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Author: Subject: Preparation of anhydrous AlCl3 in DCM - photos
peach
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[*] posted on 19-7-2010 at 11:21


No, it behaves as a solid crystalline.



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anotheronebitesthedust
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[*] posted on 19-7-2010 at 12:50


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:
Quote:

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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[*] posted on 19-7-2010 at 13:56


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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[*] posted on 19-7-2010 at 14:05


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. :o. 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.





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[*] posted on 19-7-2010 at 15:40


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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[*] posted on 19-7-2010 at 16:09


Quote: Originally posted by peach  
No, it behaves as a solid crystalline.


OK but that starts me thinking about aluminum hydride again.




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[*] posted on 19-7-2010 at 16:14


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]





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[*] posted on 19-7-2010 at 19:07


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.

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[*] posted on 22-7-2010 at 05:58


kmno4,

Thank you for doing that experiment!


Tim, S2Cl2 corrodes aluminum rapidly, at least if the atmosphere is damp.




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[*] posted on 26-7-2010 at 03:26
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 1331 times
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[*] posted on 26-7-2010 at 06:27


Quote: Originally posted by rrkss  
Thought this might be useful as it makes a difficult reagent easy to get.



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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[*] posted on 28-7-2010 at 05:41


This method is pretty heavy on iodine consumption :(



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[*] posted on 28-7-2010 at 08:56


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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[*] posted on 28-7-2010 at 10:58


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.... :cool:




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[*] posted on 28-7-2010 at 12:50


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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[*] posted on 28-7-2010 at 15:20


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>

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.




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[*] posted on 28-7-2010 at 16:01


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.


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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[*] posted on 29-7-2010 at 23:00


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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[*] posted on 30-7-2010 at 03:43


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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[*] posted on 30-7-2010 at 05:41


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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[*] posted on 30-7-2010 at 19:53


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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[*] posted on 11-8-2010 at 11:07


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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[*] posted on 11-8-2010 at 13:26


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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[*] posted on 11-8-2010 at 21:48


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.

Quote: Originally posted by kmno4  

As they say - Experiment is a king :)


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. :o. 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.

Quote: Originally posted by IrC  

OK but that starts me thinking about aluminum hydride again.


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.

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:


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. :P

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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[*] posted on 12-8-2010 at 03:53


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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