Sciencemadness Discussion Board - Preparation of anhydrous AlCl3 in DCM - photos

I'm speculating, but if this inspires anybody else for alternative mechanisms, I'm happy. It does seems that getting Al out of its crystal matrix is hard with just isolated single chlorine atoms in HCl. I get the impression that a bit of Cl2 in solution might help initiate the reaction, since adsorbed Cl2 would give three chlorine atoms adsorbed in proximity far more frequently. If that works, it would certainly be more benign as an initiator than HgCl2.

The experiment I performed sometime ago(sorry most data was lost in a computer crash) was based on two test one as a control and one with HgCl2 as an activator. The DCM was saturated with Cl2 till it was yellow in both test tubes. Al foil was shreaded in a blender along time prior to this experiment. An equal amount of Al was added to both with a trace amount of HgCl2 in one of the test tubes. About 1 hour later they where observed and there was almost no reaction in the control which was still yellow with Al sitting on the bottom and some has rose to the top .

The one with HgCl2 was no longer yellow and much of the Al had turned to a grey(ish) powder/gitter. Im sure an excess of Al was used and im pretty confident that if Cl2 was feed in I would have sustained a reaction and managed to consume all the Al present.

It was all done in the absence of light at RT BTW which could have caused a different outcome due to Cl radicals attacking the Al easier.

As a result I am not very confident that Cl2 would be efficient in starting the reaction safely. Perhaps best to just try to heat some Al in a testtube while feeding in Cl2 and scrape the subliminate off the edge to get the needed AlCl3 to catalyse the reaction. Or feed only a small amount of HCl at first into your reaction and allow it to age before proceeding.

peach - 5-7-2010 at 14:31

[every post needs a photo]

Next time add abit of activated Al using HgCl2 to avoid the runaway and allow the formation to start right away. I did this a while back on a small scale using Cl2 and small amount of activated Al and it appeared to proceed nicely so If I where to ever scale up, something which I plan on doing one day I would indeed add a small amount of Al/Hg just to get the ball rolling and that would provide you with the AlCl3 needed to initiate the rest of the reacton.

That's an interesting idea. Where you simply gassing it at room temperature in solvent? DCM? It's surprising Cl2 worked, since that seems to need hundreds upon hundreds of C going directly against Al, whereas HCl works quite a lot lower.

Love the pics. It that rainbow there as a color chart or just an artifact of the photo? Its a good idea either way and I think I may start adding ones to photos to give a better idea of the reaction mixtures color under the light I was under.

If you like that, you'd love my underpanties. They won me an award, $20, for a photo of them, which I believe has now been publicly circulated.

I didn't add the book on purpose, but that is an extremely good point you make. As that was actually an issue with this, the camera didn't pick the green of the DCM up at all well! Before the reaction started, it was strong limey green.

The book is "Designing the 21st Century"

I'm guessing at least some AlCl3 may have formed then, so it may pick up again if regassed. I'll give it a try once I have some more sulfuric. If that doesn't work, I'll prime it with some Hg salt as suggested.

When the reaction really started to pick up, it was very impressive. It had been doing absolutely nothing for hours as I stared at it. Then I looked back and there was a tiny, tiny stream of bubbles. Some time later, it was going like a shook up bottle of coke. I did worry it was going to run away, but it didn't significantly warm up or accelerate past there.

Using DCM as the solvent probably made it impossible for the temperature to go very high, but it didn't even reach DCMs boiling point. And everyone knows DCM can be refluxed by a warm fart.

Quote: Originally posted by DJF90

If you look carefully in the picture you can see its the cover of a book. Although the aim is to obtain AlCl3, its nice to see attempts that are simple like this, but which do not use mercury or its salts; as chemists we have the responsibility to experiment safely and in an environmentally sound manner, and safe disposal of mercury wastes (especially in solution) is not easy.

[Edited on 5-7-2010 by DJF90]

Well spotted, can I assume you've seen Blade Runner? "enhance, pan... print copy". Now if you look even more closely...
closer...
you'll see there's someone naked standing in the room.

I have some mercury chloride, but you're right, I'd prefer not to open it if I don't have to really. Whenever I'm doing something involving gases or vapors, I try to wash the exhaust of the reaction through something I know will entirely neutralize it to a salt, for instance. But it's better to avoid a risk as opposed to patch it up.

My guess is that the dissociation of HCl is far greater with AlCl3 present than not, because of the equilibrium AlCl3 + HCl <--> AlCl4(-) + H(+). Electron transfer in the net reaction Al(n) + H(+) --> Al(n-1) + H(0) would then proceed at a much higher rate.

I wonder if a small amount of iron salt would be enough to start the disassociation and allow the Al to pick it up.

I have much less an idea what the intermediates are, but it seems that the important facts are that the Al is a solid and a good electronic conductor, so it's useful to look at half-reactions at the surface that involve free electrons

Exactly, it seems like a pretty basic thing to make, but there was certainly something being produced hours after the gassing had stopped. And it wasn't being produced earlier.

The only thing I could think of is that some of it had been converted to AlCl3, because I can't imagine what else it'd so readily form into along the way, involving the evolution of what I suspect was H2.

Apparently, AlCl3 it's self is a rubbish conductor. For reasons I think someone has already hinted at (covalent like ionic bonding).

It does seems that getting Al out of its crystal matrix is hard with just isolated single chlorine atoms in HCl. I get the impression that a bit of Cl2 in solution might help initiate the reaction, since adsorbed Cl2 would give three chlorine atoms adsorbed in proximity far more frequently. If that works, it would certainly be more benign as an initiator than HgCl2.

That's an idea as well. I'd prefer not to start using the hot processes if possible. Although, I could give them a go too. I'll have to visit the local neon sign store for some borosilicate tubing first.

I'm also idly curious as to whether there's any significant or measurable chlorination of the DCM to chloroform.

Urban Dictionary;
"Does This Smell Like Chloroform?

Jack: "Hey Dorothy, does this smell like chloroform to you?"
Dorothy: "Hmm let's see..." *smells, faints*
Jack: "Great Success!" *unzips pants*"

There is indeed an odd smell to the DCM. When I first smelt it, I thought it smelt eggy. Smelling it some more, it is somewhat ethereal and sweet smelling. It's certainly not HCl(g) or DCM it's self.

I was using neoprene tubing to direct the gas around, but that shouldn't have produced that kind of smell. I don't really like the neoprene, I'd like some Tygon or something like that.

I'll investigate when recovering the solvent. Hopefully there's enough there to show up on the thermometer, and I get a solid number out of it. It'd be fun if it was hey! UV illuminating the DCM could accelerate the process if so; perhaps explosively, so probably best tested small scale and out of the glass first. Actually, I wonder if UV would help the disassociation here. [adds deuterium lamp to shopping list]

Quote: Originally posted by smuv

A more efficient way to do this reaction, might be to do it in a 2 neck flask, or even a filter flask. Put a balloon over one neck and affix a hose leading to the HCl generator to the other. Fill the balloon with HCl, clamp the hose, and allow everything to stir under an atmosphere of HCl or Cl2. The solubility of HCl in DCM is fair, however I fear most of the HCl you generated did not dissolve.

P.S. Just a nit picky correction: Anhydrous AlCl3 should not be called a salt, the bonds are mostly covalent in character.

[Edited on 7-5-2010 by smuv]

I can't use the balloon trick. My gas generator is glassware, so it'll pop if I restrict the outlet. I try to gas things as slowly as I can get the gas generated. I'd much, much prefer a fritted wash head, but the frit fell off mine after some harsh cleaning. That another thing on the shopping list, which is getting longer by the minute...

I think a decent amount of the HCl(g) went into solution. As I say, the camera doesn't pick up how green the DCM was very well.

But I'm considering building a HCl generator for the ease of preparation, measurement and flow control; the generator isn't particularly convenient or study even with a pressure equalizing funnel. I googled it and found a document about "Clandestine HCl generators". A ton of people use old propane bottles, which seems extremely risky; there's no way I'd do that myself. The authorities show the correct way to do it in their own PDF, plastic pipe fittings with a pressure relief. The pressure never seems to go over 60psi (although I note, they didn't try charging it with more reactants IIRC).

Personally, I'd have the blow off valve connected to another bit of tubing full of strong base in solution, to deal with all that acid should it pop. A safety feature I also note the authorities not using.

And yes, you're correct about the bonding.

DCM = dichloro methane?

Is AlCl3 supposed to be soluble in DCM?

CaCl2 isn't the greatest drying agent, if pure anhydric, unhydrolysed AlCl3 is your goal (for organic catalysis?)

DCM = Dichloromethane. My last container of 'paint stripper' (pure DCM) arrived in this box. Dun den deeeeeeerrrr..... But is it toxic watson?

I'm only pulling your leg.

AlCl3 is supposed to be soluble in DCM yep, you'd expect it to be really soluble on first appearances. But it's not all that soluble at all. This makes me wonder why my DCM may have clouded. I doubt it's the salt, but maybe it's some of watson's chloroform?

I chose DCM because it's polar aprotic and it'd readily absorb the HCl(g) and could be easily stripped away from the AlCl3, it can even be used as the solvent for work further down the line.

Enough AlCl3 will dissolve in DCM that the solvent will fume when exposed to water, however.

CaCl2 isn't the greatest drying agent, if pure anhydric, unhydrolysed AlCl3 is your goal

I should probably use some concentrated sulfuric I suppose. AlCl3 doesn't absorb airborne moisture too rapidly. It'll certainly do it, but not at an unhandleable rate. My main concern about getting moisture on it is something I think Sauron said about using it to demethylate xylene and there needed to be a Bron.Low acid present for the F.C. I'm considering using xylene and such things as a solvent with the Lewis acid, so I don't want it stripping methyls off wherever it likes.

Quote:

I wouldn't have expected the reaction in DCM to proceed at RT without catalysis: any dissolved HCl is likely to be undissociated or very poorly dissociated.

I wasn't expecting it to either, given the heating that's usually needed. I was sure it was a complete flop, until it sprang it's self into life and started evolving it's own gas. It's also curious, as I say next to one of those pictures, that the bulk of the Al isn't reacting at first, yet something in the wash head is fizzing away. Maybe that's moisture? Hmmm. A conundrummer.

I'm guessing that anyone who can set up a HCl generator and lead the HCl through a suspension of Al in DCM without poisoning himself is also capable of leading said HCl over hot Al filings or powder, collecting the subliming AlCl3 and safely dispose of the H2.

Thanks for the vote of confidence. You're right, in terms of a proven method, that's pretty much unbeatable and yields a sublime result!

I'm a huge advocate of wash heads and actually neutralizing things before they escape over relying on a fume hoods to catch and absorb things that get airborne. So for this example, I was blowing the exhaust through strong KOH solution, there was absolutely no sign of HCl at the output.

And it's not poisoning! It's burning! That's something of a special joke which watson alone will get.

I'm avoiding using the hot methods because I need to sort out my glassware to sublime in a dry atmosphere. Also, a cold method would be easier just in terms of handling the glass and so on.

