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Author: Subject: Preparation of Anhydrous Aluminum Chloride from Cupric Chloride and Aluminum
gdflp
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[*] posted on 17-12-2015 at 18:18


AlCl<sub>3</sub> will condense acetone and leave a mixture of red polymers. Strong acids and bases are not compatible with acetone.



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blogfast25
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[*] posted on 17-12-2015 at 18:50


I think the OP should attempt reaction of neat anh. CuCl2 with Al powder (or finely shredded foil?) at test tube level, to get an idea of how exothermic the reaction is.

It might even be possible, should the reaction prove too fast/exothermic, to include heat sinks like CaF2 to the mix, to 'tame' things a little.

This page (scroll down a little) does give an idea of what to expect:

http://www.amazingrust.com/Experiments/background_knowledge/...

[Edited on 18-12-2015 by blogfast25]




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[*] posted on 17-12-2015 at 21:12


Quote: Originally posted by gdflp  
AlCl<sub>3</sub> will condense acetone and leave a mixture of red polymers. Strong acids and bases are not compatible with acetone.


I am not surprised to read that... I was discussing the idea of using acetone as a solvent for AlCl3 with a professional chemist earlier, and he thought it would work well... I was skeptical... I think I might still give it a go.

Ethyl acetate might work better.

[Edited on 18-12-2015 by JJay]
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[*] posted on 17-12-2015 at 21:16


Quote: Originally posted by blogfast25  
I think the OP should attempt reaction of neat anh. CuCl2 with Al powder (or finely shredded foil?) at test tube level, to get an idea of how exothermic the reaction is.

It might even be possible, should the reaction prove too fast/exothermic, to include heat sinks like CaF2 to the mix, to 'tame' things a little.

This page (scroll down a little) does give an idea of what to expect:

http://www.amazingrust.com/Experiments/background_knowledge/...

[Edited on 18-12-2015 by blogfast25]


I have some shredded aluminum foil set aside for that purpose, but first I am going to try the reaction without a test tube. It should be easy enough to introduce CaCl2 to the mix.
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[*] posted on 17-12-2015 at 21:57


Quote: Originally posted by JJay  
I can put it into some terms that you will understand. Let's say that Sam, Jill and Jen are ions. They don't like hanging out at the bookstore much, especially when their annoying relatives are there. They are more likely to hang out at the bookstore when members of the opposite sex are present. But they don't hang out at the bookstore when it is packed.


I have no idea what you are trying to babble about, but ions prefer a solution with other ions in it.




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Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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[*] posted on 18-12-2015 at 00:20


Quote: Originally posted by DraconicAcid  
Quote: Originally posted by JJay  
I can put it into some terms that you will understand. Let's say that Sam, Jill and Jen are ions. They don't like hanging out at the bookstore much, especially when their annoying relatives are there. They are more likely to hang out at the bookstore when members of the opposite sex are present. But they don't hang out at the bookstore when it is packed.


I have no idea what you are trying to babble about, but ions prefer a solution with other ions in it.


To a point. After that point, they prefer a solution with fewer ions. You might want to try a hobby that isn't as technically demanding like paper mache or perhaps rock collecting.
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[*] posted on 18-12-2015 at 05:55


Quote: Originally posted by JJay  
You might want to try a hobby that isn't as technically demanding like paper mache or perhaps rock collecting.


DA is a degreed chemist, as am I, for what that's worth.

One the one hand we have Debye Huckel, a sophisticated piece of now mainstream chemical theory, elegantly and rationally explaining that and why ionic strength increases solubility.

On the other hand we have a psycho-babble explanation by JJay.

Rara, who to believe?




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[*] posted on 18-12-2015 at 11:48


Quote: Originally posted by blogfast25  
Quote: Originally posted by gdflp  
Ferric chloride might be another option to consider as well, with an HoF of -400kJ/mol(Wolfram Alpha). It shouldn't be as exothermic as the reaction with cupric chloride, but it's more widely available(at least in the US) than zinc chloride as it's sold by some pottery suppliers.


The fact that FeCl3 sublimes at 315 C could be a recipe for contaminated AlCl3.


But could you distill the crude product, or do AlCl3 and FeCl3 form an azeotrope?
CRC Handbook: BP AlCl3 = 180 °C (sublimes); BP FeCl3 ≈316 °C (wiki: decomposes).
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[*] posted on 18-12-2015 at 13:58


I mixed together 0.5 grams (slight molar excess) of aluminum foil that had been shredded in a coffee grinder and 3.2 grams of anhydrous CuCl2 and set it on fire with a propane torch. It burned with green and orange sparkly flames and gave off considerable smoke, likely a mixture of aluminum chloride and hydrogen chloride formed by its contact with water vapor. The reaction was pretty lively but not so vigorous that I would be afraid to do it in a test tube with good ventilation (although with really finely powdered aluminum, I am not so sure). The ashes were a mixture of white aluminum chloride that reacted violently with water, a tiny bit of unreacted aluminum, a trace of light green copper compounds (possibly containing the impurities), and nuggets of copper.

