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peach
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No, it behaves as a solid crystalline.
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anotheronebitesthedust
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Maybe the neoprene tubing is adding unforeseen variables into the reaction. DCM penetrates neoprene and it was mentioned earlier in the thread that
neoprene often contains sulfur. Is aluminum sulphide possible? Were all your reagents free of contaminants?
From Wiki:
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2 Al + 3 S → Al2S3
This colorless species has an interesting structural chemistry, existing in three different forms. The material is sensitive to moisture, hydrolyzing
readily to hydrated aluminium oxides/hydroxides.[1] The hydrolysis reaction also generates the odoriferous and toxic gas hydrogen sulfide (H2S).
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kmno4
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I have just done experiment with amalgamated Al and DCM.
I was interested if any reaction occurs.
Amalgamated Al foil wiped off HgCl2 solution (becoming immediately hot in contact with air) was immersed in DCM. Some small bubbles appeared (probably
traces of water).
After 12 hours at 20 C, there was no change in colour of DCM or Al, nor bubbles.
When Al foil was taken out of DCM, immediately started getting hot.
So:
Al is inert to DCM (under given conditions).
Al remains active in DCM (no some passive layer)
As they say - Experiment is a king 
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Sedit
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Nice work this goes nicely with what I want to attempt soon.
I tryed just recently to place a long glass tube on a flask filled with H2SO4 an NaCl. The tube was filled with some shredded Al foil and the HCl was
allowed to escape thru the top.
Oddly enough what I got was patches of Al turned white.  . There would be white
patch in the middle but the bottom was un reacted ect ect... it really makes little sense because it should be Al mostl reacted at the lowerend and
progressivly get less reacted as the HCl is used up.
It was just a random experiment and I intend to eliminate variables shortly.
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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ziqquratu
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Perhaps you could try this on a small scale and see if the evolved gas burns (indicating H2)? You'd have to do it in a closed container, I suspect -
charge with HCl gas, seal and stir for a while, then try to ignite it. Of course, it may not work, so a negative result would not be conclusive, but a
positive would give some useful information.
I think that you'll probably want to use a different solvent - DCM will usually react happily with active metals, so you'd be better of with something
else. I'd probably reach for ether, or THF, perhaps. Something polar but inert to the active Al, the HCl, and the AlCl3. I'm of two minds as to
whether EtOAc would be suitable. DMSO might work and would be nicely polar. DMF is probably out under the conditions. Obviously, you have to have
access to these solvents...
You know what might work, is acetic acid... very polar, dissolves everything well, shouldn't have any compatibility problems... I'd try this or an
ether. Someone let me know if I'm missing something here.
I also think you'd be aided by a little bit of water in your solvent, to assist in proton transfer - traces of water are often necessary in reactions
using active metals. Usually it only needs to be a few ppm, but in this instance, adding a little might be helpful. Shouldn't need it in AcOH,
though... and ether straight from the bottle should be damp enough. Or if you wanted to control the amount, saturate a little solvent by storing it in
a closed container over water, then taking a little of the organic layer and diluting it with dry solvent (in ether at 25 *C, the concentration of
water will be about 1.5% - w/w, I think).
By the way, it occurs to me, isn't AlCl3 soluble in DCM? If I'm right (which at this time of the morning...), this suggests that your precipitate is
NOT AlCl3.
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IrC
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OK but that starts me thinking about aluminum hydride again.
"Science is the belief in the ignorance of the experts" Richard Feynman
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Sedit
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This page offers the suggestion that DCM is viable solvent for AlCl3, I honestly beg to differ when free halogen is used to produce Alx3 insitu
because from what i'v seen in test tube reactions it appears to polymerize but my reaction could very well be tainted.
I wish to gather more Iodine to perform a few more test but iodine isn't exactly something I can just get on a whim anymore and waste so I will have
to take much care in working with it.
PS: can anyone explain why some AL foil(all from the same source) reacted with HCl while most did not? Some is completely turned to an off white
powder while some is shinny as can be.