I want to make my own firstly for the fun of doing it, but also because this is the state the last batch I ordered turned up in. This is straight out of their sealed bottle. I suspect that's iron contamination, which is common with commercial AlCl3 apparently. I don't know how iron gets in there, maybe it's sludge graded Al they use from scraping out ladles or something. But iron trichloride is not brown, and this stuff is only sitting on the surface of the grains.

I'd had it sitting in a puddle of DCM for weeks, changing it whenever it went brown again. I must have had it in DCM for three weeks or more, and no visible quantity of it dissolved. Look at the state of the solvent! Incidentally, in it's time in the Hotel de Flask, it chemically etched the taper shut. I had it in there for weeks because I was trying everything I could to get the stopper back off. I had to smash it in the end. Yes, some idiot forgot to grease / tape the taper before it jammed. I'd like some PTFE keck clips, bottle caps and tapers.

Because it seizes glass, and does this to rubber... Silicon or some other plastic would probably work, but virgin PTFE bar is cheap, turning a taper is easy and it's great stuff.

This is after weeks of being repeatedly shaken. Letting it sit and soak is the best way to clean it up if you have some like this. The yellow color is either water or iron trichloride. I would vote for iron contamination, as all the grains would be greeny yellow if it was moisture.

Something moderately more clean. If it's iron trichloride, it shouldn't be too much of a problem for what I'm using it for. But it makes determining colors difficult, because it's horrendously dark brown whatever it is. The bottle claims to be 99% Al, CP grade. It's still extremely reactive, it's not over exposed; it violently fumed when exposed to water at a later date. I don't understand why the brown only appears to be on on the surface though. Also note that there are white grains in there, which'd be yellow if they were damp. In fact, it looks like multicolor popping candy. Consistency does not appear to be a factor for the manufacturer. Although, some of those grains still look brown. A mixed batch? There's three different colors in there.

(for organic catalysis?)

Indeed sir, indubitably.

I've been investigating the use of various Lewis acid for Hard Acid demethylations.

AlCl3 has been a weapon of choice, but I haven't been getting very promising results in the past.

BF3... jeez... too expensive... too nasty... too aggressive, even AlCl3 seems too aggressive.

I may try FeCl3, I was drying some today under HCl. The hexahydrate melts at tens of degrees and boils at around 280. The anhydrous salt melts over 300 and boils tens of degrees later. I managed to obtain something that stayed dry and solid at 280C, but I didn't see any of it melt as it went on to 320C. I did see some vapors coming from it. I may have managed to reduce it to iron (II) chloride, which melts over 600. Or iron oxychloride. There was an ominous leafy green tint to the flask after it cooled (indicating the (II) chloride), but it's now disappeard. Haven't tested it with some water yet, I only made about 10g so there's not a whole lot to test. I have pictures, but this is getting off topic...

I've also used AlI3. I doubt something could make more mess if it tried! And it consumes expensive I2 at quite a rate. At least it's super simple to prepare.

I recently had some luck with AlI3 though, swapping round solvents, and it made me wonder if the same would apply for AlCl3 or other Lewis acids, like FeCl3. I may try running these with the addition of some dry ice if tar production continues with different solvents. Ice / salt baths don't seem to work for AlCl3.

I expect someone is about to ask, so I'll say now that the substrate I'm talking about is primarily 4-Allyl-2-methoxyphenol. I'm sure you're all bored out of your minds seeing posts about ideas on that. I'm specifically not using the common name so the thread doesn't go down 'that route'. Also, I'm genuinely not interested in this for any profitable reasons, it's simply something of a curiosity to me. Again, that's a different topic though, this is about Lewis acids.

[Edited on 6-7-2010 by peach]

Sedit - 5-7-2010 at 16:51

AlI3

That could be your safe alternative to using HgCl2 I would suspect. Add Iodine until its consumed and that should give you the catalytical amount of Alx3 compound needed to start the reaction when the solution is gassed.

I would assume more then likely incorrectly so that trace iodine could also speed the reaction up somewhat and allow it to proceed easier.

I thank you peach you are inspiring me by filling my head with so many "what ifs" that I may start experimenting with this again since it is a very useful reagent for a number of things yet as I understand its a bitch to store and somewhat of a bitch to prepair so if a quick an simple cold lab prep could be worked out it could indeed open alot of doors. I remember attempting my HgCl2 experiment in the past because someone told me that you could not generate AlCl3 by gassing withou AlCl3 being present to act as a catalyst so I had to prove that a catalyst of somesort could be generated in situ and I do believe my testtube results confirmed some reaction took place.

Keep going man before you know it I maybe working riht along side of you to tackle these various means but first I have to finish my new lab bench which I making a fume hood for to YEHHHHHH I feel so excited for the first time in years.

blogfast25 - 6-7-2010 at 06:05

You should be able to leach out any FeCl3 with dry acetone in which it is readily soluble (but I'm not sure whether AlCl3 is acetone soluble or not).

peach - 6-7-2010 at 10:56

You should be able to leach out any FeCl3 with dry acetone in which it is readily soluble (but I'm not sure whether AlCl3 is acetone soluble or not).

I genuinely have no real idea what's going on with the brown coating. It seems to permeate some of the grains in that last photo, but I seriously shook the shit out of that 25g worth over three weeks. I'd walk past it and shake away. The DCM would turn brown almost immediately. It took weeks worth of shaking to get rid of most of it, and some still remained in the valleys of the grains.

At first I tried washing it with an excess of DCM, but it just wouldn't move. It kept staining but not moving. When the taper sealed, I couldn't replace the solvent anymore and I had to let it sit whilst I awaited an order of something other things (which took a month to arrive). I ended up using half a liter of CP grade DCM to try and clean that, and repeated, quick rinses do not work.

I tried SO many things to open the taper. Heating it, cooling it, dropping solvent onto it and covering it with film. Cooling it to draw the solvent in. Tapping, force. KOH to dissolve the glass in the taper. Absolutely nothing budged it. It was 100% jammed. I eventually had a 25mm wrench on it, with the flask wrapped in a towel to avoid cuts. It just didn't move AT ALL. So I had to put a screw driver through it in the end.

I tried to save the flask but, even with the stopper smashed to pieces, the remaining glass wasn't going anywhere. I seriously doubt that's a mechanical bond, I think the acid has actually etched the taper shut somehow.

I should have wrapped the glass taoer in PTFE, I do so for most work. Instead, I just stoppered it and went to bed. The next day, it was jammed forever. PTFE, win!

I do think that commercial sample is contaminated with iron. As I say, if it was water, all the grains would have gone greeny yellow.

What I don't understand is why the surface is so brown. The only thing I can think of is oxygen, but chlorine should displace that. Maybe it was removed from the forming process whilst still hot. But then, why would oxygen displace the chloride?

I began to suspect I'd been sold some form of counterfeit product, as it's from a big branded supplier. AlCl3 is white, not filthy brown.

I would seriously like an answer as to what that is, I can't stand not knowing.

If you get some AlCl3 like that, do this...

Put it in a flask and just about cover it with DCM. Stopper the flask with something wrapped in PTFE or perhaps use a silicon / HDPE / polyprop stopper. I don't think polyacetate will work, as HCl(g) ripped my polyacetate clips to bits in no time. PTFE plumbing tape works well if you don't have PTFE stoppers / cap liners. The gas rated kind (in yellow packs) is thicker and doesn't tear so easily.

Don't add excess DCM. Let it sit in the DCM until the solvent is filthy brown / black. Then decant the solvent (it can be poured off without loosing a grain of the lewis acid) replace and repeat. Once the solvent isn't staining filthy brown anymore, boil the remains off under vacuum.

AlI3

Yeah, you know it. It's brutally dirty stuff. I'd buy a fume hood if only to keep the damn staining to a minimum, I need some of that lab bench paper. I2 smells really nice, and water with it in tastes really nice, but damn... get it ANYWHERE and you'll need to buy some sand paper to shift the stains. I was working as carefully as I could, to 0.01g and using grade A's, and it still ruined things around it.

But it performs really nicely for this kind of work. If I use AlCl3, I usually end up with a lot of what looks like tar and I don't see any Al elemental returned. If I use AlI3, I get a nice sandy / beige colored result and a mountain of Al returned. I have tried AlCl3 at salt / ice temperatures, but it seems to yield tar. Dry ice is an option, as that seems to help with BF3, which I don't want to get involved with.

I have a ton of photos from a recent work up, I'll send you a copy. PM me your email; since it's off topic.

I can't think of anything that has formed during this DCM based method that isn't AlCl3, in terms of the bubbling and reaction I saw occur. Can you see the videos? That took hours to start, but you can see it's running. It may be that the initial AlCl3 simply takes a while to start forming. I need to try regassing that sample really and see if it picks up where it left off. I think a decent amount of my HCl(g) made it into the DCM, but at least some of it will have escaped. Especially as I'm not using a frit.

I think I reached the solubility limit of DCM fairly soon after I started gassing, as it was lime green from early on and didn't get a whole lot darker as the generator ran over night. Indicating it has reached the limit and was going to waste.

I don't like using iodine. It's just far, far, far too messy. It's the only thing that stained my school's chemically resistant lab benches. It's even stained my PTFE bar. Messy stuff. Shame it smells and tastes so nice, and is such an amazing antiseptic (You know iodine is effectively the cyanide of bacteria? Great stuff!).

I would assume more then likely incorrectly so that trace iodine could also speed the reaction up somewhat and allow it to proceed easier.

Depends which reaction you're talking about. If you mean the 'salt' formation, then probably. If you mean the demethylation, I'm less sure. And I can promise you, that latter reaction DOES NOT need speeding up when AlCl3 is concerned; including via a soft base pusher.

I thank you peach you are inspiring me by filling my head with so many "what ifs" that I may start experimenting with this again since it is a very useful reagent for a number of things yet as I understand its a bitch to store and somewhat of a bitch to prepair so if a quick an simple cold lab prep could be worked out it could indeed open alot of doors. I remember attempting my HgCl2 experiment in the past because someone told me that you could not generate AlCl3 by gassing withou AlCl3 being present to act as a catalyst so I had to prove that a catalyst of somesort could be generated in situ and I do believe my testtube results confirmed some reaction took place.

I don't think the sublimation method is prohibitive at all to some of the hundred post + members on this board. They could undoubtedly run it with their eyes closed and one hand behind their back, literally.

I just thought I'd try the DCM method because it'd been discussed on other forums.

Additionally, generating a catalytic amount of AlCl3 is NOT hard at all. If you can't do that, you certainly shouldn't be messing with HCl(g) in the first place, or any of the downstream options. I've only chosen a cold DCM method to see if it'd work, not because I don't know about the hot methods. I'll probably use the hot methods to get clean AlCl3 if further gassing doesn't work.

One thing that concerns me is that switching the stir bar back on produces a suspension similar to the original Al suspension. There won't be any intermediate Cl 'salt' there, so that indicates that a potentially large amount of HCl(g) went to the scrubber, as a large amount of Al elemental remains. Not an issue once you have some catalytic quantity of AlCl3 I know, but it just itches at me is all; that gas isn't free. I have a few grains of AlCl3 left from the CP sample, but thought I'd try this without it first.

I wonder about heating and UV however.

Keep going man before you know it I maybe working riht along side of you to tackle these various means but first I have to finish my new lab bench which I making a fume hood for to YEHHHHHH I feel so excited for the first time in years.