Next I am going to try dissolving the CuCl2 in acetone and dropping it on aluminum to see what happens.
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[*] posted on 18-12-2015 at 15:18


CuCl2 in acetone doesn't appear to react very readily with aluminum, at least not at room temperature without any special effort to initiate the reaction.
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[*] posted on 18-12-2015 at 15:29


Quote: Originally posted by JJay  
CuCl2 in acetone doesn't appear to react very readily with aluminum, at least not at room temperature without any special effort to initiate the reaction.


Does your CuCl2 actually dissolve in (Me)2CO?

Try adding a spec of iodine to get the party started?

[Edited on 18-12-2015 by blogfast25]




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[*] posted on 18-12-2015 at 15:39


Quote: Originally posted by blogfast25  
Quote: Originally posted by JJay  
CuCl2 in acetone doesn't appear to react very readily with aluminum, at least not at room temperature without any special effort to initiate the reaction.


Does your CuCl2 actually dissolve in (Me)2CO?

Try adding a spec of iodine to get the party started?

[Edited on 18-12-2015 by blogfast25]


I don't have any iodine, but I think that would probably work... a drop of water might work also. I'd rather not break out the mercury for this reaction.
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[*] posted on 18-12-2015 at 16:22


Remember that the SRPs for these reactions are all calculated/measured for water-solvated species, not solutes in acetone.

[Edited on 19-12-2015 by blogfast25]




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[*] posted on 7-1-2016 at 16:31


Hi, i know this doesnt help produce AlCl3 from copper chloride but about a month ago i reacted NH4Cl with Al metal powder (one or other was in excess but i dont really remember and i didnt write proportions in my lab book) and upon heating with a torch the mixture began to react vigorously but not violently and produced a voluminous white cloud that stank of HCl, no ammonia smell was noticed but i was trying to keep my face out of the foul cloud. there may have been a little water in my NH4Cl but it was freshly opened that day, so i believe it was a cloud sublimed/recondensed AlCl3.
If you try this i would recomend using NH4Cl in excess (which i think is what i did) because this should prevent run-away reaction by keeping the reaction out of stoich' and the sublimation of of NH4Cl should also cool the reaction, but will give more contaminants. if it is done in stoich' proportions it may get out of control.
anyhoo just a thought

...also... sorry if this has been tried, i must admit that i hadnt really read the thread that thoroughly, but ive read a few times on here about people trying to make anhydrous AlCl3 and i thought this might be helpful

[Edited on 8-1-2016 by moominjuice]
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[*] posted on 7-1-2016 at 16:42


@moominjuice:

Assuming the vapours could be condensed, it's still a recipe for AlCl3 with strong contamination with NH4Cl, as it too sublimes quite easily.




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[*] posted on 7-1-2016 at 17:07


Well the reaction didnt flare up or go mad... in fact it looked cool enough that a glass vessel with sand in the bottom would easily survive. condensation would therefore be fairly simple, also AlCl3 and NH4Cl have quite different temperatures at which the sublime (thus separable), and for most uses i dont think NH4Cl would be a problematic containment (unless measuring how much you have). the reason i like this approach is that it is simple, requires no solvent (but solvent could purify it) and makes more sense than any other method on here ive read about just because it works quickly and uses easily obtained reagents... ive tried making aluminium halides via other methods (eg I2 in toluene with Al) and they dont work that well.
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[*] posted on 7-1-2016 at 17:18


This is one of three chemistry projects I have sitting on a shelf right now. I'm currently working on putting together a lab with appropriate ventilation before I pursue this further.

All this might be rendered moot by Schlessinger's preparation of aluminum chloride from Inorganic Laboratory Preparations. It looks easy and doesn't require any reagents that I can't buy at the grocery store, nor does it require any special apparatus that everyone doesn't have already.

Untitled.png - 152kB

[Edited on 8-1-2016 by JJay]
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[*] posted on 7-1-2016 at 17:28


Ok point taken, ill see if i can do the reaction with excess Al, thus removing the contaminant, and see if that will occur easily enough as to not require special equipment. for the record i would be surprised if the first reaction i tried got hot enough to sublime the NH4Cl (after i removed the torch) but ill play around some time soon and find out for sure.
i need a project at the moment anyway.
perhaps you have to have been there to appreciate it but it was fun to make such a big cloud of an apparently hard to make substance, with so little effort.

(your edit makes my comment seem quite random lol, never mind, ill do some experiments and get back to you)

[Edited on 8-1-2016 by moominjuice]
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[*] posted on 7-1-2016 at 17:39


Yeah, if you set off a half gram of aluminum with a stoichiometric amount of copper (II) chloride, you'll definitely want to do it in extreme ventilation.
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[*] posted on 7-1-2016 at 17:44


Quote: Originally posted by moominjuice  


(your edit makes my comment seem quite random lol, never mind, ill do some experiments and get back to you)



LoL, sorry about that.
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[*] posted on 7-1-2016 at 18:29


Quote: Originally posted by moominjuice  
Ok point taken, ill see if i can do the reaction with excess Al, thus removing the contaminant, and see if that will occur easily enough as to not require special equipment.


