Sciencemadness Discussion Board

Preparation of Anhydrous Aluminum Chloride from Cupric Chloride and Aluminum

JJay - 16-12-2015 at 07:34

I've been trying to find information on preparing aluminum chloride from copper (II) chloride and aluminum. A couple of reputable sources have stated that it is possible, but I haven't found many details.

Preparing anhydrous copper (II) chloride is easy enough... it can be made from hydrogen peroxide, hydrochloric acid, and copper metal - or it can be made from calcium chloride and copper sulfate. Either preparation is easy. I couldn't find any 30% hydrogen peroxide locally, so I used the calcium chloride / copper sulfate method:

--

100.00 grams (0.77519 mol) of calcium chloride monohydrate were added to 200 mL of water and stirred until it dissolved. The solution was allowed to cool to room temperature, and the specific gravity was measured to ensure that the salt had not been contaminated with additional water. 193.55 g of copper sulfate pentahydrate was dissolved in 700 mL of water, producing a dark blue solution. The calcium chloride solution was added 10 mL at a time with stirring. Addition produced a green color that quickly dissipated, along with considerable precipitate. The mixture was allowed to stand overnight and then vacuum filtered. The clay-like filter cake was washed with 100 mL of water and discarded. The cyan filtrate was found to be unreactive to additional calcium chloride solution.

The volume of the filtrate was reduced over heat to 400 mL. A small amount of white precipitate formed. The mixture was vacuum filtered, and the presumed gypsum filter cake was discarded. Reducing the volume to 300 mL did not produce a significant amount of additional precipitate. The solution was evaporated to dryness in an oven with occasional stirring to produce 85.46 g of an amorphous brown solid, a yield of 82.00%, which was considered rather low due to mechanical losses.

--

So anyway, now I'm looking at around 85 grams of anhydrous copper chloride, and I have a roll of aluminum foil. Anyone have any ideas as to how to best prepare aluminum chloride from these materials, or should I just start torching assorted stoichiometric compositions in various apparatus and see what works?




20151216_061650.jpg - 785kB

Amos - 16-12-2015 at 08:13

This is something I was thinking about trying at some point, as well. My idea is to stir anhydrous CuCl2 and aluminium powder in chloroform, after which, theoretically, copper metal or cuprous chloride would settle out and aluminium chloride would be in solution in the chloroform. I have no idea if it would work, though.

blogfast25 - 16-12-2015 at 08:22

Well done on the anh. CuCl2.

The reaction between Al powder and anh. CuCl2 has actually been reported on this forum somewhere (but I can't find it).

I can't vouch for the veracity of the results but as I recall the reaction proceeds but is far too exothermic to be of great use for preparing anh. AlCl3.

Only one way to find out: try! You will need Al powder though, no matter how course. Foil won't work, I think.

What has been tried and tested (by me) is the preparation of anh. AlCl3 by reaction of Al powder and anh. ZnCl2:

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

This runs nicely contained and at leisurely pace.

It too has been reported here, with one member replicating my results but again I can't find the thread (the search facility is truly crap).

JJay - 16-12-2015 at 08:23

Quote: Originally posted by Amos  
This is something I was thinking about trying at some point, as well. My idea is to stir anhydrous CuCl2 and aluminium powder in chloroform, after which, theoretically, copper metal or cuprous chloride would settle out and aluminium chloride would be in solution in the chloroform. I have no idea if it would work, though.


It looks like it would take extremely large volumes of chloroform for that to work... I only have about 120 mL, so I don't think I will be able to do that... but it might be possible to use anhydrous ethanol.

blogfast25 - 16-12-2015 at 08:24

Quote: Originally posted by Amos  
This is something I was thinking about trying at some point, as well. My idea is to stir anhydrous CuCl2 and aluminium powder in chloroform, after which, theoretically, copper metal or cuprous chloride would settle out and aluminium chloride would be in solution in the chloroform. I have no idea if it would work, though.


Worth a shot but I think the reaction is too exothermic, while requiring some temperature to get going.

JJay - 16-12-2015 at 08:27

Quote: Originally posted by blogfast25  
Well done on the anh. CuCl2.

The reaction between Al powder and anh. CuCl2 has actually been reported on this forum somewhere (but I can't find it).