I would love for you to do so. I've been working away with these acids for half a decade +, but they so often yield tar against multi-functional group substrates.

I post a lot on other forums, but have been avoiding this one for a while as it's so intolerant of stupidity. But I've come to appreciate why it's so and that there are some far more knowledgeable people here. I am very happy to be chatting with such intelligent people for a change. I've spent so many years being the only person trying this, it's really annoyed me.

This method is CONSTANTLY ranted on about on forums, yet no one seems to ever actually see if it works. I can tell you, having spent over half a decade messing around with it, that the majority of journal entries don't seem to function when strong lewis acids are transplanted onto delicate substrates. They work for single functional group substrates, not multi. On multi-group substrates, the Lewis acids seem to go after anything that'll give them electrons. I was called modest peach and a doormat at school, but I have never seen someone actually running these reactions for themselves.

There seems to be some odd things occuring with the more powerful hard acids (I think they're interacting with the phenol / tail double bond). I expect the balance of hard / mild is important and that a soft base pusher is not all that essential. A soft acid also doesn't seem to work, as it won't attack the hard base site of the ether linkage without microwave irradiation; which shifts the charges around in the substrate provided there is an electron withdrawing group on the tail (perhaps, like a double chloride, which can be transformed to a ketone with a base).

I have a sample of material that I suspect to be the demethylation product. This sample is strongly oxygen / light sensitive from my initial impressions. It reacts, strongly, with a FeCl3(aq) test for phenols.

Alone, it'll turn from beige to dark brown in hours. In it's raw, fresh state, it looks exactly like sand or mild curry. As it's exposed to light / atmospheric oxygen, it turns dark brown. I say, 'seems to yield tar' as previous tests yielded something like intractable tar. Yet I'm fairly sure, now I have the suspect product, that it can be easily mistaken for tar once it touches the ambient atmosphere or UV for any significant period (catechols are often used for photosensitive chemistry). I'm not sure which is most important at present. I'll try vacuum distilling it as soon as I have some B14 available, which will hopefully yield a white result.

The sample stinks of coal tar, and it'll easily solidified in the freezer and flows at room temperature. I suspect some of the substrate has passed over and that the sample needs cleaning up. I washed with cyclohexane during filtration, then extracted with ethyl acetate. Again, a vacuum filtration should clear this up; I'd rather vacuum distill than recrystallize, given how little seems to be known as fact regarding this material.

However this is, again, off topic. This is about lewis acid not demethylation.

Also, you don't need a fume hood for this kind of work. You can accomplish a better effect by sealing your joints well and scrubbing the exhaust through a strongly basic solution. I have done this in person. It works perfectly. Absorption is not the same as neutralization via a strong reaction. Wash heads are incredible, better than fume hoods for a lot of work; you can reduce things like nerve agent to toothpaste with a wash head, a fume hood won't do that (wash heads are also faaaaaaar cheaper, can be stored away easily and can be used for other reactions).

I would recommend storing AlCl3 (or any other similar, hard Lewis acid) in a glass bottle and wrapping the threads with PTFE tape. My CP bottle had strongly discolored it's own label over year, indicating that the CP seal had failed to some extent. Gas PTFE on the threads and a PTFE cap liner would almost certainly fix that.

PM me with an email address anyway. I've already gone far enough off topic, I'm not going any further in this thread.

I want to try distilling this DCM to see if watson's chloroform idea is right. SEDIT, please PM me and buy some fritted wash heads.

[Edited on 6-7-2010 by peach]

Sedit - 6-7-2010 at 12:16

"That took hours to start, but you can see it's running. It may be that the initial AlCl3 simply takes a while to start forming"

Sorry don't hve much time to fully read whats typed but this caught my eye. Yes it does take AlCl3 a whil to start forming and once it does it acts like a catalyst hence the reason people say that the solvent method is ineffective unless you already have AlCl3 to initiate it. However AlI3 should serve the same catalytical use and I found last night that when I2 is added to Al grain in dry DCM almost right away there was evolution of bubble(don't really know what these are just yet more then likely an artifact of trace H2O reacting somehow) and the Deep purple color of the DCM started to fade to a more redish color. Sadly I broke mylast test tube (damn this batch was cheep) just crackin the corner so I quickly sealed it with tape and am going to feed HCl into it later tonight to see if it indeed consumed the Al in a reasonable manner.

[Edited on 6-7-2010 by Sedit]

Chainhit222 - 7-7-2010 at 02:47

I recommend you buy some metal glassware clips. I bought mine on expediglass.com. I use them for all synthesis which involve acid vapors. Kech clips LOVE to fall apart from HNO3.

peach - 10-7-2010 at 09:16

I found some PTFE clips, but they're £7. EACH

Having left this reaction to stand in an amber, PTFE liner bottle, I discovered the following

The unreacted aluminium has sunk to the base, but there is very clearly a white precipitate. Which I am virtually certain will be the anhydrous AlCl3. The camera doesn't show the definition between the two colours well.

Once I have some more sulphuric, I'll regas this and get the rest of the elemental aluminium done.

It's prooving more and more difficult to find. It used to be on sale at B&Q at 96% for drain clearing. Then, last year, some woman poured it down her plughole (probably with her face over it) and received a large amount of it back. Hot.

She was severely burned and B&Q have pulled all of it off the shelves. I had a scout round in there last night and they're getting rid of any of the more pure, useful chemicals. The brick cleaner now proudly state 'Does not contain hydrochloric acid' and the only one that does is full of detergents. Sodium Chlorate greenery killer seems to have also gone. I still have a few bags of that, however.

The acid had it directly stated on the pack that it shouldn't be added directly to water, that'd it get hot and react. I doubt she bothered to read that.

Some google searching of sulphuric based drain cleaners reveals a worryingly large number of posts that are unsure about how to use it. Perhaps it was worth taking off the shelves. Doesn't make my life much easier though!

[Edited on 10-7-2010 by peach]

aonomus - 10-7-2010 at 12:42

To remove that nasty brown stuff from the AlCl3, you could always try to use a soxhlet with DCM, you wouldn't need much heat, but you'd need a really good condenser.

peach - 10-7-2010 at 21:41

That's an idea.

It may take a while. DCM boils not a lot above room temperature, as you undoubtedly know. I tried flushing that sample with very large amounts of DCM, around 500ml's worth for 25g of it, with vigorous shaking between rinses (at around 25C) and the brown would simply stain the solvent, but remain thickly in place on the grains.

When the stopper jammed, I was forced to let it sit in the DCM for a while. And that seemed to be the best option in terms of effort, solvent use and removal of the muck, to literally let it soak in just enough to cover it with a shake whenever I walked past over the next three weeks.

Sublimation may work. If it's some iron compound, the aluminium will sublime at much lower temperatures.

I don't have anymore of that CP grade acid left, and I won't be buying anymore. The convenience of having it turn up premade is outweighed by the number of variables it induces in the kind of things I'm using it for, which is experimental work.

Or as we used to say "It's pure, whatever it is!"

That reminds me of the Simpsons episode where the Mafia are selling milk to the schools, sucked out of rats, and Chief Wiggum busts the place then licks one...

"It's pure!"

[Edited on 11-7-2010 by peach]

peach - 15-7-2010 at 15:14

I regassed the material above today.

It did start reacting again very quickly but, after putting a significant amount of gas through it, I still don't see the acid.

I'll filter and clean it up tomorrow, but I'm not all that hopeful.

I have already used far more gas than is theoretically needed. A fritted wash head would be a good thing to have, but I have used molar excesses so far, adding it at the slowest rate my funnel will drip the conc. H2SO4 onto the salt, and don't have the result.

I had to extend my wash head's down tube with a little neoprene tubing. After a while, the solvent began turning brown. I'm not sure if that's contaminants in the Al source or if the neoprene has been chemically attacked. It does seem remarkably similar to the brown of the commercial acid, just not as dark. The tubing doesn't seem to have degraded, but I'll inspect it closer tomorrow.

I'm sick to death of running this cold solvent method. I should easily have a catalytic amount of the acid by now. I don't think it's particularly viable. I think a well thought out hot method is the way to go.

[Edited on 15-7-2010 by peach]

Sedit - 15-7-2010 at 16:25

Why does the solvent have to be cold?

peach - 16-7-2010 at 01:23

It's not actually cold (this time anyway). By the time it'd started running again it was gently refluxing with it's own heat.

I suppose I could use something like cyclohexane, which has about double the BP and would probably be the solvent I'd be using later anyway, but... I don't think such a none polar solvent is going to pick the HCl(g) up very well. And there already seems to be an issue with wasting gas and a solvent method.

When I store up the energy to do it again, I might try it with a sprinkle of mercury chloride in there from the start. But I think it's going to have the same issues. I suspect there is already enough salt to catalyze, as I'm reasonably sure that's what I can see in the bottle photo. Also, if you can get mercury chloride, I doubt you need to make your own AlCl3.

I have a question for the guys who've tried blowing it over hot Al alone. Does it tend to cake up and stick to the glass where it solidifies again? And am I going to have problems piping it around if I insulate the tubing? I'm not doing anything extreme, I'd just like to simultaneously sublime it out of my test tube and to somewhere easier to get at it. I'd like advice from someone who's done this first hand, because it's a major pain in the ass when things clog up when it involves gases and concentrations of acid like this.

kmno4 - 16-7-2010 at 10:08

Your posts are too complicated for me, too much reading about unimportant and off-topic details (sorry).
If you ask about reaction of hot Al with HCl - I did this many years ago. Condensed AlCl3 does not stick to glass and can be removed from glass pipe with aid of curved steel wire (as I did).
With pipe with not too samall diamater and on laboratory scale there should not be any problems
Besides it is very nice reaction. You can read a little more about it in Brauer.

BTW.
I would not try to do this in DCM at all. I would be affraid of CH2Cl2 reaction with Al. It is known that CH2Cl2 is using as solvent in organometallic chemistry but many halo-methanes react with Al, Zn or Mg. CH2Br2 gives violent reaction with Al.
Your coloured solution may be just result of such reactions (but not violent) giving some polymeric shit (especially during prolongated time of contact Al with DCM).

peach - 16-7-2010 at 11:10

The far too ghetto setup for my liking. The glass is all off for repairs and modifications next week, this isn't what I'd usually be happy with and it doesn't work all that well. Note the ACS grade HB pencil. This is a lazy setup, don't copy it verbatim. There should be drying tubes and siphon traps on there, not just because health and safety madness says so, but because it doesn't work so well like this

The wash on the vent is full of strong KOH solution to neutralize any HCl(g) trying to wander off. It works perfectly well. Where HCl(g) escapes, is the seals. They need fully greasing, and preferably with that chemically resistant Molykote (I think it's called)

Peach finally has RBFs

Fizzz.... this method isn't so bad, but it's not perfect either. I would certainly prefer a known, bottled quantity. I'm considering either making my own liquid storage vessel from 316 or renting a cylinder.

Death acid. I thought it was pretty the way there were tentacles falling from the down tube.

This is why wash heads with long stems are utter balls. They also seem to encourage siphoning. If anyone can explain the angles on QuickFits heads like this, please elucidate. I note they've stopped doing this. I'll be having the glassblower rip them off, straighten them and fit barbs or screw threads, probably barbs.