I think as well im going with Cl2 as my halogen and not screwing with HCl when I try to scale because I have shown better results in the past this HCl
work seems to have spotty results at best.
[Edited on 20-7-2010 by Sedit]
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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12AX7
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Concerning the sulfurous products, I suggest sulfur chlorides or related compounds (what does S2Cl2 do to aluminum metal?). Justification: at least
when the aluminum isn't reacting, the redox conditions in solution are fairly neutral, so maybe the HCl is hydrolyzing the vulcanized rubber and doing
funny things to it.
Tim
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Fleaker
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kmno4,
Thank you for doing that experiment!
Tim, S2Cl2 corrodes aluminum rapidly, at least if the atmosphere is damp.
Neither flask nor beaker.
"Kid, you don't even know just what you don't know. "
--The Dark Lord Sauron
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rrkss
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Microscale Prep of Anhydrous AlCl3
Thought this might be useful as it makes a difficult reagent easy to get.
Attachment: alcl3.pdf (1MB)
This file has been downloaded 1345 times
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Nicodem
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Threads Merged
26-7-2010 at 04:57 |
blogfast25
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So it involves (after fluxing with inert gas) refluxing a mixture of Al and I2 with DCM, at BP = 40 C. Overall reaction:
Al + 3/2 I2 +3 CH2CI2 --> AlCl3 + 3 CH2ICI
Quantities used: 1.85 mmol Al; 2.56 mmol I2, 31.2 mmol CH2Cl2
Chloroiodomethane has a higher BP: > 100 C
I think I might try that...
[Edited on 26-7-2010 by blogfast25]
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Chainhit222
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This method is pretty heavy on iodine consumption 
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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Lambda-Eyde
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Wouldn't a stream of chlorine gas displace the iodine in chloroiodomethane to yield elemental iodine and dichloromethane, and therefore be the only
reagent consumed in the reaction apart from the aluminium?
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Chainhit222
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Quote: Originally posted by Lambda-Eyde  | | Wouldn't a stream of chlorine gas displace the iodine in chloroiodomethane to yield elemental iodine and dichloromethane, and therefore be the only
reagent consumed in the reaction apart from the aluminium? |
So you think I can recycle the solvent by gassing it with chlorine? I could totally try this right now.... 
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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Lambda-Eyde
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I don't see why this reaction shouldn't occur:
Cl<sub>2</sub> + 2CH<sub>2</sub>ICl --> I<sub>2</sub> +
2CH<sub>2</sub>Cl<sub>2</sub>
And therefore iodine and dichloromethane are both acting as catalysts in this reaction. I'd love to try Peach's experiment myself, but I'm not doing
anything involving chlorine gas before my fumehood is up and running.
The only thing that bothers me is the low temperature of the reaction. In the classical preparation of aluminium chloride quite high temperatures are
needed for the reaction to proceed. This reaction is of course entirely different if the active chlorinating agent isn't elemental chlorine itself,
but dichloromethane (which is in fact known to react with certain metals), which could explain this.
It would be interesting to see if AlCl<sub>3</sub> could be prepared with HgCl<sub>2</sub> and a chlorinated solvent,
analogous to the preparation of aluminium isopropoxide.
This is of course just a wild theory, I have no idea what the equation for the reaction would look like.
Does anyone with some more insight in organic chemistry have any comments?
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Nicodem
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| Quote: Originally posted by Lambda-Eyde | I don't see why this reaction shouldn't occur:
Cl<sub>2</sub> + 2CH<sub>2</sub>ICl --> I<sub>2</sub> +
2CH<sub>2</sub>Cl<sub>2</sub>
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You don't see why it shouldn't occur and I don't see how it could occur. I doubt it is only a difference in opinion as your equation calls for a
electrophilic substitution on an alkyl iodide which is very unusual to say the least. Alkyl iodides are known to get I-chlorinated by Cl2 under proper
conditions to give R-ICl2 compounds, but no such electrophilic substitution as above was ever reported to my knowledge.
| Quote: |
It would be interesting to see if AlCl3 could be prepared with HgCl2 and a chlorinated solvent, analogous to the preparation of aluminium
isopropoxide.