I can't vouch for the veracity of the results but as I recall the reaction proceeds but is far too exothermic to be of great use for preparing anh. AlCl3.

Only one way to find out: try! You will need Al powder though, no matter how course. Foil won't work, I think.

What has been tried and tested (by me) is the preparation of anh. AlCl3 by reaction of Al powder and anh. ZnCl2:

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

This runs nicely contained and at leisurely pace.

It too has been reported here, with one member replicating my results but again I can't find the thread (the search facility is truly crap).


I've seen your work, and I do like the preparation from zinc chloride, but it is just as hard to get zinc chloride here as it is aluminum chloride (I can get it, but i'd have to pay completely outlandish hazardous material fees).

I have also read that it is possible to prepare zinc chloride from zinc and copper chloride, but I haven't seen the details on that either... might be interesting to look into at some point in the future.

Praxichys - 16-12-2015 at 08:27

Be careful. Heating mixtures of aluminum and highly halogenated hydrocarbons have been known to cause violent decomposition. Both carbon tetrachloride and tetrachloroethylene form detonable mixtures with aluminum powder. It has also been documented that aluminum forms an explosive mixture with chloroform, especially in the presence of AlCl3.

http://webwiser.nlm.nih.gov/getSubstanceData.do?substanceId=...

I don't think it would be too dangerous to try, but definitely make sure the container can vent safely if there is a sudden exothermic reaction. The product of this reaction is the AlCl3 you're after, so maybe intentionally starting and controlling this reaction is what you need.

blogfast25 - 16-12-2015 at 08:39

Quote: Originally posted by Praxichys  
Be careful. Heating mixtures of aluminum and highly halogenated hydrocarbons have been known to cause violent decomposition. Both carbon tetrachloride and tetrachloroethylene form detonable mixtures with aluminum powder. It has also been documented that aluminum forms an explosive mixture with chloroform, especially in the presence of AlCl3.

http://webwiser.nlm.nih.gov/getSubstanceData.do?substanceId=...

I don't think it would be too dangerous to try, but definitely make sure the container can vent safely if there is a sudden exothermic reaction. The product of this reaction is the AlCl3 you're after, so maybe intentionally starting and controlling this reaction is what you need.


Interesting.

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

blogfast25 - 16-12-2015 at 08:46

The Standard Heat of Formation of CuCl2 is - 206 kJ/mol (NIST), of ZnCl2 it's - 415 kJ/mol (Wolfram Alpha).

That difference of 200 kJ/mol clearly indicates that the reaction with CuCl2 runs MUCH hotter than with ZnCl2.

gdflp - 16-12-2015 at 08:58

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.

blogfast25 - 16-12-2015 at 09:18

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.

Amos - 16-12-2015 at 09:18

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.


Rakugoldpottery on eBay sells anhydrous ZnCl2 as well as FeCl3, both quite cheaply.

JJay - 16-12-2015 at 09:43

It looks like perhaps the easiest way to react CuCl2 with aluminum while capturing the product in an anhydrous state is to dissolve the copper (II) chloride in dry ethanol and then drip it onto the aluminum. I think aluminum chloride should go into solution with the copper precipitating out. Once the reaction is complete, the reaction mixture could be filtered and the aluminum chloride precipitated out of the ethanol with chloroform. Right?

[Edited on 16-12-2015 by JJay]

blogfast25 - 16-12-2015 at 10:53

Quote: Originally posted by JJay  
It looks like perhaps the easiest way to react CuCl2 with aluminum while capturing the product in an anhydrous state is to dissolve the copper (II) chloride in dry ethanol and then drip it onto the aluminum. I think aluminum chloride should go into solution with the copper precipitating out. Once the reaction is complete, the reaction mixture could be filtered and the aluminum chloride precipitated out of the ethanol with chloroform. Right?

[Edited on 16-12-2015 by JJay]


I'm wondering, just wondering, if AlCl3 with ethanol will not perhaps give Al ethoxide? Or at least partly? Hmmm...