That's neoprene. I've used it before with HCl(g) and it's been absolutely fine. I checked for compatibility. Cole Parmer don't have any information on it, and other sources say it'll be fine with HCl. It's chlorinated and the poor mans PTFE. It's also annoying to work with for a number of reasons. I can't see what's inside it, so I can't stop a problem until it's about to drop through a wash head. I can't tell if it's clean when I wash it out. It doesn't 'flow' very nicely in terms of flexibility and it tends to stick to the glass when it gets warm. It also has quite a 'rubbery' , 'grainy' surface, which encourages things to gather up in it. I'll be going with tygon or similar PTFE like tubing in the future I think. The neoprene swells up in DCM, but that's fairly normal and not chemical attack; PTFE will do it. But, as you'll see, things don't go quite so to plan, and this tubing may be to blame

DCM attemping to escape. It's yellow because there's FeCl3 muck in the tubing. I surrounded the tube with ice when I noticed white fumes coming from it, which was the DCM escaping. It's possible some of that is Watson's chloroform, I still have to distill the solvent. But I'm not entirely sure if I want to be bothered with this one. And I've thrown out the contents of that wash tube

Either the Al or the neoprene is fizzing a lot more than it appears here. The flask is about as warm as luke warm tea, and I can see DCM refluxing off the wash head. So something is reacting

The next day, I put more excess HCl(g) through it. There is a distinct, scary, 'burnt rubber' or 'rotting egg' smell abount the exhaust, which is certainly not the solvent or HCl(g). I am concerned

I'm not gassing it anymore, it's had it's chance. Note the discolouration of the solvent. Impurities in the Al (unlikely) or the neoprene reacting (more likely)

The particulate has gone from stirrable to caking the walls. It's the wrong colour and it's the wrong smell

But it does fume like crazy when I open it to the atmosphere; far worse than AlCl3. But the fumes are incorrect. It's not HCl(g), which it should be if it was AlCl3. The fumes stink of burning rubber or egg. I am now concerned and wish I had a fume hood, as I have no idea what that is or how carcinogenic it is

The neoprene. Despite being covered in something crystalline looking and the odd smell, the polymer seems unaffected. It's swollen, but that's the solvent. It's not rotted away, degraded or brittle. The coating on it does aggresively fume on exposure to water but, again, the smell is entirely wrong

Trying to get a clear picture of the insides is very difficult. In this picture, I'm blowing through one of the necks. AlCl3 fumes on exposure to moisture, but this is doing it far worse. It's possible the yellow is from traces of FeCl3 in the tubing, but there's more there than I'd expect from traces on the inside of the tubing. That wasn't there when I was using pure glass. What could the neoprene be producing that'd steal chlorides from the Aluminium and fume more aggressively than AlCl3? With such an odd, clearly not HCl(g), smell and behavior?

The only thing I can do is hit it with the hose. Again, this is where I'd like a fume hood, this is an unexpected and uncontrollable result. The photo does this no justice at all, it was much worse than it looks here. There were clouds of white, persistent, burnt rubber, sulphur, egg smelling fumes pouring out of the flask for an hour or two. It was so bad I thought someone was going to call the fire engines out. Worse, I tried 'misting' the fumes out of the air. If it was HCl(g), I would certainly recognize it and the water would knock it down. This definitely isn't. I'm worried about what that was, and I'm not sure if I want to know! :mad:

FAIL! NO WHITE PRECIPITATE, NO HCl(g) FUMES ON EXPOSURE TO THE ATMOSPHERE, MASSIVE EXCESSES OF HCl(g) TO REACH THIS STAGE. IF THIS WORKS, I DOUBT IT'S MUCH EASIER THAN COOKING THE AL. AND I SEVERELY DOUBT IT'S ANYWHERE NEAR AS EFFICIENT IN TERMS OF HCl(g). NEOPRENE APPEARS TO BE INCOMPATIBLE WITH HCl(g) WHEN IN DCM. MAKES ME WONDER HOW THIS WOULD HAVE WORKED WITH PURE GLASS. BUT I'M DONE MESSING AROUND. GO HOT METHOD!

[Edited on 16-7-2010 by peach]

watson.fawkes - 16-7-2010 at 11:39

Sulfur crosslinking holds vulcanized neoprene together. Even though there's no sulfur in the chloroprene monomer, sulfur is added to the polymer for vulcanization in order to stiffen or harden it. If you're smelling sulfur compounds like H2S, it's coming from degradation of your tubing. Given that you had a piece of it sitting in close contact with a Lewis acid, it's hardly surprising there may have been some depolymerization or other attack. If you knew the smell of chloroprene, it would help diagnose what happened.

blogfast25 - 16-7-2010 at 13:04

See, at this point I'd quit while you're ahead, if I was you.

This whole experiment reminds me of another thread in which a serious and knowledgeable experimenter wasted hours trying to follow a patent that... wait for it... called for the reduction of KCl with Mg at RT in organic solvent media. It too sounded about as plausible as Cold Fusion, it too didn't work. I'm all for thinking outside the box but it doesn't always lead to satisfaction.

Diligent effort though... :cool:

[Edited on 16-7-2010 by blogfast25]

entropy51 - 16-7-2010 at 17:17

Given the simplicity of the 120 year old methods for successfully making AlCl3 in the lab, this thread is well, inconclusive. Gassing Al with dry HCl works, every time.

[Edited on 17-7-2010 by entropy51]

Sedit - 16-7-2010 at 17:21

Entropy there are a gazzlion writeups spaking of the solent method and many claim success. I myself have had limited success as well but I think his down fall here is the use of DCM bar non.

Th solvent method has a quaint charm about it that you dont get with the ho method and I would love to see this synthesis become fruitful.

Ill share a bt of stuff monday, nothing important really just a few notes that seem strange with DCM as the solvent.

aonomus - 16-7-2010 at 17:52

I think one of the problems with the recent experiment was the rubber tubing contaminating the reaction.

If you're desperate for a dispersion tube of some sort, check out your pet store. I did find what appeared to be a HDPE plastic frit which *may* be solvent resistant. (in the form of a plastic air 'stone' for fish tanks)

peach - 16-7-2010 at 20:05

Sulfur crosslinking holds vulcanized neoprene together. Even though there's no sulfur in the chloroprene monomer, sulfur is added to the polymer for vulcanization in order to stiffen or harden it. If you're smelling sulfur compounds like H2S, it's coming from degradation of your tubing. Given that you had a piece of it sitting in close contact with a Lewis acid, it's hardly surprising there may have been some depolymerization or other attack. If you knew the smell of chloroprene, it would help diagnose what happened.

Thanks watson! I did think sulphur had to be involved somehow, given the smell. I don't think it was H2S it's self, it wasn't rotten eggy enough, but there was definitely 'that smell' about it. The predominating smell was burning rubber. As I say, I've used it around HCl(g) before and not had problems. But the combination of the solvent or the Lewis acid has not helped. There was an awful amount of curious fuming (more than I'd expect from AlCl3), lack of HCl(g) and that horrible smell upon exposure to moisture. The yellowing in the solid could also be due to something sulphur based. I was hoping I might find some of your chloroform in this, and wanted my DCM back anyway. But I'm sorry, I'm going to have to abort on redistilling this one.

Given the simplicity of the 120 year old methods for successfully making AlCl3 in the lab, this thread is well, inconclusive. Gassing Al with dry HCl works, every time.

[Edited on 17-7-2010 by entropy51]

I bet gassing hot Al alone works like a charm. Not too sure about the solvent methods. If anything, it seems to be fairly inefficient in terms of absorption. I may repeat this when my glassware is finally back and I can do it with an all glass setup, as things changed when that piece of neoprene came into contact with the DCM + possibly Lewis acid. It wasn't doing this earlier, when I was only using it to direct the dry HCl(g) and exhaust around.

Entropy there are a gazzlion writeups spaking of the solent method and many claim success. I myself have had limited success as well but I think his down fall here is the use of DCM bar non.

Th solvent method has a quaint charm about it that you dont get with the ho method and I would love to see this synthesis become fruitful.

Ill share a bt of stuff monday, nothing important really just a few notes that seem strange with DCM as the solvent.

You're taking the weekend off from Science Madness?! How dare yar! Yar Land lubber YAR!

But seriously, I'd really like to hear those observations. I thought DCM would be quite a nice choice, as it seems to pick the gas up well, doesn't react with the Lewis acid and the acid isn't particularly soluble in it either, making recovery easy.

Apart from avoiding the need to heat the glassware up, I would also like this method to work as it seemed to be producing fine white precipitate earlier on. Which would make recovery and subsequent measurement easy; not having to scrape sublimed and then resolidified lumps off the glass or grind them up would be good.

I do suspect, however, the failure is the result of that piece of tubing.

See, at this point I'd quit while you're ahead, if I was you.

This whole experiment reminds me of another thread in which a serious and knowledgeable experimenter wasted hours trying to follow a patent that... wait for it... called for the reduction of KCl with Mg at RT in organic solvent media. It too sounded about as plausible as Cold Fusion, it too didn't work. I'm all for thinking outside the box but it doesn't always lead to satisfaction.

Diligent effort though... :cool:

[Edited on 16-7-2010 by blogfast25]

This is the point where I tend to agree. Watson's heard me going on about toxic versus harmful or irritant, but clouds of odd smelling fumes pouring off polymers are generally not great at all.

There is ZERO chance I'll be using neoprene in contact with the DCM + acid again.

There's a patent out there about Wack'ing eugenol through to a ketone. It actually lists tens or a hundred plus different substrates and is essentially someone listing off everything with a double bond that they've Wackered with Methyl Nitrite. I tried their eugenol method, down to the letter, and just kept getting brown sticky gunk. I even tried swapping the oxidizers around. No luck.

It's worth noting that a common 'trick' when patenting something is called 'ring fencing', which is to try and list off as many other possibilities as possible to block others from making slight changes and repatenting it. Having distilled my third or four batch of suspect tar, I began wondering if they'd actually run this experiment or just made up the numbers; or to be kind, omitted some important detail. I also discovered a Tetrahedron entry explaining why it wouldn't work for eugenol.

I don't remember where this solvent suggestion came from in particular. I simply saw it being mentioned in forums and thought I'd give it a go.

My brother happens to be an IP lawyer, so PM me if you need some free legal advice for patenting your new super product.

Quote: Originally posted by aonomus

I think one of the problems with the recent experiment was the rubber tubing contaminating the reaction.

If you're desperate for a dispersion tube of some sort, check out your pet store. I did find what appeared to be a HDPE plastic frit which *may* be solvent resistant. (in the form of a plastic air 'stone' for fish tanks)

OooooOooo no.

Even having experience using this tubing with this gas, sucessfully, I did find myself questioning how wise it was to use it within the reaction space it's self.

I needed to do so as my wash bottles are still in the post and the downtube wouldn't reach into the solvent.

I've heard from people using aquarium frits before, and having them fall apart on exposure. Even borosilly frits will fall apart around strong bases.

A lesson I didn't really need explaining again, it has to be glass all the way.