This is of course just a wild theory, I have no idea what the equation for the reaction would look like. |
In one of the previous AlCl3 threads, I proposed that this kind of an reaction might be a viable source of AlCl3. But in view of kmno4's interesting experiment described above, I'm not
any more convinced this would be practical or possible at all. In regard to this topic and also this thread's topic, see also the post by Greenimp in
that same thread just slightly bellow mine.
…there is a human touch of the cultist “believer” in every theorist that he must struggle against as being
unworthy of the scientist. Some of the greatest men of science have publicly repudiated a theory which earlier they hotly defended. In this lies their
scientific temper, not in the scientific defense of the theory. - Weston La Barre (Ghost Dance, 1972)
Read the The ScienceMadness Guidelines!
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Lambda-Eyde
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| Quote: Originally posted by Nicodem |
You don't see why it shouldn't occur and I don't see how it could occur. I doubt it is only a difference in opinion as your equation calls for a
electrophilic substitution on an alkyl iodide which is very unusual to say the least. Alkyl iodides are known to get I-chlorinated by Cl2 under proper
conditions to give R-ICl2 compounds, but no such electrophilic substitution as above was ever reported to my knowledge.
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Thanks. That makes sense. I was just thinking along the lines of "F > Cl > Br > I", but when carbon comes into play it's a whole different
game.
I'm reading quite a bit of organic chemistry, but I don't have the "feel" for it as you and quite a few other members have. I don't look at that
equation and automatically think "Aha! Electrophilic substition on an alkyl halide!".
I can only hope that that sixth sense comes with time (and reading). 
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Formatik
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Has anyone tried reacting liquid HCl with aluminium to get AlCl3? It sounds easy enough, and powder might not even be needed. The reaction has been
described in: Proceedings of the Royal Society of London, Vol. 14, p. 209: "Metallic aluminium became dull in the gas, and quickly dissolved,
with evolution of gas, when the liquid acid came into contact with it and formed a colourless solution".
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rrkss
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What you are referring to as liquid HCl is HCl gas dissolved in water. The preparation of AlCl3 needs to be water free to work so that procedure will
not work.
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blogfast25
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| Quote: Originally posted by rrkss | | What you are referring to as liquid HCl is HCl gas dissolved in water. The preparation of AlCl3 needs to be water free to work so that procedure will
not work. |
No, I think he's referring to water free liquid hydrogen chloride (atmospheric BP = - 85 C). This could still be kept liquid at higher pressures. But
I can't see how it would react with Al without a powerful Lewis acid present as catalyst.
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Formatik
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Yes, liquid, not aqueous HCl. Strong H2SO4 and NH4Cl was used to generate the gas. The HCl liquefied under high pressures (500 to 1100 psi). For
further details, that ref. is here.
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Methansaeuretier
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Anyone ever tried to use chlorine in DCM instead of HCl in DCM? Chlorine is very easy to prepare and usually it is also very reactive:
Reactivity --> F > Cl2 > Br2 > I2
Also Cl2 does not react violent with air when heated.
[Edited on 11-8-2010 by Methansaeuretier]
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blogfast25
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| Quote: Originally posted by Methansaeuretier | Anyone ever tried to use chlorine in DCM instead of HCl in DCM? Chlorine is very easy to prepare and usually it is also very reactive:
Reactivity --> F > Cl2 > Br2 > I2
Also Cl2 does not react violent with air when heated.
[Edited on 11-8-2010 by Methansaeuretier] |
Look higher up in this thread.
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peach
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| Quote: Originally posted by ziqquratu |
By the way, it occurs to me, isn't AlCl3 soluble in DCM? If I'm right (which at this time of the morning...), this suggests that your precipitate is
NOT AlCl3. |
Not really. I've left commercial acid sat in it for three weeks, shaking it every time I walked past, to clean some brown muck off it's surface. I
think I mentioned this at the start, avec pictures.