MolecularWorld - 16-12-2015 at 11:41

@OP: Good job!
We all know what the product is, and that it's purity is acceptable for the intended use, but my experience requires me to ask the following pedantic questions:
How do you know the CuCl2 is anhydrous? What tests did you perform to prove this? Did you test for decomposition after heating? You didn't even state the temperature or the length of time spent in the "oven".
You mention sourcing reagents locally. I interpret that to mean your reagents are OTC. Everybody knows that all OTC products are highly contaminated, enough to throw off the results of most experiments (I didn't think so, but I was wrong). You must repeat the experiment with reagent grade chemicals, and/or verify the product spectroscopically, to ensure contaminants didn't give a misleading result.
What was the purity of the product? How did you determine this?
How do you know your product is anhydrous copper(II) chloride? Don't tell me you judged this based mostly on color, that would be "beneath contempt."
Edit: Also, those one-piece plastic mason jar lids are far from air-tight (I also use them for chemical storage). You'd be better off using a regular gasketed lid, otherwise your "anhydrous" product soon won't be (unless you store the jar in a desiccator).
Quote: Originally posted by blogfast25  
Well done on the anh. CuCl2.
"Talk about an aversion to evidence..."

[Edited on 16-12-2015 by MolecularWorld]

blogfast25 - 16-12-2015 at 11:58

Quote: Originally posted by MolecularWorld  
Everybody knows that all OTC products are highly contaminated, enough to throw off the results of most experiments (I didn't think so, but I was wrong).


Two broad, sweeping statements and as such inaccurate. There are plenty reasonably pure OTC chemicals and plenty of reactions that are quite robustly resistant to contaminants.

Your attempt at creating some equivalence here is pathetic, as anyone who has ever prepared anh. CuCl2 will know.

Asking questions isn't pedantic, doing so out of spite though is silly.

[Edited on 16-12-2015 by blogfast25]

MolecularWorld - 16-12-2015 at 12:06

Quote: Originally posted by blogfast25  
There are plenty reasonably pure OTC chemicals and plenty of reactions that are quite robustly resistant to contaminants.

You see, that's what I thought, and to some extant continue to believe; i'll be sure to quote this post the next time someone tries to convince me otherwise.

blogfast25 - 16-12-2015 at 12:09

Quote: Originally posted by MolecularWorld  

You see, that's what I thought, and to some extant continue to believe; i'll be sure to quote this post the next time someone tries to convince me otherwise.


The only way to deal with purity of chemicals, OTC or otherwise, is on a case-by-case basis and in the correct context.

[Edited on 16-12-2015 by blogfast25]

MolecularWorld - 16-12-2015 at 12:12

Quote: Originally posted by blogfast25  
Your attempt at creating some equivalence here is pathetic, as anyone who has ever prepared anh. CuCl2 will know.
Asking questions isn't pedantic, doing so out of spite though is silly.


My questions are genuine, though pedantic, and not purely "out of spite".
But I would never have asked them without my previous experience, which was silly indeed.

[Edited on 16-12-2015 by MolecularWorld]

Detonationology - 16-12-2015 at 12:19

Quote: Originally posted by blogfast25  
The only way to deal with purity of chemicals, OTC or otherwise, is on a case-by-case basis and in the correct context.

MSDS sheets are terribly unreliable when it comes to OTC chemicals. The listed range of a chemical concentration in a product always seems to fall on the lower end. In the case of household ammonia labeled on the MSDS as 3-10% NH4OH, titration found that the concentration was between only 3-3.5% concentration. It's false marketing! If the MSDS says 99-100%, I could only assume that that would mean that the only other "1%" of ingredients would be unintentional contamination, and relatively negligible, but it's probably more like 5% contamination, which would be considered quite significant.

[Edited on 12-16-2015 by Detonationology]

blogfast25 - 16-12-2015 at 12:25

Quote: Originally posted by Detonationology  

MSDS sheets are terribly unreliable when it comes to OTC chemicals. The listed range of a chemical concentration in a product always seems to fall on the lower end. In the case of household ammonia labeled on the MSDS as 3-10% NH4OH, titration found that the concentration was between only 3-3.5% concentration. It's false marketing! If the MSDS says 99-100%, I could only assume that that would mean that the only other "1%" of ingredients would be unintentional contamination, and relatively negligible, but it's probably more like 5% contamination.



MSDS sheets are NOT CoAs or QA statements. They have nothing to do with marketing either.

And last time I looked, 3-3.5 lies within 3-10.