When the glass turns up, and my patience has leveled up to normal again, I may try it using fritted wash bottles, as it would be good if it worked. But I'm not improvising anything involving the DCM + acid combination with regards to glassware.

Thanks for the replies everyone. It takes a lot of effort to gas something for hours on end, then do the washing up, then the photos, then the post, so it's really good to hear feedback.

Now I'm off to get a massive mug of tea (at 4am), recover from all those beers and eat my syrup soaked banana fritters from the chippy, which have probably gone even soggier by now. And maybe shoot that magpie if it doesn't shut up.

It's going to wake the damn rock doves up, I'll never hear the end of it then.

"Coo Cooooo Coo" "You said that last time"

[Edited on 17-7-2010 by peach]

kmno4 - 17-7-2010 at 08:52

....there are a gazzlion writeups spaking of the solent method and many claim success....

"gazzlion writeups" - it can be interesting.
Can you give here some references, links, for let's say... eeee... 2 ones ?
Of course,please do not give any patents, nobody is interested in rubbish.

IrC - 17-7-2010 at 12:01

I think the ones saying this approach is a waste of time are missing the point. I think she is doing the right thing, experimenting and learning. How many lights did Edison try? If he quit after 3 or 4 we would have Russian bulbs (since there was an inventor over there inventing the same thing during those years) and Edison might not be so well known? For all we know Peach will come with a new procedure or some new catalyst which makes her solvent method work? If you do not try you never succeed?

Note I said she referring to Peach as it is the only way this quote makes sense.

"If you like that, you'd love my underpanties. They won me an award, $20, for a photo of them, which I believe has now been publicly circulated."

Thinking about reactions possible I do know most of the possible products like to eat organic tubing. DCM alone is going to do this, especially if hot. Years ago I bought DCM in 5 gallon buckets to depot laser power supplies which to keep you out were a mixture of various silicone compounds and sand. It would eat anything I tried from latex to neoprene to taigon tubing (unsure of it's composition), in general everything you could buy in the form of rubber or plastic so she needs to work on that part of the experiment. I quit using DCM after monoxide nearly took me out. It would eat every glove I bought eventually (usually sooner than later) allowing the stuff to leak in and stay in skin contact as well as breathing the fumes even though I really tried to have positive ventilation. I am sure most know DCM in our body converts to carbon monoxide carried by the blood from either inhalation or skin contact. So you should really have someone to monitor things since by the time you realize it you could very possibly need a blood transfusion to survive.

I imagine some of the reaction products can eat her tubing faster than DCM alone yet reading the quote below from Wiki I wonder about temperature and if these reactions listed below can all occur. One would think she did not have enough heat, since 400–500 °C is stated?

From Wiki:

Dichloromethane is produced by treating either methyl chloride or methane with chlorine gas at 400–500 °C. At these temperatures, both methane and methyl chloride undergo a series of reactions producing progressively more chlorinated products.[1]

CH4 + Cl2 → CH3Cl + HCl
CH3Cl + Cl2 → CH2Cl2 + HCl
CH2Cl2 + Cl2 → CHCl3 + HCl
CHCl3 + Cl2 → CCl4 + HCl

The output of these processes is a mixture of methyl chloride, dichloromethane, chloroform, and carbon tetrachloride. These compounds are separated by distillation.

I like your methodical well documented work Peach keep it going as to me it is interesting reading and fits in line with the reason we love to experiment.

blogfast25 - 17-7-2010 at 12:57

I think the ones saying this approach is a waste of time are missing the point. I think she is doing the right thing, experimenting and learning. How many lights did Edison try?

Let me tell you this, son, and I'm pretty sure Peach will agree with me: the difference between Peach and Edison was that he had something to go on: an electrical resistance heats up when a current passes through it. From there to a working prototype was hard technological work.

In the case of RT production of AlCl3 we know that without a catalyst to strongly reduce the activation energy of the chlorination reaction no reaction could take place without violating the Laws of Physics. Don't take my word for it: read up on 'activation energy' and atomic theory in general.

If Peach wants to continue this line of investigation her first line of enquiry must be the development of a suitable catalyst. That requires fundamental understanding of the chlorination reaction mechanism at the atomic/molecular level.

If she can pull that off, hats off and in the air to her. But until that becomes the approach, AlCl3 at RT in solvent media is like flogging a dead horse: it WILL NOT come alive by sheer will power.

In essence all she's doing right now is leading dry, cold HCl over dry, cold Al: the DCM is nothing but a spectator. We know that that doesn't work, if anything does her works proves that very convincingly...

[Edited on 17-7-2010 by blogfast25]

peach - 17-7-2010 at 13:25

As far as it not working, have a look at the first set of photos and the photo of the amber bottle. That was elemental Al, dry HCl(g), no catalyst and all glass at room temperature. The solvent is green and the aluminium is reacting. When the reaction is finished, the green is gone and the solvent is opaque. There is pure white precipitate in the solvent the next day, separating from the aluminium layer, and the solvent has cleared to transparent.

I'm open to suggestions as to what that is if it's not AlCl3. Your possible reagents are elemental Al, HCl(g), DCM and borosilicate glass.

The solvent does make a difference, by changing the availability of the gas. E.g. concentrated acids often don't attack things until you add a little water to disassociate them. What happens if your pour hydrochloric acid on Al?

The evidenced is given by that piece of tubing. Neoprene doesn't react with dry HCl(g). None of the tubing before that piece has responded in a manner anyway near similar. I've used it before with dry HCl(g) and it doesn't react like that. Look at the tubing before and after the flask, it's fine. In fact, the tubing in the flask is fine as well. Only the portion below the surface of the liquid seems to have reacted.

But it does in the presence of DCM or the lewis acid. Either way you look at that, you're point is wrong. If it was the DCM that indicates the importance of the solvent in changing the reactivity of the gas. If it was lewis acid, that means I formed some of the acid in the first instance.

So no, based on that factual evidence and some logical reasoning, I don't agree.

the difference between Peach and Edison was that he had something to go on: an electrical resistance heats up when a current passes through it. From there to a working prototype was hard technological work.

An ironic statement to make given the above observations.

As to Edison, he 'borrowed' the observations of one of his students with regards to thermionic diodes.

[Edited on 18-7-2010 by peach]

blogfast25 - 18-7-2010 at 07:54

I'm open to suggestions as to what that is if it's not AlCl3. Your possible reagents are elemental Al, HCl(g), DCM and borosilicate glass.

[Edited on 17-7-2010 by peach]

Did you isolate/analyse the white precipitate? No, so we don't know what it was. I've seen a lot of 'precipitates' that turned out to be a trick of the light or pure perception. Recently I saw a green precipitate I thought might have been PrCl3. Couldn't prove it though... Maybe next time.

What you need indeed is a proton transferring solvent, much like water in a sense. A Lewis acid indeed. Or a Lewis acid than can be added to the mix or is soluble in a particular solvent. That's what's needed. I don't think you're close, TBH...

[Edited on 18-7-2010 by blogfast25]

peach - 18-7-2010 at 09:38

I'm no where near some of the guys on this forum in terms of theory, but you're no where close to me in terms of being able to use your own senses; like your eyes to look at those photos and your mind to interpret what might be happening.

You have Al, HCl(g) and glass to work with and you're still telling me I have something that comes out of a green solvent, turns the solvent opaque (like a salt falling out of it), then clears to white and crystallizes to a white ppt and it's not AlCl3?

I get the feeling there are a lot of arm chair chemists here. Tell me, specifically, what else could have formed. That is both white, precipitated form aluminium and involved those reactants. You're a brave boy to be telling me what I observed.

If you can't reply with a specific answer, you loose. If you don't reply, you loose.

The fact is, those initial pictures are of AlCl3 forming.

I asked you specifically to tell me what it was if it wasn't AlCl3 forming, and you entirely ignored me. I have prodived the evidence. The reagents, that the solvent was green and reacting, went opaque and then cleared, released a white precipitate (AlCl3 is opaque white). Where's the alternative answer?

A trick of the light? Come on! Even those photos show the Al is separating from the precipitate. And they're poor compared to the reality.

Tell me what it is if it's not AlCl3. You can't, because it is. Because it's the only thing that could have formed from those reagents. Stop flogging the dead horse that is you ego. You're an armchair chemist, and quoting my own results like they're yours.

I pointed out the neoprene tubing. The solvent and lewis acid. Both ignored.

Have you tried dissolving a Lewis acid in a polar aprotic solvent like water? Plumes of HCl(g) are the answer. You think that might be why I've chosen a polar aprotic? You have zero idea what you're talking about.

Water with AlC3? WHAT!? You're talking about hydrated salts. And AlCl3 in a hydrate form doesn't function as a hard lewis acid.

I have one question alone for you to answer. What is it if it's not AlCl3. Again, you have Al, HCl(g) and borosilicate glass as your reagents. If you don't answer, you're trolling.

I'm a SHIT chemist, and this is what I'm competing with?

[Edited on 18-7-2010 by peach]

blogfast25 - 18-7-2010 at 11:18

I get the feeling there are a lot of arm chair chemists here. Tell me, specifically, what else could have formed. That is both white, precipitated form aluminium and involved those reactants. You're a brave boy to be telling me what I observed.

If you can't reply with a specific answer, you loose. If you don't reply, you loose.

The fact is, those initial pictures are of AlCl3 forming.

I asked you specifically to tell me what it was if it wasn't AlCl3 forming, and you entirely ignored me. I have prodived the evidence. The reagents, that the solvent was green and reacting, went opaque and then cleared, released a white precipitate (AlCl3 is opaque white). Where's the alternative answer?

Oh dear, oh dear, oh dear. Feeling a little piqued, Peach? Arm chair chemists? As opposed to you producing stinky goo and a white precipitate you claim to be AlCl3?

Remind yourself not to become a scientist: you'd be in for far greater criticism than I have meted out. In science we don't 'lose' or 'win', not if we're honest about things.

You've not 'prodived' the evidence: the material wasn't isolated and characterised. Why not? Not sure how to? Too much of an 'armchair chemist'???

Please also tell me how Al + HCl + DCM + borosilicate glass result in 'something green'?

IrC - 18-7-2010 at 12:43

blogfast25 please explain why you are attacking this nice lady for coming here, spending long hours experimenting, taking pictures, documenting, posting here, which I thought was the reason for this site? She did her best, asked others of greater knowledge for their opinions, yet the gist of what I get from your posts is you are calling her an idiot who should give up chemistry.

You think your head is so very big.

Almost forgot to mention but it looks to me that unlike yourself she is actually doing something in her quest for learning, something you are not doing. Why don't you do the experiments yourself if your head is so very big and show her proof of her error if it exists.

[Edited on 7-18-2010 by IrC]

blogfast25 - 18-7-2010 at 12:51

you are calling her an idiot who should give up chemistry.

[Edited on 7-18-2010 by IrC]

And where did I do that, dearest? Learn to read.

Fleaker - 18-7-2010 at 12:53

People are really convinced this can't be done in solution?

People are really convinced that HCl gas (let alone chlorine) doesn't attack aluminum ?

Might try this in chloroform, with a small crystal of I2 (or droplet of Br2) added--that would be my choice for performing this reaction.

[Edited on 18-7-2010 by Fleaker]

IrC - 18-7-2010 at 13:01

you are calling her an idiot who should give up chemistry.

