Sciencemadness Discussion Board - How to make relatively pure AlCl3

How to make relatively pure AlCl3

Adas - 8-10-2011 at 00:37

Hello,
I want to make relatively pure AlCl3 as a strong lewis acid from industrial grade HCl and ordinary Al foil. I wanna know the yield (if it is high enough).

When I put excess Al foil into HCl, it creates AlCl3 in good yield, but there is a problem: Aluminium foil is nor pure aluminium and contains also hard metal particles (which are not soluble in HCl), and maybe some other metals. The insoluble particles can be filtered off, but I want to know how pure AlCl3 will remain in the solution. Is it 95%? Or 75%?
Any infos will help, thanks.

Endimion17 - 8-10-2011 at 04:53

You can't get anhydrous (pure) aluminium chloride from an aqueous solution. One visit to Wikipedia would've inform you on that. You'd get a mixture of chlorohydrates which would decompose at higher temperatures to oxide.

Direct synthesis from metal and chlorine or metal and hydrogen chloride at almost 700 °C is the only way.

Adas - 8-10-2011 at 05:11

So you say, that HCl(aq) and Al do not make AlCl3 but other products? I don't need anhydrous AlCl3, because I wanna use it in solution. It can be a hydrate, but I don't want other products.

Vogelzang - 8-10-2011 at 05:12

https://www.hyperlab.info/inv/index.php?s=309f6d356ac3dee658...

Some aluminum chloride patents.

Patent US1818839 Process for manufacturing anhydrous aluminum chloride
aluminum sulphate and MCl

Patent US1647446 Process for collecting and utilizing aluminum chloride

[Edited on 8-10-2011 by Vogelzang]

Nicodem - 8-10-2011 at 11:22

Quote: Originally posted by Adas

So you say, that HCl(aq) and Al do not make AlCl3 but other products? I don't need anhydrous AlCl3, because I wanna use it in solution. It can be a hydrate, but I don't want other products.

AlCl3 does not exist in aqueous solution. By dissolving aluminium in hydrochloric acid you get a solution of chloride anions and a complex mixture of various aluminium cations complexed with water molecules and chloride anions. The water obviously can not be removed by evaporation, as H2O is much more basic that Cl-. Even if you dissolve AlCl3 in water, you still can not obtain an aqueous solution of AlCl3, because it quantitatively reacts with water (AlCl3 is an extremely strong acid and it reacts with weak bases such as water).

Adas - 8-10-2011 at 11:24

So, is there any other strong and easy to get Lewis acid?

Nicodem - 8-10-2011 at 11:25

AlI3

turd - 8-10-2011 at 12:07

Hi Nicodem,

of course AlCl3 can not be obtained from aq. solution because HCl goes off before H2O, but I don't understand your explanation.

Quote: Originally posted by Nicodem

Yes, but couldn't you say the same thing about NaCl? The Na will be hydrated as well. Or do you make the difference because the Al-O bond is very strong? [AlOx(OH)y(H2O)z] vs. [Na(H2O)x] ?

Quote:

The water obviously can not be removed by evaporation, as H2O is much more basic that Cl-.

If it was only a matter of H2O and Cl-, the same argument should work for NaCl, no?

My intuitive explanation would be that Al is extremely oxophilic. That's what makes AlCl3 very acidic in aqueous solution. Seems to me a quite a different reason than the Lewis acidity that organikers use for Friedel-Crafts, etc. in non-aqueous solvents?

Retard-3000 - 8-10-2011 at 14:41

If you want a lewis acid you'll want the anhydrous form of aluminium trichloride, NOT the hydrate. My hypothesis it that it can be made by bubbling chlorine gas through aluminium tribromide in a suitable anhydrous solvent. Aluminium tribromide being easily made from aluminium and bromine at STP.

Endimion17 - 8-10-2011 at 14:58

The point of Friedel-Crafts is in anhydrous aluminium chloride which strips chlorine anion from an alkyl group. You need the whole AlCl3 as a species, because it turns into AlCl4- giving the radical (alkyl) which then enters the main reaction. Later that tetrachloroaluminate anion reacts with a hydrogen dumped from the molecule which was meing alkylated. and you get aluminium chloride and HCl. Aluminium chloride is a catalyst because it's being replenished.