This is a good example of why I'd like to make it myself. It isn't too hard to get, but it does require a supplier account a lot of the time, which is
becoming near impossible for the public to obtain (without telling some big porkies). The reason I'm mainly thinking of (for making it myself), is the
amount of discolouration in that CP sample. That's no where near pure. The colour likely won't produce much of difference in terms of the results
it'll produce, as it's probably some other Lewis acid on the surface. But, as you can see, it produces extremely dark staining of the solvent, which
pure AlCl3 won't do. That makes it very tricky to determine what's happening in any subsequent work that involves watching for colour changes.
Particularly if tar is a potential result of those organics, with the dark brown significantly increasing the chances of a false judgment being made.
If the work is actually experimental (no real references), it gets worse again.
Here's a sample of commercial (CP grade) AlCl3, fresh out of the bottle and into a flask. Note the disgusting discolouration;
And after 3 weeks of soaking and swirling (in fact, I used around half a liter of DCM attempting to clean this up);
Solvent decanted and AlCl3 rinsed multiple times with fresh DCM;
Stripped of DCM under vacuum;
| Quote: Originally posted by anotheronebitesthedust | | Maybe the neoprene tubing is adding unforeseen variables into the reaction. DCM penetrates neoprene and it was mentioned earlier in the thread that
neoprene often contains sulfur. Is aluminum sulphide possible? Were all your reagents free of contaminants? |
There were only three components in the original experiment. The aluminium I've tried digesting with KOH and then the hydroxide with sulphuric.
Behaves perfectly. Pure white hydroxide that clears in acid. The DCM is CP grade.
When I dipped the neoprene into the solvent, that's when the problems began. It will be the neoprene. And
specifically, it being in contact with the solution. I don't have anymore commercial AlCl3 to try this test with though.
As I pointed out to blogfast, neoprene out of the solution is fine for porting HCl(g) around, it doesn't do anything. It's also okay
around DCM and aluminium. So, the only option is, there's AlCl3 in the solution or the DCM + HCl(g) is creating the compatibility issue.
Before I had the neoprene touching the actual solution it's self, it was going fine.
Indeed it is.
Below is the experiment that started this thread. In the flask is clean, atomized aluminium, the DCM is CP grade, the down tube from the wash head is
blowing in dried HCl(g). That's all, nothing else is there but for the borosilicate glass. The plate's element is switched off.
As I've said previously in this thread, bare in mind that the solvent something is in changes it's reactivity, sometimes drastically. Using the
example I gave earlier, conc. H2SO4 becomes more reactive as it's diluted down with water, to allow it to disassociate. Similarly, concentrated nitric
will passivated metals until a little water is added. The only thing changing in these examples is the solvent for the acid.
VIDEO OF ALCL3 COLD FORMING IN DCM
>>>>>>>CLICK ME<<<<<<
VIDEO OF ALCL3 COLD FORMING IN DCM
| Quote: Originally posted by Sedit |
I tryed just recently to place a long glass tube on a flask filled with H2SO4 an NaCl. The tube was filled with some shredded Al foil and the HCl was
allowed to escape thru the top.
Oddly enough what I got was patches of Al turned white. . There would be white
patch in the middle but the bottom was un reacted ect ect... it really makes little sense because it should be Al mostl reacted at the lowerend and
progressivly get less reacted as the HCl is used up.
It was just a random experiment and I intend to eliminate variables shortly. |
As much as I'd like to assure you that's AlCl3, it may be due to moisture in the HCl(g) stream producing hydrochloric on the metal, with it being so
close to the gas generator it's self. However, as you can probably guess, I'm willing to consider the possibility this reaction does occur at room
temperature, just slowly.
I would also recommend ziqquratu's suggestion of igniting the exit stream to look for hydrogen. I'll give this a go next time as well.
My chemistry knowledge runs out as to where it'd be picking the hydrogen up from. But, if the solvent going green is free chlorine, it'd suggest the
HCl(g).
| Quote: Originally posted by Sedit | This page offers the suggestion that DCM is viable solvent for AlCl3, I honestly beg to differ when free halogen is used to produce Alx3 insitu
because from what i'v seen in test tube reactions it appears to polymerize but my reaction could very well be tainted.