[Edited on 7-18-2010 by IrC]

And where did I do that, dearest? Learn to read.

I have been reading for almost 60 years. Along the way I learned it is not just each individual word but rather how they are structured to convey a point. From "Remind yourself not to become a scientist" You've not 'prodived' the evidence: the material wasn't isolated and characterised. Why not? Not sure how to? Too much of an 'armchair chemist'???

This is insulting to her or to anyone this would be spoken to. How is armchair in the picture looking at the well documented work she posted here. Where is your work?

One can only assume after reading the lengthy and hard work she posted you are saying she should give up. Only an idiot would do so having walked as far as she already has along the path of greater enlightenment so one can only surmise you are calling her an idiot. You are trying to discourage her and this pisses me off.

The problem is not my ability to read it is your ability to speak. Or write.

[Edited on 7-18-2010 by IrC]

blogfast25 - 18-7-2010 at 13:03

Quote: Originally posted by Fleaker

People are really convinced this can't be done in solution?

People are really convinced that HCl gas (let alone chlorine) doesn't attack aluminum ?

Might try this in chloroform, with a small crystal of I2 (or droplet of Br2) added--that would be my choice for performing this reaction.

[Edited on 18-7-2010 by Fleaker]

The people aren't convinced by this series of experiments, Fleaker, that's not the same thing.

Where would the I2 come into it? Soluble in CHCl3 and thus an initiator? It's possible but I've never seen it...

blogfast25 - 18-7-2010 at 13:05

The problem is not my ability to read it is your ability to speak. Or write.

[Edited on 7-18-2010 by IrC]

And I think you've got very long toes. And that pisses me off...

IrC - 18-7-2010 at 13:12

Your honor, the defense rests it case on the last testimony by the defendant.

Thinking out loud

Fleaker - 18-7-2010 at 13:19

No, there are many things wrong with this set of experiments.

I2 is very soluble in chloroform, and will certainly react with aluminum--it (and bromine) does so neat and vigourously to give the polymeric halides. Even tin metal reacts with iodine in refluxing toluene, and Sn is much less active than aluminum. I suspect this reaction would be better done using chlorine gas, particularly if one adds a catalytic amount of mercuric chloride to remove the oxide. Perhaps also a better coordinating inert solvent is needed--issue being, few solvents are inert to a potent Lewis acid. I think iodine crystal would promote the formation of the aluminum trichloride.

Still, kmno4 said something quite salient--that the solvent may react with the metal to form an organometallic, and polymerize or react with trace oxygen. This could be some of the darkening that was seen by peach and explain any ebullition in the solution rather than hydrogen gas.

I've never tried to make AlCl3--it's so cheap and commercially available. Shipping is more expensive than the damn compound.

EDIT: Let's all keep it civil, friends. We're here to discuss science, and our respective failings in our methods, not showcase immaturity. Let's not insult each other, nor peach's efforts to at least attempt this reaction. Peach has done more (home chemistry) than I've done this past year.

[Edited on 18-7-2010 by Fleaker]

12AX7 - 18-7-2010 at 14:48

Blogfast25, I thought you were better than this.

Sedit - 18-7-2010 at 19:10

OK someone requested the work I said I would add but its late and tommorows monday so I got alot of shit to get ready in the morn.

But What I think I did conclude is DCM sucks. I used Napthas for my first test I spoke of earlier with catalytical amounts of HgCL2. Howver using DCM and an attempt a I2 as the initiator give me completely different results. AS the DCM evaporated I expected to see some fuming of somekind and perhaps I2 leave the test tube.

However I got a nasty thick purple goop instead that would char indicating polymerization. The iodine was locked in this polymerization as well if that helps anyone conclude what might be going on. No amount of heating this goop would free any iodine omly cause charing of the goopy mess.

Im going to give it a larger scale run just to put and end to it but I must setup first and could take me a week before I prepaired because I want a control, I2 initiator and one with HgCL2 as initiator.

For references I talked about look over at Erowid there are more then a few claiming success and they provide the references and patents at the end of the writeups as far as the solvent method is concerned.

IrC - 18-7-2010 at 20:37

I know I am not the chemist most here are and I still have not figured out the green color Peach mentioned, but is it possible the white solid is aluminum hydride?

Possibly this site has info you can use on the subject of Catalysts.

National Centre for Catalysis Research http://203.199.213.48/

While I know the experiment is not the same, he is using DCM and AlCl3 and talking about a green colored result, possibly there is something in the following thread which may give you ideas about what is going on in yours? Whatever reaction is occurring, as someone else mentioned I think the tubing involved in the reaction should be considered.

http://www.sciencemadness.org/talk/viewthread.php?tid=11607

[Edited on 7-19-2010 by IrC]

Eclectic - 19-7-2010 at 02:30

What happens if you use Et2O? As I recall the adduct of AlCl3 and Et2O is stable and distillable.

http://www.freepatentsonline.com/3843774.html

[Edited on 7-19-2010 by Eclectic]

Sedit - 19-7-2010 at 05:51

There is a thread somewhere that speaks of using hydrated AlCl3 in DMSO to form an adduct which anhydrous AlCl3 is the decomposition product of perhaps its worth looking into. I tryed small scales of it but never got my DMSO dry enough to get it to work correctly.

blogfast25 - 19-7-2010 at 06:03

Quote: Originally posted by 12AX7

Blogfast25, I thought you were better than this.

Tim, I don't think you've really read what I wrote, at least not all of it. And some of us being called 'armchair chemists' wasn't exactly Peach's finest hour either...

As far as I'm concerned it's water under the bridge now, I'm not good with grudges

peach - 19-7-2010 at 11:02

You still haven't offered a single option for what may have been forming if it wasn't AlCl3. In fact, you haven't a single piece of actual scientific thinking, at all.

Al + dry HCl(g) in borosilicate glass effervesces and -----> white precipitate in DCM

Green solvent is a sign of chlorine or a chloride being in solution. The solution is green, then turns opaque after effervescing, then clears when a precipitate falls out of solution. And AlCl3 is not particularly soluble in DCM, and it's a white crystalline solid.

Tell me what the white precipitate is if it's not AlCl3. A very simple question, which you seem known the answers to.

This is the third time I'm asking you the same question.

I'm sick to death of this random, baseless arguing polluting my thread. If someone gave me a specific answer and said "No, it hasn't worked because of .....", I'd care. But that hasn't happened.

I don't hold grudges either. I'm too busy to be bothered remembering peoples' user names on internet forums. I'll argue with you alone, then agree with you alone in another thread. But I haven't seen any attempt to explain this, other than complaining.

I don't want any more arguing, I want Blog alone to reply with specific alternatives. And to explain why it was likely those and not AlCl3; as you seem sure it wasn't, so I assume you have evidence for that.

Knock out the arguing and look at those photos. AGAIN, FOR THE THIRD TIME, Al + dry HCl(g) in borosilicate ----> green solvent, effervescence, white precipitate.

Lets keep in mind you're the person who suggested doing this in "like water" solvent, shall we.

Give some scientifically reasonably answers as to what it is if it's not AlCl3.

I would urge everyone else to stop replying now. Give Blog alone a few days to reply.

FOURTH TIME, SPECIFIC ANSWERS.

If you don't reply with something scientifically reasonable and post some more complaining, I'll speak a moderator about you. Reply with something scientifically worthwhile or don't reply.

[Edited on 19-7-2010 by peach]

IrC - 19-7-2010 at 11:13

Does this white precip have a jelly consistancy as in aluminum hydroxide?

peach - 19-7-2010 at 11:21

No, it behaves as a solid crystalline.

anotheronebitesthedust - 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).

kmno4 - 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

Sedit - 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.

. 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.

ziqquratu - 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.

IrC - 19-7-2010 at 16:09

No, it behaves as a solid crystalline.

OK but that starts me thinking about aluminum hydride again.

Sedit - 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]

12AX7 - 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.

Tim

Fleaker - 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.

Microscale Prep of Anhydrous AlCl3

rrkss - 26-7-2010 at 03:26

Thought this might be useful as it makes a difficult reagent easy to get.

Attachment: alcl3.pdf (1MB)
This file has been downloaded 1262 times

blogfast25 - 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]

Chainhit222 - 28-7-2010 at 05:41

This method is pretty heavy on iodine consumption

Lambda-Eyde - 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?

Chainhit222 - 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:

Lambda-Eyde - 28-7-2010 at 12:50

I don't see why this reaction shouldn't occur:

Cl2 + 2CH2ICl --> I2 + 2CH2Cl2

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 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.

Does anyone with some more insight in organic chemistry have any comments?

Nicodem - 28-7-2010 at 15:20

Quote: Originally posted by Lambda-Eyde

I don't see why this reaction shouldn't occur:

Cl2 + 2CH2ICl --> I2 + 2CH2Cl2

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.

Lambda-Eyde - 28-7-2010 at 16:01

Quote: Originally posted by Nicodem

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).

Formatik - 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".

rrkss - 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.

blogfast25 - 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.

Formatik - 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.

Methansaeuretier - 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]

blogfast25 - 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.

peach - 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

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.

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.

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).

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.

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]

blogfast25 - 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...

peach - 12-8-2010 at 16:48

I realize it's very easy to spot AlCl3 once it's isolated in it's natural habitat and in the metaphorical headlights, but that neoprene entirely ruined what appeared to be working very well. No simple way to separate the AlCl3 from the unreacted aluminium sprang to mind, so I was attempting to push it forward to completion, to get said sample for testing. I also felt comfortable I could achieve that after seeing it effervescing as it did.

I purposefully stopped gassing and left it to finish effervescing on it's own. When I came back to finish it, I thought "Oky doky, I'll give it a try in one of my newly arrived, better sized flasks". That meant extending the down tube, as the flask wasn't better sized in this sense; only in it's volume. Based on how I'd seen the neoprene behave with the components separate from each other, I assumed (remembering what assumptions make you and me) it'd probably be okay to dip a bit into the solution. Wrong, wrong, wrong, I will certainly admit that critical failure, although, it was an error based on some degree of reasoning.

I don't just suspect the neoprene, I am confident it was that specific contact that ruined the attempt to finish the material. I was so genuinely concerned about the foul smelling, unpredicted gas that resulted with the neoprene in contact, I wasn't in anyway going to attempt carrying on with it. If not for that, for the fact it was now becoming obvious the tubing had caused some significant deviation from what should have been happening, and so the product I was testing would also be heavily contaminated.

As much as I go on about fume hoods being overused and safety being taken out of perspective, this is one of the few examples were I would have liked one to store the flask in whilst it gassed. As I've also said, the reason for that is not the HCl(g) (which I can easily deal with outside of a hood, using very little effort) but that the neoprene result went totally out of control in terms of knowing what the fumes were (with plastics and rubbers emitting some of the most disgustingly carcinogenic materials known to man as they degrade). As I had no idea what was coming off, I also had no way to safely scrub it, or assess it's risk potential, so assumed (this time for the better I expect) to 'run away'.

There was most definitely something in there that was violently reacting with water. On trying to rinse the flask out, there was a gigantic volume of persistent fuming and heating, which went on for hours.