[Edited on 16-12-2015 by blogfast25]

Detonationology - 16-12-2015 at 12:31

Quote: Originally posted by blogfast25  
And last time I looked, 3-3.5 lies within 3-10.

If your chemical supplier sold you a chemical that varied in concentration by 7%, would you be happy with your purchase? Would you deem it reliable enough to use in a reaction where purity is paramount? It's absolutely about marketing: the manufacturer make more money when a higher concentration is labeled.

UC235 - 16-12-2015 at 13:33

Household ammonia solution is not a reagent and was never meant to be used as one. MSDS is intended to be a statement of hazards about a product for use by the general public in the event of an emergency. Manufacturers are not required to disclose all ingredients or exact quantities as doing so would make duplicating their product trivially easy. They are required to list hazardous components and a range of concentrations for which safety measures are going to be identical. I'm not sure if the numbers they can put for ranges are dictated by law somewhere, but for normal end consumers, it is irrelevant. If the ammonia cleans their oven fine, then they don't care if it's 3% or 10%.

The ammonia bottle doesn't advertise 10% and then contain 3%. It doesn't advertise any percentage. If it did, the story would be different. If my 8.25% bleach is 6% when I buy it I am not happy. But probably it's going to be 8-9% with the 8.25% being a lowballed estimate to account for degradation over time.

[Edited on 16-12-2015 by UC235]

DraconicAcid - 16-12-2015 at 13:37

Quote: Originally posted by Detonationology  
Quote: Originally posted by blogfast25  
And last time I looked, 3-3.5 lies within 3-10.

If your chemical supplier sold you a chemical that varied in concentration by 7%, would you be happy with your purchase? Would you deem it reliable enough to use in a reaction where purity is paramount? It's absolutely about marketing: the manufacturer make more money when a higher concentration is labeled.

No, they just don't want to have to print up different MSDSs for their ammonia if they happen to sell 3% and 8% seperately.

Besides- what the MSDS says isn't important for marketing, since that's provided as a safety issue, not a sales pitch. Does it give a concentration on the bottle?

[Edited on 16-12-2015 by DraconicAcid]

blogfast25 - 16-12-2015 at 13:38

Quote: Originally posted by Detonationology  

If your chemical supplier sold you a chemical that varied in concentration by 7%, would you be happy with your purchase? Would you deem it reliable enough to use in a reaction where purity is paramount? It's absolutely about marketing: the manufacturer make more money when a higher concentration is labeled.


If it was thus labelled, I'd have no choice or seek out another supplier.

You're conflating purity and concentration. A solution could contain almost only, say, B and solvent and yet B's concentration could be low and variable. Worksheets can be adjusted for concentration of reagents, usually very easily.

Concentrations of certain chemicals (volatile ones e.g.) may vary inevitably, depending on product. You won't find any HCl 37.50 w% any time soon.

The marketer is only dishonest if he sells a product that falls outside of the specification as advertised.

Detonationology - 16-12-2015 at 14:25

Quote: Originally posted by blogfast25  
You're conflating purity and concentration. A solution could contain almost only, say, B and solvent and yet B's concentration could be low and variable. Worksheets can be adjusted for concentration of reagents, usually very easily.

How could the concentration of a solution be "low and variable" if the solution contains "almost only" a salt (I would assume) dissolved in a solution? I do not follow this "B" analogy. Could you please explain with a relative example how purity and concentration are not related?

[Edited on 12-16-2015 by Detonationology]

blogfast25 - 16-12-2015 at 15:30

Quote: Originally posted by Detonationology  
I do not follow this "B" analogy. Could you please explain with a relative example how purity and concentration are not related?



I didn't say purity and concentration are not related, I said you're conflating them, i.e. they're not the same.

Take an ammonia solution. Highly pure it would evaporate to literally nothing. But if some solids are left on evaporation that would point to some impurity (obviously not water, and not ammonia either!). Those impurities could be expressed with respect to the ammonia (e.g. '99.9 % ammonia'), while also expressing the concentration of ammonia with respect to the solvent (e.g. '25 w% NH3' or '2M NH3').

MolecularWorld - 16-12-2015 at 18:54

Quote: Originally posted by blogfast25  
Take an ammonia solution. Highly pure it would evaporate to literally nothing. But if some solids are left on evaporation that would point to some impurity (obviously not water, and not ammonia either!).