I wish to gather more Iodine to perform a few more test but iodine isn't exactly something I can just get on a whim anymore and waste so I will have
to take much care in working with it.
PS: can anyone explain why some AL foil(all from the same source) reacted with HCl while most did not? Some is completely turned to an off white
powder while some is shinny as can be.
I think as well im going with Cl2 as my halogen and not screwing with HCl when I try to scale because I have shown better results in the past this HCl
work seems to have spotty results at best.
[Edited on 20-7-2010 by Sedit] |
With regards to solubility, see the pictures above for a graphic demonstration of that. And there was a lot more DCM used there than
you see in the photos.
But that's not to say it doesn't dissolve to some extent. If I pour that DCM off, it'll fume as it goes into the sink.
Thinking about your aluminium foil, was this a brand spanking new piece harvested from underneath the first wrap or two, and were you wearing gloves?
It's possible moisture from your fingerprints has helped the gas form hydrochloric acid at those sites. If the foil is scrunched up at all, it'll also
encourage 'concentrating points'. Anything that's not a sphere has points where energy of all forms all precipitate, heat, pressure, stress, strain
etc. That's why bathyscaphes, firework shells and the cores of nuclear weapons are spherical. It's also why pryotechnicians will pay a lot more for
'german / indian blackhead flake' aluminium for their shells, because the flattened out profile, with sharper edges, ignites noticeably easier and so
it burns very rapidly. Seems like a stupidly minor factor, but it does influence the real world, sometimes dramatically; e.g. the score line down a
piece of glass.
| Quote: Originally posted by Lambda-Eyde | I'd love to try Peach's experiment myself, but I'm not doing anything involving chlorine gas before my fumehood is up and running.
The only thing that bothers me is the low temperature of the reaction. In the classical preparation of aluminium chloride quite high temperatures are
needed for the reaction to proceed. This reaction is of course entirely different if the active chlorinating agent isn't elemental chlorine itself,
but dichloromethane (which is in fact known to react with certain metals), which could explain this.
|
A fume hood is a good idea, but an even better idea is to scrub the chlorine out of the gas leaving the glassware using something it'll dissolve or
absorb well in. I'm running this kind of thing in the kitchen, stood right beside it with absolutely no protective gear on. I can't smell, see or in
anyway sense the gas because I'm scrubbing the excess out with a base wash on the exhaust.
Tightly fitting tapers, greased, are a very good idea however. As is keck clipping them to make sure they're all seated. Whatever you're doing, it's
usually also a good idea to leave one or two strategically chosen joints unclipped, in case there's any unexpected, odd excursion in the pressure;
e.g. boiling goes mental, reaction runs away, exhaust gets clogged etc...
I don't think it's the DCM alone that's doing this. I think, at most, the HCl(g) could be helping it. I think the far more likely option is that the
DCM has changed the availability and reactivity of the HCl(g) with regards to the metal.
| Quote: Originally posted by Nicodem | | In regard to this topic and also this thread's topic, see also the post by Greenimp in that same thread just slightly bellow mine.
|
I used to talk with greenimp all the time on another forum.
He was very dedicated to trying things, but he (like myself) also wasn't super knowledgeable in terms of the theory compared to some of you, and he
gave up with it all after a while.
But there is a critically important factor he mentions in his post, which demonstrates that he has actually tried what he's talking about. In that he
mentions the reaction carried on for hours on it's own, even after the generator was out.
That's identical to my own experiences using DCM. I think this is one where the flow rate either has to be absolutely minute, or it needs gassing and
then leaving to sit, then regassing once it's stopped. Otherwise, it'll be easy (I expect) to pour a ton of gas through, have the solvent saturate and
the rest to go out the exhaust.
Yes, before you ask, I was producing the gas very slowly. As slowly as I could get the drips out of the funnel after spending ten minutes tweaking it
like a kinky nipple addict. It was still, I expect, too quick. Another possible benefit to actually having a regulator controlled supply of the gas.