But yes, in need of an all glass, tygon run it is. I am still watching this thread and working on this. It's just taking a while for me to piece everything together; I want to make sure it either fully confirms or strikes off the idea. I'm very busy working on some other things, so it's a kind of side project that's coming together bit by bit as things turn up in the post.

Once I have the glass and bits I feel happy with, I'll certainly rerun it and post up the results, whether it functions or fails. I'll make a video for you, I'm finding it's by far the quickest, easiest and most detailed way to record and convey information.

JOHN

[edit]I'm not sure what you mean by the dark stuff. You mean in the videos? That 'dark' muck will be the unreacted aluminium sat at the bottom (I was switching the stir bar off so you could see it churning it's self). The fluffy stuff floating around like little flocculated puffs of fluff will be the acid, if it's functioning. If you mean in the photos, I have no idea about that either. I have asked here before about it, and not got an answer. It's supplier supplied CP AlCl3, it's production had nothing to do with me. The dark colour of the solvent is only me rinsing it back off (which took ages, a lot of effort, and entirely seized a taper, beyond repair). That's part of the reason for me wanting to make the acid myself. I feel relatively sure I can do a lot better than the commercial material I bought, and that pure white precipitate in the all glass run hints that this may be true. I also don't like the big nodules the supplier's AlCl3 comes in, they basically don't stir, and jam up the bar, even in a lot of DCM (which an equal 'a lot' of people expect it'll dissolve fine in).

[Edited on 13-8-2010 by peach]

Magpie - 12-8-2010 at 17:43

But yes, in need of an all glass, tygon run it is.

I'd be careful which Tygon tubing I selected. The stuff you get at the hardware (DIY) store may be loaded with leachable plasticizers.

http://www.tygon.com/tygon-traditional-tubing.aspx

peach - 13-8-2010 at 06:18

Carrrr, carrrr, jack jack jack jack...

Warning is accepted, as I know numerous people are using standard tygon and wondering why it's going wonky.

However, as I say, this is the latest addition to their ultra-chemically resistant line.

Click the link to this PDF on the Saint site and have a peruse of the relative chemical resistance table on the right, just below the one at the top

The 2375 is supposedly even better than this.

As I also say, I've actually spent a long time looking at the even more resistant options, like pure PFA, FEP and PTFE. But, whilst they're slightly better in terms of ultimate resistance, they're mechanical properties also make them a pain in the ass to work with in terms of connecting to bits of glass that needs taking apart, cleaning, setting back up in different patterns and so on; they're too rigid (so they'll tend to try and pull glass over or put strain on it), usually opaque (can't easily see if there's a condensation problem developing) and don't stretch or flex very much (meaning they can't be forced over gas ports, won't seal very well and will tend to slip off). They're good for things like HPLC, but that's because the equipment is in a pretty much permanent setup, weighs a ton and has specialized fittings to grip the tubing.

All this makes the ultra-chemically resistant tygon the best option I can think of in terms of laboratory glass. It's not as truely inert as the true fluoropolymers, but it's very close. And it's far more suited to draping and push fit connections. Added bonus, it's smooth as glass inside (making it very easy to clean, important) and it's crystal clear (so it's easy to see what's going on inside).

It's also why I have that pricey PTFE thickened vacuum grease. When a single experiment like this is using up pages of people debating every detail and waiting for one or two people to report back, I can do without the tubing / grease variable, thank you thank you.

I'm having a chat with the girlies at Saint about 2375. I'll report back on how it performs. If any of you are interested and thinking of doing some corrosive / reactive gas work, shoot me a PM just to say (yep). If I get enough PMs showing some interest, I may buy a big spool in a size and grade I know will work and chop it to order up for yars. The typical spool lengths are far too long and far too expensive for amateurs on a budget; you just need to think about what lengths you're going to use prior to firing up the scissors and binning it after one go.

"CARRR, that jumper's almost shiny, I'll take that as well!"

[Edited on 13-8-2010 by peach]

blogfast25 - 13-8-2010 at 07:16

'Black stuff': I'm referring to your second photo of a conical flask on page 3, with a dark liquid at the bottom, erm... 'John'.

blogfast25 - 13-8-2010 at 12:19

My best bet regarding the whitish material you claim is anh. AlCl3, is that it is in fact a mixture of Al hydroxy chlorides (Al(OH)nCl3-n) and possibly some plain old hydrated alumina. I think water has gotten into your system, causing the reaction by the usual means: 3 H3O+ + Al + n H2O ---> Al(H2O)n 3+ + 3/2 H2 + 3 H2O. Any AlCl3 formed would be partially or wholly hydrolysed.

I'd bet your product would not sublime and would not work as a Friedel Crafts catalyst.

DJF90 - 13-8-2010 at 14:55

Well I'd bet that in order for you to be correct, Blogfast, that there would have to be a shit-load of water kicking about. We're talking saturated DCM, moist aluminium powder and wet HCl gas stream. Now I don't know about you, but peach doesn't appear to be that sloppy of an experimenter. Just what's the problem with him actually making anhydrous AlCl3? Its hardly impossible, and in my personal opinion, I believe he's done just that. Until he runs it again sans neoprene, and has obtained some concrete results with analysis, lets give the lad a break - he's obviously making good progress and I'd hate for the negative comments from someone like yourself to put him off trying it again, or even worse, prevent him from posting the results here when he does finally get it working ok.

peach - 13-8-2010 at 15:27

'Black stuff': I'm referring to your second photo of a conical flask on page 3, with a dark liquid at the bottom, erm... 'John'.

erm... 'blogfast', as I've now explained, at least twice, that 'black stuff' is from commercial grade, supplier bought AlCl3. Not a product of this reaction. Indeed, I stated this (I believe) on the first page, with similar photos, whilst I was explaining why I was trying to run this method.

Are you actually reading prior to hitting reply?

I've supplied written evidence, photos and now videos, and you're still repeating the same questions.

I'm perfectly willing to be wrong, and for you to be right, but I've explained the dark solvent now a few times. How many more times do I need to repeat this? What's this? The third time?

[Edited on 13-8-2010 by peach]

entropy51 - 13-8-2010 at 15:42

I've supplied written evidence, photos and now videos, and you're still repeating the same questions.

I'm perfectly willing to be wrong, and for you to be right, but I've explained the dark solvent now a few times. How many more times do I need to repeat this? What's this? The third time?

Where is the "evidence" that you've made any AlCl3 whatsoever? At least you got the thread title correct: "Preparation of anhydrous AlCl3 in DCM - photos". There are photos and lots of chit chat, but nothing to show that your mystery gunk is AlCl3. Instead of personally attacking me for "complaining", why don't you cough up your alleged "evidence" that the emperor is indeed clothed.

peach - 13-8-2010 at 23:50

There is indeed a lot of chit chat, however, I'm sure you'll agree a lot of it is either coming from pure speculation (as opposed to me having actually tried it, written it down, taken photos and made a video; albeit, without a test of the result video as of yet).

We have people who aren't sure about the solubility of it in DCM telling me it probably hasn't worked because the product would go into the solvent. Fairly basic stuff to be making comments on the reaction.

I've made it quite clear I'll be using ultra-chemically resistant tygon in the future, but have been told to be careful which tygon I choose. I'd expect you to make some comment about this being pointless, then demand I use it if the result functions. So I'll go with the 'pointless' chemical inertness in anticipation of any complaints should the result function.

I've posted evidence all over this thread showing that aluminium does react, in some way, with what appears to be aluminium, DCM and HCl(g) alone in an all glass setup. That is somewhat speculative evidence, agreed. Throwing a pure white crystalline result into some water and watching for fuming is less so, but I'm sure someone could still find room to argue about that result as well.

The 'evidence' I'm referring to in response to blogfast is the commercial grade material being impure and not part of this reaction; the photos (which I've now posted twice, with explanations of what it is and what I'm doing to it).

Quote:

I'm perfectly willing to be wrong, and for you to be right, but I've explained the dark solvent now a few times. How many more times do I need to repeat this? What's this? The third time?

And, in effect, you are still complaining. I know a pure white, material that violently fumes hasn't been shown yet, I've only mentioned the odd fuming when washing the glass out. A number of other people have already commented on that not being a solid result yet, which I've acknowledged and replied to with the following;

Quote:

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.

It's fairly obvious I am actually reading the replies (as evidenced by the massive one on page 3, replying to all the questions about things like solubility, with first hand experience of these properties, and why someone might want to do this).

In essence, you're moaning about repetitive observations that are already being taken into account, again. And, again, this is something you'll have a go at other people for when they 'state the obvious' in threads you're taking part in.

Look, I get the point. It's been made a lot of times already, I haven't made any solid claim as to having produced it since people asked for an all glass style run, I've actively told others what they have may not be anhydrous and I'm spending money and time getting the bits together so the other people in this thread will know for sure if it does or doesn't work; for free. I'm not a millionaire, it will take a little while to get a result with a very low degree of potential error.

[Edited on 14-8-2010 by peach]

blogfast25 - 14-8-2010 at 04:07

John,

You don't seem to understand. The difference between anhydrous AlCl3, a very useful product, and slightly moist goo (a mixture of Al(OH)n Cl3-n.m H2O and Al(OH)3.k H2O) is HUGE.

Try and repeat the experiment with just one (1!) drop of water added to the DCM: reaction will be slow, yet decisive. The water will act like the catalyst you need to chlorinate Al with HCl, except it will also hydrolyse the product.

Your experiment took place over quite a long duration, even with quite dry reagents you're probably accumulating water.

The reaction between Al and HCl at RT in the presence of an inert 'solvent' is simply something that according to the known Laws of the universe isn't possible.

The onus is on you to provide unequivocal proof of the miracle: dry AlCl3. Until further notice some of us will keep betting on those safe Laws of the universe...

peach - 14-8-2010 at 08:03

You're explaining the properties of Lewis acids to someone who has spent years working with them, more often they than B/Ls.

I know the difference between hydrate and anhydrous.

One drop of water will not hydrolyze a batch of Lewis acid. One drop equates to 0.05ml, which is roughly 0.05g, which is roughly 0.0028 moles of H2O. I'm dealing with tens to a hundred grams of hard Lewis acid on a regular basis. That's a 26 molar excess, minimum.

Maybe I've misread what you're saying, but the water can not catalytically hydrolyze Lewis acid in that fashion, for the simple reason that there's nowhere else for the hydrogen to come from to produce the resulting hydrogen chloride fumes, but additional water.

Where in this seemingly catalytic hydrolysis you propose is the HCl(g) getting it's hydrogen from other than the water? How does more than an order of magnitude less water hydrolyze that quantity of acid? Why do the workups involving Lewis acids hydrolyze their reaction with at least an equimolar amount of water to acid? Why does my flask cease fuming when left, yet fume in a linear fashion as I add water?

As more water contaminates the sample, it degrades the Lewis acid acids proportionally and, perhaps more importantly, yields B/L acids that do then cause the acid to behave in different fashions. E.g. Lewis acids on their own (as I understand it from a post by Sauron or Nicodem) have a hard time demethylating toluene, and require a B/L present to function. Indicating a partially hydrolysed Lewis acid will not only skew the yields, but also allow for side reactions.

