cnidocyte
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Making anhydrous AlCl3 - Al + HCl in polar aprotic solvent
I'm trying to think of a way to make anhydrous AlCl3 and all I can really think of is reacting Al with HCl in a polar aprotic solvent like DMSO. If
I'm not mistaken the DMSO should allow the HCl to ionize and oxidize the Al. What I don't know is how the DMSO will react. Will it react with any of
the reactants or products of this reaction?
I found a few threads on anhydrous AlCl3 production on this forum and most people seem to find it tricky so I'm guessing theres a hole in my theory
here.
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Eclectic
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No one seems willing to try Ethyl Ether....it makes a stable, distillable adduct with AlCl3.
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Sedit
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DMSO also forms an adduct that according to some patents provided by Uno1me IIRC a while back can be precipitated and dryed even if hydrated AlCl3 is
used. Distillation of the Adduct frees the anhydrous AlCl3 from the complex recovering your DMSO.
Knowledge is useless to useless people...
"I see a lot of patterns in our behavior as a nation that parallel a lot of other historical processes. The fall of Rome, the fall of Germany — the
fall of the ruling country, the people who think they can do whatever they want without anybody else's consent. I've seen this story
before."~Maynard James Keenan
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Sedit
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DMSO also forms an adduct that according to some patents provided by Uno1me IIRC a while back can be precipitated and dryed even if hydrated AlCl3 is
used. Distillation of the Adduct frees the anhydrous AlCl3 from the complex recovering your DMSO.
Knowledge is useless to useless people...
"I see a lot of patterns in our behavior as a nation that parallel a lot of other historical processes. The fall of Rome, the fall of Germany — the
fall of the ruling country, the people who think they can do whatever they want without anybody else's consent. I've seen this story
before."~Maynard James Keenan
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Chainhit222
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I would try this if I had dmso.
Can someone give me a procedure for doing this with ether, I will do it for you if you give me some details. I have a decent bit of anhydrous diethyl
ether.
[Edited on 21-9-2010 by Chainhit222]
The practice of storing bottles of milk or beer in laboratory refrigerators is to be strongly condemned encouraged
-Vogels Textbook of Practical Organic Chemistry
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Eclectic
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Pass dry HCl into aluminum foil pieces in dry ethyl ether in reflux apparatus fitted with drying tube. Maybe use a bit of I2 to kickstart it. If the
Aluminum dissolves and the ether gets warm that's a good indicator.
Vigorous mag stirring may help.
[Edited on 9-21-2010 by Eclectic]
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cnidocyte
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I'm gonna try it with DMSO once I can find DMSO. I hear they stock it in health food shops but I haven't found a health food shop that has it yet.
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Chainhit222
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Oh, peach tried that. I thought adduct meant something other then just gassing Al in a solvent.
http://www.sciencemadness.org/talk/viewthread.php?tid=14111&...
What I meant to ask, is what is sedit talking about using AlCl3 hydrate to make AlCl3 anhydrous. I know you can dehydrate the FeCl3 salt, but I have
no knowledge of doing this to AlCl3.
[Edited on 21-9-2010 by Chainhit222]
The practice of storing bottles of milk or beer in laboratory refrigerators is to be strongly condemned encouraged
-Vogels Textbook of Practical Organic Chemistry
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peach
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A moderator needs to step in and arm up the merge hammer on these threads, they're getting out of control.
As soon as I have some free time and get another receiver bend, I'll make up some ether and try that.
I have a load of DMSO. I think Len was talking about the drying the hydrate idea using a DMSO adduct in another thread on this.
You can usually dry chloride salts that won't survive baking in the atmosphere by doing it under HCl(g). I believe the problem with this and aluminium
trichloride is that it'll start subliming all over the place again. I'm not sure.