This is elementary organic chemistry I was learning in high school.

Just pass some chlorine over heated aluminium wire. If you need it as a catalyst, you don't need much of it.

[Edited on 8-10-2011 by Endimion17]

peach - 8-10-2011 at 15:06

Quote: Originally posted by Adas

Hello,
I want to make relatively pure AlCl3 as a strong lewis acid from industrial grade HCl and ordinary Al foil. I wanna know the yield (if it is high enough).

When I put excess Al foil into HCl, it creates AlCl3 in good yield, but there is a problem: Aluminium foil is nor pure aluminium and contains also hard metal particles (which are not soluble in HCl), and maybe some other metals. The insoluble particles can be filtered off, but I want to know how pure AlCl3 will remain in the solution. Is it 95%? Or 75%?
Any infos will help, thanks.

Water is precisely the opposite to what you want anywhere near AlCl3, or any other Lewis acid.

In fact, you will struggle to make / keep it anhydrous at home without a glove box / manifold; very expensive gear (thousands) that takes a lot of practice (years).

The only way to produce it from an aqueous solution would be to dry it under a stream of dry hydrogen chloride. I suspect the aqueous product would start to sublime as it started going to anhydrous, adding yet another complexity; I can't find any information on the temperature needed to dehydrate the aqueous products.

Have a look at;

This thread

And this one

The only other two possibilities I know of are;

- The Royal Society of Chemistry mentioned dripping liquefied hydrogen chloride gas on aluminium

- Doing it in a solvent at room temperature using catalysts

Someone else also mentioned a distillation method for drying it to anhydrous. The precise details of which I can't remember, but it wasn't a standard distillation.

You can create and handle Lewis acids in the atmosphere. However, they will begin loosing their Lewis acid properties as they're exposed. Perhaps more importantly, they are fuming acid gases that draw water towards themselves and then dissolve to form (for the AlCl3 example, hydrochloric), so you will gain Bronsted Lowery acid properties that may drastically skew your results.

FeCl3 is a Lewis acid. It is also sold at just about any electronics store worth it's salts

; Maplin / RS / Farnell / RadioShack (used to be Tandy's in the UK until it disappeared) etc. It comes as it's hydrate; orange nodules. You could have a practice drying such things with that as a starting material. Note the interesting melting / boiling points for FeCl3. They're almost the same number.

[Edited on 9-10-2011 by peach]

Nicodem - 8-10-2011 at 23:42

Quote: Originally posted by turd

Hi Nicodem,

of course AlCl3 can not be obtained from aq. solution because HCl goes off before H2O, but I don't understand your explanation.

Quote: Originally posted by Nicodem

Quote:

The water obviously can not be removed by evaporation, as H2O is much more basic that Cl-.

Yes, of course, the same principle holds true for NaCl, but the example of NaCl decomposing during the dissolution is already described in nearly every textbook for students under the section about solvation. Being fanatically strict, it is not correct to talk about a solution of NaCl, as what defines sodium chloride (its "ionic bond" and crystal structure) is lost in its reaction with water (only the Na vs. Cl ratio is what is left).
There is however a very important difference. The dissolution of NaCl is reversible and after solvent removal you regenerate sodium chloride. The reaction of AlCl3 with water is not reversible. It is more similar to dissolutions of SOCl2 or CH3COCl in water, they dissolve and you get solutions, but these do not contain any more of the original solute after the equilibrium is reached. You also do not get the solutes back by solvent evaporation.

Adas - 9-10-2011 at 09:14

So, can I use FeCl3? Is it a good and stable Lewis acid?

Nicodem - 9-10-2011 at 09:44

Quote: Originally posted by Adas

So, can I use FeCl3? Is it a good and stable Lewis acid?

How is anybody going to answer when you forgot to say what you want to use it for?

Adas - 9-10-2011 at 10:51

Quote: Originally posted by Nicodem

Quote: Originally posted by Adas

So, can I use FeCl3? Is it a good and stable Lewis acid?

How is anybody going to answer when you forgot to say what you want to use it for?

I want to use it as a type of catalyst, but that's not really important. More important is, that I am gonna use it underwater, so it must be stable.

turd - 9-10-2011 at 14:24

Quote: Originally posted by peach

Water is precisely the opposite to what you want anywhere near AlCl3, or any other Lewis acid.

