| Pages:
1
2
3
4 |
Ostwald
Harmless
Posts: 1
Registered: 22-5-2004
Member Is Offline
Mood: No Mood
|
|
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!). And if you have soil sieves (find your local geology or soil department), you could separate
sand grains out by size. Which may help.
|
|
|
Cyrus
Hazard to Others
Posts: 397
Registered: 24-4-2004
Location: Ancient Persia
Member Is Offline
Mood: No Mood
|
|
How do you seperate B(OH)3 and NaCl?
I dissolved about 500g of borax in boiling water (the soln. was about 800ml) and added the correct amount of HCl, by a day later the beaker was filled
with a solid chunk of loose crystals, which glow both green and yellow/orange in a flame. Now I ought to have a mass of boric acid mixed with sodium
chloride.
To seperate them what should I do next?
|
|
|
S.C. Wack
bibliomaster
Posts: 2419
Registered: 7-5-2004
Location: Cornworld, Central USA
Member Is Offline
Mood: Enhanced
|
|
Add as much water as it takes to dissolve everything hot, or more, then cool. I see Merck says that any remaining HCl will increase sol. in water, so
hope there isn't any left. And the flame tests, rather sensitive, especially sodium. You may still get some yellow with your boric once they are
separated.
|
|
|
Cyrus
Hazard to Others
Posts: 397
Registered: 24-4-2004
Location: Ancient Persia
Member Is Offline
Mood: No Mood
|
|
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]
|
|
|
chemoleo
Biochemicus Energeticus
Posts: 3005
Registered: 23-7-2003
Location: England Germany
Member Is Offline
Mood: crystalline
|
|
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]
Never Stop to Begin, and Never Begin to Stop...
Tolerance is good. But not with the intolerant! (Wilhelm Busch)
|
|
|
Cyrus
Hazard to Others
Posts: 397
Registered: 24-4-2004
Location: Ancient Persia
Member Is Offline
Mood: No Mood
|
|
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]
|
|
|
Cyrus
Hazard to Others
Posts: 397
Registered: 24-4-2004
Location: Ancient Persia
Member Is Offline
Mood: No Mood
|
|
| Quote: | Originally posted by chemoleo
IIRC I made over 500 g of it, nice and clean?!?
|
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  .
|
|
|
BromicAcid
International Hazard
Posts: 3246
Registered: 13-7-2003
Location: Wisconsin
Member Is Offline
Mood: Rock n' Roll
|
|
I was thinking about making some more boron soon seeing as how I recently came across an abundance of magnesium.
But there is one step in the process that always goofs me up. Dehydrating the boric acid as many of us know, creates a hard glassy mass that has to
be pulverized before use. However, what about another method to dehydrate the acid?
How about making a suspension of boric acid powder in something inert, something that can be heated to a high temperature. What comes to mind at
first, vegtable oil, really doesn't get high enough but you see where I'm going with this. With efficent stirring and high temps it may
well dehydrate the boric acid to the oxide and keep it in a nice divided state. However nucleation on the smaller particulates may take place leading
to larger particle sizes overall, but I wouldn't think as large as I usually end up with from trying to crush B2O3 formed the regular way.
|
|
|
JohnWW
International Hazard
Posts: 2849
Registered: 27-7-2004
Location: New Zealand
Member Is Offline
Mood: No Mood
|
|
| 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.
|
|
|
12AX7
Post Harlot
Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline
Mood: informative
|
|
*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!
|
|
|
BromicAcid
International Hazard
Posts: 3246
Registered: 13-7-2003
Location: Wisconsin
Member Is Offline
Mood: Rock n' Roll
|
|
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
|
|
|
Blackout
Harmless
Posts: 25
Registered: 26-11-2004
Location: Canada
Member Is Offline
Mood: No Mood
|
|
http://www.freewebs.com/akexperimental/chemexperiments.htm
go down where it's write "Magnesium and sand"
\"Si vis pacem, para bellum.\"
\"If you wish for peace, prepare for war.\"
|
|
|
Cyrus
Hazard to Others
Posts: 397
Registered: 24-4-2004
Location: Ancient Persia
Member Is Offline
Mood: No Mood
|
|
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.
|
|
|
Cyrus
Hazard to Others
Posts: 397
Registered: 24-4-2004
Location: Ancient Persia
Member Is Offline
Mood: No Mood
|
|
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
|
|
|
12AX7
Post Harlot
Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline
Mood: informative
|
|
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
|
|
|
Cyrus
Hazard to Others
Posts: 397
Registered: 24-4-2004
Location: Ancient Persia
Member Is Offline
Mood: No Mood
|
|
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.
|
|
|
12AX7
Post Harlot
Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline
Mood: informative
|
|
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
|
|
|
Cyrus
Hazard to Others
Posts: 397
Registered: 24-4-2004
Location: Ancient Persia
Member Is Offline
Mood: No Mood
|
|
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]
|
|
|
chemoleo
Biochemicus Energeticus
Posts: 3005
Registered: 23-7-2003
Location: England Germany
Member Is Offline
Mood: crystalline
|
|
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]
Never Stop to Begin, and Never Begin to Stop...
Tolerance is good. But not with the intolerant! (Wilhelm Busch)
|
|
|
12AX7
Post Harlot
Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline
Mood: informative
|
|
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
|
|
|
Cyrus
Hazard to Others
Posts: 397
Registered: 24-4-2004
Location: Ancient Persia
Member Is Offline
Mood: No Mood
|
|
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.
|
|
|
12AX7
Post Harlot
Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline
Mood: informative
|
|
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
|
|
|
chemoleo
Biochemicus Energeticus
Posts: 3005
Registered: 23-7-2003
Location: England Germany
Member Is Offline
Mood: crystalline
|
|
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!
Never Stop to Begin, and Never Begin to Stop...
Tolerance is good. But not with the intolerant! (Wilhelm Busch)
|
|
|
12AX7
Post Harlot
Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline
Mood: informative
|
|
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
|
|
|
12AX7
Post Harlot
Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline
Mood: informative
|
|
Update-ish
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
|
|
|
| Pages:
1
2
3
4 |
![[*]](images/xpblue/default_icon.gif){kind=icon}
