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Adas
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How to make relatively pure AlCl3
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.
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Endimion17
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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.
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Adas
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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.
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Vogelzang
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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]
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Nicodem
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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. |
AlCl<sub>3</sub> 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 H<sub>2</sub>O is much more basic that Cl<sup>-</sup>. Even if you dissolve AlCl<sub>3</sub> in
water, you still can not obtain an aqueous solution of AlCl<sub>3</sub>, because it quantitatively reacts with water
(AlCl<sub>3</sub> is an extremely strong acid and it reacts with weak bases such as water).
…there is a human touch of the cultist “believer” in every theorist that he must struggle against as being
unworthy of the scientist. Some of the greatest men of science have publicly repudiated a theory which earlier they hotly defended. In this lies their
scientific temper, not in the scientific defense of the theory. - Weston La Barre (Ghost Dance, 1972)
Read the The ScienceMadness Guidelines!
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Adas
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So, is there any other strong and easy to get Lewis acid?
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Nicodem
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AlI<sub>3</sub>
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turd
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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 | | AlCl<sub>3</sub> 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. |
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? [AlO<sub>x</sub>(OH)<sub>y</sub>(H2O)<sub>z</sub>] vs. [Na(H2O)<sub>x</sub>] ?
| Quote: | | The water obviously can not be removed by evaporation, as H<sub>2</sub>O is much more basic that Cl<sup>-</sup>.
|
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?
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Retard-3000
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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.
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Endimion17
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The point of Friedel-Crafts is in anhydrous aluminium chloride which strips chlorine anion from an alkyl group. You need the whole
AlCl<sub>3</sub> as a species, because it turns into AlCl<sub>4</sub><sup>-</sup> 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]
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peach
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| 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]
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Nicodem
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| 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 | | AlCl<sub>3</sub> 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. |
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? [AlO<sub>x</sub>(OH)<sub>y</sub>(H2O)<sub>z</sub>] vs. [Na(H2O)<sub>x</sub>] ?
| Quote: | | The water obviously can not be removed by evaporation, as H<sub>2</sub>O is much more basic that Cl<sup>-</sup>.
|
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? |
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.
…there is a human touch of the cultist “believer” in every theorist that he must struggle against as being
unworthy of the scientist. Some of the greatest men of science have publicly repudiated a theory which earlier they hotly defended. In this lies their
scientific temper, not in the scientific defense of the theory. - Weston La Barre (Ghost Dance, 1972)
Read the The ScienceMadness Guidelines!
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Adas
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So, can I use FeCl3? Is it a good and stable Lewis acid?
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Nicodem
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How is anybody going to answer when you forgot to say what you want to use it for? 
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Adas
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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.
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turd
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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
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Adas
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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 
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Endimion17
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Adas, every pharmacy uses boric acid. It's not like you'd ask them for potassium...
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Nicodem
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| 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 HgCl<sub>2</sub>, or if
you use FeCl<sub>3</sub>, AlCl<sub>3</sub>, etc. In any case, a "Lewis acid" in water becomes just a normal "Brønsted acid",
so you can just use H<sub>2</sub>SO<sub>4</sub>, 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)<sub>4</sub> and
"HBF<sub>3</sub>(OH)". With the hydrolysable acids like FeCl<sub>3</sub>, AlCl<sub>3</sub>,
ZrCl<sub>4</sub>, etc., it results in an acidic mixture of complexes and oxychlorides in dynamic equilibrium.
FeCl<sub>3</sub> hydrolyzes in water.
…there is a human touch of the cultist “believer” in every theorist that he must struggle against as being
unworthy of the scientist. Some of the greatest men of science have publicly repudiated a theory which earlier they hotly defended. In this lies their
scientific temper, not in the scientific defense of the theory. - Weston La Barre (Ghost Dance, 1972)
Read the The ScienceMadness Guidelines!
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peach
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You're right, over generalisation.
Okay, the types of Lewis acid he wants to use can't be in or around water.
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unionised
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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 ...
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turd
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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.
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Panache
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 | 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.
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Prepuce
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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.
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Nicodem
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| 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!
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