What?? HgCl2 and B(OH)3 - two classic Lewis acids - are perfectly compatible with water. The problem with AlCl3 is not its Lewis acidity, but that it hydrolyses like a MF, since Al2O3.xH2O is very stable.

Lewis Acid/Base:
A + :B --> A-B

Hydrolysis:
AX + H2O --> AOH + HX

Adas - 10-10-2011 at 09:02

Quote: Originally posted by turd

What?? HgCl2 and B(OH)3 - two classic Lewis acids - are perfectly compatible with water.

Maybe, but HgCl2 is toxic (I will NEVER mess with mercury!) and B(OH)3 is impossible to get for me. But I will try FeCl3 someday

Endimion17 - 10-10-2011 at 10:28

Adas, every pharmacy uses boric acid. It's not like you'd ask them for potassium...

Nicodem - 10-10-2011 at 11:15

Quote: Originally posted by Adas

I want to use it as a type of catalyst, but that's not really important. More important is, that I am gonna use it underwater, so it must be stable.

There is no need for being oxymoronic, as it is not going to help your cause.
It is pretty obvious you don't understand the most basic concepts, if you think that it does not matter what acid you use or that the nature of the base is irrelevant. There is a huge difference if you use a weak acid like boric acid or an supersoft acid like HgCl2, or if you use FeCl3, AlCl3, etc. In any case, a "Lewis acid" in water becomes just a normal "Brønsted acid", so you can just use H2SO4, or any such, to have the same effect. If you would have read the basic acid-base concepts, you would have realized by now that water is a base and that it reacts with acids, particularly with hard acids. When it does so with an (relatively) unhydrolysable acid like boric acid or BF3, it forms the corresponding "Brønsted acid" like HB(OH)4 and "HBF3(OH)". With the hydrolysable acids like FeCl3, AlCl3, ZrCl4, etc., it results in an acidic mixture of complexes and oxychlorides in dynamic equilibrium.

Quote: Originally posted by Adas

But I will try FeCl3 someday

FeCl3 hydrolyzes in water.

peach - 10-10-2011 at 12:00

Quote: Originally posted by turd

What??

You're right, over generalisation.

Okay, the types of Lewis acid he wants to use can't be in or around water.

unionised - 10-10-2011 at 12:22

Last time I looked boric acid wasn't a very good lewis acid and it's practically not a bronsted acid.
In any event, since water is quite a good lewis base ...

turd - 11-10-2011 at 06:18

Well yes, it isn't. I mentioned it, because usually it is the first acid students learn about that is not a Brønsted acid. And Hg ions are the archetypical soft acids.

Quote:

In any event, since water is quite a good lewis base ...

So? A colleague of mine spent his whole PhD combining Lewis acids and bases in water.

And IMHO he had more exciting results than the Schlenk technique people. Water is great - cheap, plentiful, stable and not even poisonous.

Panache - 15-10-2011 at 23:13

Quote: Originally posted by Adas

I want to use it as a type of catalyst, but that's not really important. More important is, that I am gonna use it underwater, so it must be stable.

Are you a fish? If not I suggest performing chemistry above water, like everyone else does. I saw two people get married once underwater, but they were diving instructors, are you a diving instructor?

You will have considerable difficulty using a Bunsen burner underwater, likewise for most electronics, also heating things is very difficult as you must not only heat your reaction but the entire volume of water you are under.

Hope this helps.

Prepuce - 16-10-2011 at 17:24

Might you me able to add aluminum filings and choloform
to a Parr bottle and shake while injecting chlorine rather
than hydrogen ? Open and rotovap to dryness.

Nicodem - 17-10-2011 at 07:58

Quote: Originally posted by Prepuce

Might you me able to add aluminum filings and choloform
to a Parr bottle and shake while injecting chlorine rather
than hydrogen ? Open and rotovap to dryness.

Have about a little more sense before suggesting potentially suicidal plans? What if aluminium starts reacting with chloroform and the enormous exotherm is rapidly released in a pressure vessel?
Have you done that? Do you have a reference? If not, please do not suggest self injury to others before you try it on your own body!