The onus is not solely on me to provide evidence for everyone else, until I start claiming it as fact. I'm simply providing the evidence as I get it myself, making judgments on that and asking reasonable questions about what it could be if it's not yet the acid.

And now, for the billionth time, yes... I will be re-running it, I have not only hinted at that, but described in detail (which seems to be surpassing a few of you) how I'm going to do it. And I will be posting results whether it's successful or fails.

I was purposefully not replying to this thread to avoid speculation that goes beyond what's available. But, some of the replies are so blatantly incorrect they don't require any further testing (e.g. AlCl3 dissolves in DCM, so the results seen so far can't be AlCl3). I don't mind people making mistakes, or want to pick away at them and call them idiots, but that is simply wrong and needs stating very clearly as such; and I've posted the photos of why that is so, twice, with explanations, which you yourself seem to have overlooked (e.g. "solvent is dark brown", suggesting you thought that was a product of the reaction).

Your own reply, regarding one drop of water ruining the entirety of the Lewis acid sites, is equally incorrect. If it were not, hard Lewis acids would come in sure-seal bottles. They don't. Also, see the above questions, which are rhetorical in my opinion.

I'm not trying to argue for the sake of it, or prove I'm 100% correct based on some agreeably shaky results in an unfinished flask. I do 100% agree that you're correct, this does need rerunning and needs to produce solid acid that fumes at the end. But what I am demonstrating is that the aluminium is reacting with the gas in a system that is significantly cleaner and drier than anyone else in this thread has so far attempted.

This reply is not so much speculation over the reaction it's self, as what I know to be correct for Lewis acids. Which needs stating so others can see why my reasoning does or doesn't function.

To throw some speculative questions into the pot, I also routinely form anhydrous AlI3 at low temperatures using I2, the same aluminium source, solvents and same methods (or worse) and yield a result that not only fumes on contact with the atmosphere, but also functions in the subsequent reactions, and shows it's doing so on TLC plates. I2 alone will react with aluminium if left to sit for a few minutes at room temperature.

Let this be the last time I have to say this, should anyone else say it needs repeating, it will be, and you will hear the results! If the thread is quiet for a week or three, that's me waiting for the post to arrive, compiling videos and uploading things.

[Edited on 14-8-2010 by peach]

peach - 20-8-2010 at 10:00

Okay! Finally!

I decided to try running it with the 500ml wash bottle, since the others still haven't shown up.

It's all 2375 and krytoxed.

I piranha'd the glass, rinsed multiple times, then baked the entire lot dry in the oven (the aluminium was already in there as well, under nitrogen). Redistilled the DCM from sodium sulphate and ran the result directly into the still capped wash bottle. Flushed it all through with dry nitrogen, then generated the HCl from sulphuric on anhydrous, supplier bottled CaCl2. The exit on the wash bottle was going through 98% sulphuric, and continually flowing in the exit direction.

There were lots of observations I could talk about;

Strong effervescence, observations on the pH of the exhaust gas with regards to addition rates and progress, that the exhaust was ignitable, that samples of the material in there appear to strongly interact (violently so) with water

Suffice it to say, I'm not happy with this attempt and will give it another go once I have restocked on some other bits.

I suspect, water has made it into that somewhere and produced something ominously like a gel; albeit one that rapidly settles.

I haven't opened the flask yet to have a test of what's in there, but I suspect this ones a gonner, having put molar excesses of HCl through it and still not having a pure white material.

I had some teething problems running this. For a start, something solid clogged the wash head, twice. Meaning I had to open it and quickly poke the solid back out (allowing for possible water absorption, despite me doing it in a few seconds). And it was a true solid pellet, not a gel (violent poking only dislodged it, as opposed to breaking it up).

More of an annoyance, I was using a brand new addition funnel and, later into the night, I could hear a high pitch whine, like a mosquito. I assumed one of the tapers had opened and spent ten or fifteen minutes with my ear against the glass, pressing on fittings to find the problem.

To my surprise, I then discovered an absolutely minute pinhole in one of the seams of the funnel. By that, I mean I can't actually see it now with me looking for it. But it's there, because I could see HCl escaping and moisture collecting around the point, and hear it. The HCl was still leaving this hole, but it's another point of entry and variable.

I now have a replacement funnel in the post.

What I find curious is the end colour of the solvent. It certainly isn't green anymore, it's orangey brown. I have checked the aluminium, it doesn't appear to be contaminated with iron. Is that the gas generators sulphuric passing into the wash bottle?

I noted the reappearance of that odd smell, despite the change in tubing and use of an all glass setup. Sulphorous material being generated by sulphuric contamination?

I have a stockpile of dead batteries to hand, so may give Cl2 a go after this since someone asked. Failing that, I have clean Iodine and HgCl2 and could have a go at priming the aluminium. Neither of these are ideal given the extra complications they involve, but would still be good in terms of generating a fine, clean result (as opposed to the CP mucky nodules).

I will have to assume the solvent is still too damp and the stream needs further work. I've not had any gel issues preparing AlI3 with these containers of solvent and poorer methodology. The only thing changing has been the halogen source. Indicating the gas stream may be at fault, if water has been the problem here.

I could dry the solvent with Calcium hydride, but my supplier doesn't have any. Or molecular sieves.

My current guess is, the sulphuric is damp and needs redistilling. Next time, I will condense it off directly into the addition funnel, then crimp off the tubing (need some of those crimpy things surgeons use really).

With regards to watson's point, again, I haven't checked the solvent yet, but some solvent (about 5-10ml) did make an attempt to leave the bottle and puddled in the exit. I was feeling quite tired after being sat there and developed a headache, so was left wondering if perhaps there may have been (even if they were only traces), some of his chloroform present.

[Edited on 20-8-2010 by peach]

blogfast25 - 20-8-2010 at 11:20

CaCl2 is not a brilliant desiccant. I wouldn't trust it for your purposes. Water being the one thing that will lead to misleading results...

watson.fawkes - 20-8-2010 at 11:30

[...] generated the HCl from sulphuric on anhydrous, supplier bottled CaCl2. The exit on the wash bottle was going through 98% sulphuric, and continually flowing in the exit direction.
[...]
What I find curious is the end colour of the solvent. It certainly isn't green anymore, it's orangey brown. I have checked the aluminium, it doesn't appear to be contaminated with iron. Is that the gas generators sulphuric passing into the wash bottle?

Armarego 5th ed. on purification of hydrogen chloride, p 430:

Quote:

Passed through conc H2SO4, then over activated charcoal and silica gel.

My guess is that your HCl is wet. The generator you're using doesn't have adequate H2SO4 in it at all times to dry the gas. I don't know the exact cut-over point, but at some point the residual H2O in the H2SO4 of your generator is going to start preferentially attaching itself to the HCl rather than that the H2SO4. Per Armarego, concentrated H2SO4 will happily absorb residual water vapor, but that's as part of gas washing train. Silica gel will adsorb any remaining. I'm not sure how much you need silica gel for a test run, though.

The other possibility is adsorbed water on the aluminum. You mentioned that you baked it, but it's possible that it wasn't long enough. With finely divided powder, you've got to get all the water vapor out through gas diffusion, so it will bake out faster on a flat sheet than as a compact volume inside a flask.

peach - 20-8-2010 at 11:49

Thanks for the quick replies watson and blogfast.

The CaCl2 was for the HCl generation, I was using sulphuric to keep water out at the exhaust side.

I used the supplier bottle of anhydrous CaCl2 over table salt to get more gas per unit volume of salt and avoid introducing moisture by using damp kitchen salt.

Calcium hydride seems to be the preferred method of drying DCM. I'm not sure the DCM being damp is the issue however. I suspect it's that sulphuric. And the end colour and smell of the solvent suggests the sulphuric may have wandered into the wash bottle.

Edit: Thanks for those quotes, I'll look into trying it as per the suggestion. I too suspect the HCl is damp. £140 for a 220g lecture bottle of HCl, plus £500 for the monel regulator, plus a university VAT code. Generators are the only option at present. I have found it a lot cheaper and easier to get, but I also don't need 5kg of it and still don't have the regulator. I suspected activated carbon may be necessary and have a colour indicating silica tube, but wouldn't I be better replacing the silica with CaCl2?

Double edit: I was just reading this page, which is already purple on my search results and yet slipped by my memory for the ease of using the sulphuric to hand, but I could also order some more HCl next time from the supplier. He suggests the gas stream is entirely dry when dripping HCl onto CaCl2. Is that going to be true? There's a lot of water floating around in there to be potentially picked up by the exiting stream.

[Edited on 20-8-2010 by peach]

entropy51 - 20-8-2010 at 12:38

Double edit: I was just reading this page, which is already purple on my search results and yet slipped by my memory for the ease of using the sulphuric to hand, but I could also order some more HCl next time from the supplier. He suggests the gas stream is entirely dry when dripping HCl onto CaCl2. Is that going to be true? There's a lot of water floating around in there to be potentially picked up by the exiting stream.

The HCl(g) prepared by dripping Conc. HCl onto CaCl2 should be dry enough. I have used that method to prepare AlCl3 the dry way. The paper referenced in your link has been discussed and posted on this forum.

Your description of your method above sounded as if it should be adequate for removing all but minute traces of H2O. I wonder if there isn't a problem other than traces of H2O hiding somewhere? I think the purple color of your reaction may be a clue, but I don't know what it means.

blogfast25 - 20-8-2010 at 12:43

Are you referring to the dry and HOT method here, entropy51?

entropy51 - 20-8-2010 at 12:57

Yes, the hot dry way. Unfortunately, it may be the only way.

blogfast25 - 20-8-2010 at 13:00

Yes, the hot dry way. Unfortunately, it may be the only way.

Methinks so to. But he seems to be having an honest shot at it.

But HCl from conc. H2SO4 and CaCl2 already dry enough? Perhaps, if you're using a lot of H2SO4, as acid and as desiccant at once, which is what John seems to intend.

peach - 20-8-2010 at 13:01

Your description of your method above sounded as if it should be adequate for removing all but minute traces of H2O. I wonder if there isn't a problem other than traces of H2O hiding somewhere? I think the purple color of your reaction may be a clue, but I don't know what it means.

Thanks, although I suspect blogfast may be on to a problem with that method of generating it.

Dried the solvent as best as I can and redistilled prior to use, porting straight into the wash bottle.

Checked the aluminium for contamination, seems very clean.

Generated AlI3 using this and poorer work no problems.

Looks like it's the generator, as per watson's suggestion.

Not sure what you mean about the purple result, it's orange / brown. There's a piece of white paper and black duct tape that I was using as a colour reference stuck to the back. I believe the purple you're seeing is the black tape. I still don't know why it'd be such a pronounced orange, given what's in there.

This needs redoing, with watson's modifications in place. I doubt that hydrochloric onto CaCl2 method is going to be any drier. There is far less water present in the H2SO4 generator. If the gas is picking moisture up from that, it'll pick it up from the hydrochloric even better.

[Edited on 20-8-2010 by peach]

blogfast25 - 20-8-2010 at 13:06

John, is the choice of CaCl2 as anion source for a particular reason? Semi-calcined NaCl might be drier?

[Edited on 20-8-2010 by blogfast25]

watson.fawkes - 20-8-2010 at 13:10