One possible option to all this BBQ and solvent chat would be to just make the hydrated form with a bit of hydrochloric, then try
dehydrating it under HCl(g) with heat and maybe a cold finger. It'll use a fair bit of hydrogen chloride (six moles of it to every one of hydrate
going in), but then, those hot BBQ tubes and wash bottles don't absorb it 100% anyway (and they're starting at 3 moles gas to one of aluminium,
assuming 100% efficiency).
If you can dry the hydrate and catch the anhydrous easily, that would seem to be the easiest option so far.
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Chainhit222
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Ah, one would need a nice cold finger for this job, that you can take apart down the middle and easily scrape it, not some test tube microscale BS.
And yes, there is like a plague of alcl3 threads haha
But then again, I think having 10 different threads makes the community more involved, more people participate if it is popular. Perhaps leave them
around for a while, then pool all the knowledge together in order to make another "by the members" publication for anhydrous AlCl3 production.
[Edited on 21-9-2010 by Chainhit222]
The practice of storing bottles of milk or beer in laboratory refrigerators is to be strongly condemned encouraged
-Vogels Textbook of Practical Organic Chemistry
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peach
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That might be a good idea, I do see the logic there, provided it gets pooled at some point and not left splattered.
Cold fingers, I always thought they weren't designed very well for actually subliming much material. They tend to have a finger that's about the same
size as the port they poke into so, when it's pulled out, all the goodness gets stuck in the hole.....
To get round that, stick it into a port that's way bigger than the taper and use expansion sockets to make it fit. Then you can get a serious haul of
booty off in one sublimation.
Thinking about the cold solvent methods. This reaction is exothermic. The heat, I believe, is used to overcome the activation energy and drive the
reaction forwards quicker for commercial production. People have already mentioned priming the reaction with iodine, mercury chloride and others to
pass through the AE without using heat, something I will try at some point.
Personally, I wanted to avoid I2 / HgCl2 etc, despite owning enough of both in very pure forms, since that doesn't help if you don't have those
things, and they're effectively contaminants if pure AlCl3 is what you're after.
I have also read things related to AlCl3 it's self starting the reaction. My thinking, which may be wrong, was... even when a substance is a long way
from the temperature stated as 'the correct place to be', not all the molecules within it are. Maxwell & Botlzmann say the temperature you measure
is merely the average of lots of others, distributed in a curve. Some (small) percentage of the molecules there will be at very low temperatures, some
will be at extremely high temperatures.
If it's true that AlCl3 can catalyze it's own formation when HCl(g) is being bubbled through a solvent, it may only need time for those higher energy
molecules (in possession of the energy needed to overcome AE) to form the catalytic amount of the chloride.
{edit}{edit}{edit}{edit}
Whilst we're discussing these, are you guys after AlCl3 aware that there are numerous other compounds that function as both hard and
soft Lewis acids?
Anhydrous aluminium iodide is easy, easy to prepare. It's also towards the softer end of hard. Which makes it good for avoiding a complete brownout
situation on multifunctional organics. A common one in chemistry labs is BF3, which is narsty stuff if only to handle. It's also more than capable of
ripping things to brown if not cooled with dry ice.
Iodine is very simple to deal with, but expensive and tricky for a lot of people to get. There are other options though.
For example, Iron (III) Chloride, PCB etchant, is a hard Lewis acid, but slightly weaker than AlCl3.
It comes in the hydrated, yellow nodules form from electronics stores. And, unlike the aluminium version, it doesn't sublime or boil at stupidly low
temperatures.
Drying it is all of blowing hydrogen chloride over it as it approaches the decomposition point.
I have a heap of that stuff, about half a kilo, so I might add that to the list of things to make a video of.
Even metal oxides feature Lewis acid sites. I believe this is somewhat linked to their catalytic capacity.
Soft lewis acids won't attack hard lewis bases at any appreciable rate, and vice versa. It's not the same and Bronsted-Lowrys where a strong acid will
easily attack a soft alkali. Which is why you need specific bands of softness and hardness to attack differing sites.
{edit}{edit}{edit}{edit}
[Edited on 22-9-2010 by peach]
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