ScienceSquirrel - 17-10-2011 at 08:22

Aliphatic halogenated solvents and reactive metals are always a bad idea.
Chloroform explodes when mixed with quite a few eg aluminium,
magnesium, sodium, lithium, potassium, iron and zinc!

Vogelzang - 23-10-2011 at 06:20

Has anyone here tried aluminum in the Friedel-Crafts reaction? See attachment.

Attachment: JCE1989p0176Friedel-Crafts-Al.pdf (1.4MB)
This file has been downloaded 1220 times

Sedit - 23-10-2011 at 09:12

Over at The Vespiary some time back a member I believe it was V16 posted a means of extracting Anhydrous AlCl3 from an (aq) solution by using DMSO which would precipitate as a complex with the AlCl3 and would release under mild heating, anhydrous AlCl3. I once started to attempt the procedure however the complete disdain I have for the smell of heated DMSO prevented me from going any further.

Lambda-Eyde - 23-10-2011 at 09:25

Quote: Originally posted by Sedit

Over at The Vespiary some time back a member I believe it was V16 posted a means of extracting Anhydrous AlCl3 from an (aq) solution by using DMSO which would precipitate as a complex with the AlCl3 and would release under mild heating, anhydrous AlCl3. I once started to attempt the procedure however the complete disdain I have for the smell of heated DMSO prevented me from going any further.

For those interested in the procedure, see the attached reference.

Attachment: 3471250 PREPARATION OF ANHYDROUS INORGANIC METAL HALIDES.pdf (127kB)
This file has been downloaded 1863 times

smaerd - 24-10-2011 at 07:32

If no one else has confirmed that this works I will conduct tests replacing ethanol with methanol in the very near future.

[Edited on 24-10-2011 by smaerd]

Vogelzang - 24-10-2011 at 13:18

Here's an interesting AlCl3 synthesis. I wonder if it would work with tetrachloroethylene (dry cleaning fluid). It looks like it could be scaled up.

http://www.sciencemadness.org/scipics/AlCl3prep1.gif

http://www.sciencemadness.org/scipics/AlCl3prep2.gif

Vogelzang - 24-10-2011 at 13:49

Aluminum and bromine can react at room temperature. How about making and using aluminum tribromide in place of the trichloride?

http://en.wikipedia.org/wiki/Aluminium_bromide

http://pubs.acs.org/doi/abs/10.1021/j150498a017

Adas - 25-10-2011 at 04:18

Thanks, guys. But I think that dissociation is not a problem. So, what do you think? How pure is Al foil? 98% Al?

Sedit - 25-10-2011 at 07:56

Quote: Originally posted by Vogelzang

Here's an interesting AlCl3 synthesis. I wonder if it would work with tetrachloroethylene (dry cleaning fluid). It looks like it could be scaled up.

http://www.sciencemadness.org/scipics/AlCl3prep1.gif

http://www.sciencemadness.org/scipics/AlCl3prep2.gif

I ran simple test a while back on the feasibility of this reaction and IIRC someone either here or over at The vespiary again attempted this without very good results. I see no reason to use Tetrachloroethylene when the Dichloromethane they use is so available.

My results showed that when the DCM contained Cl2 it would be yellow... Placing Al into this solution did next to nothing and the yellow persisted overnight, however when a small amount of amalgumated Al was used the yellow color quickly faded.

See the problem with the solvent synthesis of AlCl3 seems to be that AlCl3 is needed as a catalyst for the reaction. This problem CAN be avoided using a small amount of Hg to get the ball rolling. I hope someone can make it work but I have yet to hear great results from the solvent methods of generating AlCl3.

If you fear Hg then perhaps Iodine would work to get things going.

SHADYCHASE54 - 1-11-2011 at 18:19

I was just wondering if refluxing aluminium chloride hydrate with a chlorinating agent like thionyl chloride would dehydrate it into an anhydrous state? If SOCl2 converts carboxylic acids to their respective chloride wouln't the same hold true in this case?

Adas - 2-11-2011 at 09:40

Quote: Originally posted by SHADYCHASE54

I was just wondering if refluxing aluminium chloride hydrate with a chlorinating agent like thionyl chloride would dehydrate it into an anhydrous state? If SOCl2 converts carboxylic acids to their respective chloride wouln't the same hold true in this case?

