's sold for making douches.

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redox - 16-4-2011 at 16:57

Boric acid is also available in hardware stores as an insecticide.

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tahallium - 9-3-2020 at 18:25

What about the electrolysis of silicates and borates? Cuz the electrolysis of sodium zincate make zinc metal and I think zinc is more reactive than silicone right?

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