I've been trying to squeeze some details out of BOC about what they can offer in terms of training and what paperwork I need to fill out to rent the
corrosive cylinders; after they sent me a message about safety training courses for regular gases. So far, they've told me they can do training with
the lab gases, but I need to book a visit from the staff and it's aimed at labs where 8 or so people will be listening in. As it's obviously something
where they'll have to send a very specific person out given their odd properties. I don't know what the actual paperwork requirements are, so I'll
have an ask about that.
That aside, I expect it will be eye wateringly expensive to rent the bastards. Coupled with the specialized regulators, rotting and then the
subsequent, quite probable, police lab visit to see what on earth I'm renting it for; first hand.
"Yeah... can I have the biggest HCl(g) and methylamine cylinders you do, and can I pay in all these used £5's as well? I don't have an account, but
my goldfish asked me to get them." "No..."
| Quote: Originally posted by Formatik | | Has anyone tried reacting liquid HCl with aluminium to get AlCl3? It sounds easy enough, and powder might not even be needed. The reaction has been
described in: Proceedings of the Royal Society of London, Vol. 14, p. 209: "Metallic aluminium became dull in the gas, and quickly dissolved,
with evolution of gas, when the liquid acid came into contact with it and formed a colourless solution". |
I haven't but, once I have a dry ice / LN2 capable condenser, I may give it a whirl.
Note that it'll react with the liquid HCl, but we're going on and on about the gas. As I have now said a few times, I suspect the DCM is allowing a
pseudo liquid form of the HCl to collect around the metal by vastly concentrating it over the normal gaseous form.
It is. However, the few notes I've seen about how it's done industrially suggest Cl2 actually needs the aluminium to be a lot hotter than it need be
with HCl(g).
I've now spent £ha£hing on a few wash bottles specifically with the intent of redoing this all in glass to a full conversion to the white
precipitate, but I'm still trying to collect a few other sizes so I don't have to run it in this monster. For a size comparison, there's a tiny
tomatoe, a slightly bigger tomatoe, a large tomatoe, a jar of beetroot, a 250ml wash bottle, the 500ml wash bottle, a 2l bottle of Vimto concentrate.
Note that the down tube still doesn't get very close to the base, so I'll need some more borosilicate on the end and a frit. I'm starting to feel like
Dr Evil getting this sorted...
"You know, I have one simple request. And that is to have sharks with frickin' laser beams attached to their heads! Now evidently my cycloptic
colleague informs me that that cannot be done. Ah, would you remind me what I pay you people for, honestly? Throw me a bone here!"
I have some Tygon 2375 in the post, which is Saint's updated version of their 2075 Ultra-chemically resistant tubing (doesn't even seem to be listed
on their site yet). Unless someone wants to paypal me funding for PTFE (which also requires special fittings) and other more high end bits and pieces,
that's the best you're getting.
Blogfast is correct in one of his suspicions, that I'm actually a guy. As much as your sexually excited PM's make me smile, their true intentions are
perhaps wasted on me. 
As usual, the difference in attitudes is notable when people, looking at my forum name, assume I'm a girl. There is a certain air of... hmmm...
?inexplicable friendliness? to them, which is remarkably lacking from those in response to posts I've signed as "John". I have even mentioned being
'bollock naked' in response to one PM, and received a reply even surer of me being a girl. I chose peach because I like the fruit and the colour, and
because orange sunshine probably isn't a great idea on chemistry fora.
John
[Edited on 13-8-2010 by peach]
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blogfast25
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The one thing that remains a weak point in your experiment, erm... 'John', is the that you don't seem to have characterised your main reaction
product. AlC3 isn't that hard to identify: dry you can sublime it easily. It should also dissolve easily and plentifully in strong HCl. And with water
it should hydrolyse quickly with considerable heat generation...
And then there's the nature of the by-products: the greenish stuff and the dark stuff.
Well worth repeating with an all-glass apparatus, IMHO...
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