I think it's a good idea. Not impossible.

thanos thanatos - 3-11-2011 at 18:23

Quote: Originally posted by SHADYCHASE54

Yes, you can use SOCl2 to produce anhydrous AlCl3, but I suspect if someone needs to produce anhydrous AlCl3 because they cannot purchase it (and it is not a listed chemical), then I think in that case it is even more likely that that same person would not be able to buy thionyl chloride (or PCl3 or PCl5, for that matter), as it is a watched chemical and more expensive than anhydrous AlCl3.

Nicodem - 4-11-2011 at 11:25

Quote: Originally posted by Sedit

Sedit, the referred reaction is something totally different from what you describe. It is apparently an SN1 reaction between AlI3 and CH2Cl2. Its peculiarity is in that the aluminium iodide acts at the same time as the acid catalyst and as the source of the iodide while dichloromethane acts as a reaction substrate and the solvent. AlI3 forms by the oxidation of aluminium with iodine. (This type of nucleophilic substitution on dichloromethane was already mentioned in the 5th paragraph of this post.)

radiance88 - 18-3-2015 at 03:52

I recently just did this experiment because it's probably the easiest to get ahold of materials for and I'm a complete noob at all this. So I placed Al foil into HCl.. What remained after the considerably vigorous emission of gas (which i assume is Hydrogen), was just metallic looking precipitate, very similar in color to the aluminum foil.

I assume that this isn't AlCl3? If so... then what was the dark gray precipitate? If there was AlCl3, was it suspended in solution? What exactly am I looking at in the beaker then?

Texium - 18-3-2015 at 08:41

The gray stuff is excess aluminum, and possibly some impurities in the foil. The aluminum chloride is in solution as hydrated aluminum ions and chloride ions. If you're going for anhydrous AlCl₃ you're barking up the wrong tree, if you just wanted some hydrated AlCl₃, filter and crystallize. Don't boil it down, or it will decompose giving off HCl and leaving you with aluminum hydroxide/oxide.

Rhodanide - 5-5-2016 at 05:38

One great way is to heat Anh. Zinc Chloride with aluminum powder and sort of "Distilling" the subliming AlCl3 into a very cold (and closed) container. The YouTube channel "Chem Player" has an awesome video on this.
Cheers!
~T

gdflp - 5-5-2016 at 06:29

That method originated from SciMad http://www.sciencemadness.org/talk/viewthread.php?tid=30150#...

chemplayer... - 5-5-2016 at 07:32

It was someone on this site who posted the original blog posting we found. Can't remember which user though. We just kept playing with it until we had an acceptable yield.

Only caveat is that the product works great for some reactions (e.g. Friedel Crafts acylation it's superb), but for others we're getting substandard results (vanillin demethylation). It's quite possible there's some subliming ZnCl2 passed over in the product and this has a detrimental effect on some uses. The reason we suspect this is because we've controlled all variables in some reactions and they're still way below literature yield - the AlCl3 source is the only thing left to blame.

blogfast25 - 5-5-2016 at 07:56

Quote: Originally posted by chemplayer...

It was someone on this site who posted the original blog posting we found. Can't remember which user though. We just kept playing with it until we had an acceptable yield.

Only caveat is that the product works great for some reactions (e.g. Friedel Crafts acylation it's superb), but for others we're getting substandard results (vanillin demethylation). It's quite possible there's some subliming ZnCl2 passed over in the product and this has a detrimental effect on some uses. The reason we suspect this is because we've controlled all variables in some reactions and they're still way below literature yield - the AlCl3 source is the only thing left to blame.

That user was me. I first published it here, then on my blog (where you found it).

Re. your purity point, the only way to 'blame' the AlCl₃ conclusively is by comparing it to a commercial grade.

Re-sublimation may be the cure here. I've a feeling that most catalytic grade commercial AlCl₃ has been re-sublimated too.

[Edited on 5-5-2016 by blogfast25]

JJay - 5-5-2016 at 08:19

I made some aluminum chloride in a quartz tube by gassing aluminum with hydrogen chloride, but there was still a lot of unreacted aluminum left in it. I think that putting the aluminum into a boat and gassing with a larger quantity of HCl would have produced better results.