Using excess Al will not necesarily remove the contaminant (NH4Cl). That's largely a question of what temperature the reaction front reaches: if it gets close or exceeds the sublimation temperature of NH4Cl it will inevitable start coming over with the AlCl3. Re-sublimation may get rid of it though.

Your reasoning would be valid if the reaction was carried out in a closed reactor but it's not.

I do think it's very worthwhile trying it again. I might do so myself.

Testing the product for NH4Cl is easy: react some of it with water, then add strong alkali like NaOH. The smell of NH3(g) is unmistakable and has quite a low olfactory detection limit.

My ZnCl2/Al experiment:

http://oxfordchemserve.com/lab-preparation-of-alcl3-reductio...

Equipment: a large test tube and a primitive condenser! ;)

[Edited on 8-1-2016 by blogfast25]




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[*] posted on 7-1-2016 at 19:09


And something else just occurred to me about that NH4Cl/Al reaction: that should yield NH4! Which would then split as:

2 NH4(g) === > 2 NH3(g) + H2(g)

Hmmm...:o

[Edited on 8-1-2016 by blogfast25]




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[*] posted on 7-1-2016 at 19:34


Thank you for your input, i enjoy discussing chemistry and i agree with your points (but they are points that i had already considered... im sorry if that sounds narcissistic). i am new to this forum and am trying to restrain myself from derailing this thread by running off at tangents, so i have been attempting to keep my non-copper based input to a limited level.
but when i say i would like to change the reaction to that of an excess of Al what i mean is that i wish to go away and think about how best to do this while driving the reaction to completion. i had not intended to imply that simply adding more Al to the reaction would achieve this, not least of all is that if i did i would have to heat the reaction to such an extent as to damage glassware, and of course apply pressure.
this reaction is exothermic, but not necessarily in a brutal sense. if excess Al were added above the main reaction in a layer it might help create a secondary reaction that is primarily endothermic, reacting with the hot ammonium chloride, condensing the less hot, while allowing a purer product to condense further up... this would complicate the proceedure greatly but it was the first thought off the top of my head.
like i say, i see and understand the issues you have raised and will think and experiment and see if i cant come back with an improvement.
i had intended on just posting an idea, and your criticisms have got me thinking about chemistry ... which has cheered me up. I really need that :) im glad i joined this forum

(edit)
yes reacting metal with ammonium salt can yield h2 but im not sure why thats such a problem.
many years ago, as i often do, i got bored and reacted Mg with NH4Cl just because i thought it might make a flame that contained no carbon and wasnt toxic. to my suprise it burned like a thermite with a normal looking flame above it with no smell of nitrogen oxides. it was really cool, a flame that makes no CO2 and no toxic gas or solid (there was no visible ammonium chloride "smoke")
i fail to see that producing NH3 or H2 is a problem.

(edit number 2)
i figured that 2NH3 H2 mix would burn nicely if you didnt get my point

(edit 3 put "nh2" not "nh3" oops, also u got ur stoich wrong :P)
[Edited on 8-1-2016 by moominjuice]

[Edited on 8-1-2016 by moominjuice]

[Edited on 8-1-2016 by moominjuice]
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[*] posted on 8-1-2016 at 07:09


Comparison of Reaction Heats:

3/2 ZnCl2(s) + Al(s) ===> AlCl3(s) + 3/2 Zn(s)

Enthalpy of Reaction (est.) at STP = - 53 kJ/mol

3 NH4Cl(s) + Al(s) === > AlCl3(s) + 3 NH3(g) + 3/2 H2(g)

Enthalpy of Reaction (est.) at STP = + 130 kJ/mol

(all based on NIST Webbook and Wolfram Alpha data)

So this would suggest that the reaction with salmiac is in fact endothermic and driven by Entropy (ammonia and hydrogen gases)


[Edited on 8-1-2016 by blogfast25]




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[*] posted on 8-1-2016 at 14:39


Quote:
'peach', many moons ago, tried to prepare AlCl3 from Al powder in DCM, gassed with dry HCl. No explosion (but no product either...)


If we're talking about the same person I'm pretty sure he eventually succeeded... but the reaction he wanted to use it for turned out to be impossible. Small world, it is. I no longer have his notes, unfortunately.

But as long as we're bubbling HCl through powdered Al suspended in a solvent... I suggest benzophenone. Polar enough to dissolve HCl, inert enough to contain AlCl3, high-boiling enough to withstand the reaction energy, and OTC via Ca benzoate. Also probably a good solvent for the reaction involving CuCl2.

Quote:
And something else just occurred to me about that NH4Cl/Al reaction: that should yield NH4! Which would then split as:


You of all people should know that AlCl3 is a Lewis acid, and NH3 is a Lewis base. What does that mean?

[Edited on 8-1-2016 by clearly_not_atara]

[Edited on 8-1-2016 by clearly_not_atara]
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