This is misleading. A highly pure ammonia solution left to evaporate in air may leave a residue from the reaction of ammonia and atmospheric carbon dioxide. If you meant "...in a stream of dry air scrubbed of carbon dioxide" or "...in a desiccator with [lots of] sodium hydroxide" you should have said so.

JJay - 16-12-2015 at 20:43

Quote: Originally posted by MolecularWorld  
@OP: Good job!
We all know what the product is, and that it's purity is acceptable for the intended use, but my experience requires me to ask the following pedantic questions:
How do you know the CuCl2 is anhydrous? What tests did you perform to prove this? Did you test for decomposition after heating? You didn't even state the temperature or the length of time spent in the "oven".
You mention sourcing reagents locally. I interpret that to mean your reagents are OTC. Everybody knows that all OTC products are highly contaminated, enough to throw off the results of most experiments (I didn't think so, but I was wrong). You must repeat the experiment with reagent grade chemicals, and/or verify the product spectroscopically, to ensure contaminants didn't give a misleading result.
What was the purity of the product? How did you determine this?
How do you know your product is anhydrous copper(II) chloride? Don't tell me you judged this based mostly on color, that would be "beneath contempt."
Edit: Also, those one-piece plastic mason jar lids are far from air-tight (I also use them for chemical storage). You'd be better off using a regular gasketed lid, otherwise your "anhydrous" product soon won't be (unless you store the jar in a desiccator).
Quote: Originally posted by blogfast25  
Well done on the anh. CuCl2.
"Talk about an aversion to evidence..."

[Edited on 16-12-2015 by MolecularWorld]


These are fair questions. I have actually not tested to see if the CuCl2 is anhydrous, but it would be possible to do so by measuring the specific gravity of a solution of known concentration. I baked it for several hours, at around 150 C until it was dry (some green crystals formed) and then at 300 C with stirring every half hour until everything turned brown and the clumps stopped sticking together, indicating that no more water was being emitted. The stirring is important to prevent the formation of a fused mass sticking to the crystallization dish.

The storage container could probably be sealed a little tighter with a gasket or PTFE tape, but I don't think CuCl2 is very hygroscopic, so I haven't worried about it. (And come to think of it, I have stored hygroscopic compounds with those lids for months with no issue.) Your point is well taken, though - I would use a better container for storing aluminum chloride. I think the cupric chloride purity is actually pretty good, but it would be possible to recrystalize the CuCl2 as the dihydrate salt from water and then dehydrate it... or it may be feasible to recrystalize the anhydrous form from an alcohol.

The reagents are OTC. The calcium chloride is not food grade but appears to be reasonably free of impurities and consistently behaves as expected in other experiments. This is the first time I have used this copper sulfate, but it also performed exactly as expected; I have little reason to believe that it contains a lot of impurities.

I am not sure how to best verify the CuCl2 spectroscopically with equipment available to the average hobbyist, but I like that idea. I wonder if ACS publishes standards for CuCl2....

I don't know what the purity is... I don't think it is analytical grade, but I do suspect it is quite good.


[Edited on 17-12-2015 by JJay]

JJay - 16-12-2015 at 22:30

Say, CuCl2 is a little more hygroscopic than I thought... it turns blue-green as it absorbs water. Pretty interesting stuff... mine is showing no signs of changing color in its current container.



[Edited on 17-12-2015 by JJay]

blogfast25 - 17-12-2015 at 08:43

Quote: Originally posted by MolecularWorld  
Quote: Originally posted by blogfast25  
Take an ammonia solution. Highly pure it would evaporate to literally nothing. But if some solids are left on evaporation that would point to some impurity (obviously not water, and not ammonia either!).

This is misleading. A highly pure ammonia solution left to evaporate in air may leave a residue from the reaction of ammonia and atmospheric carbon dioxide. If you meant "...in a stream of dry air scrubbed of carbon dioxide" or "...in a desiccator with [lots of] sodium hydroxide" you should have said so.


To illustrate a simple principle it isn't necessary to provide a complete methodology, MW. And evaporate could mean 'boil off' too...

Context is everything. I was explaining something by simple example, not putting up a method of analysis.


[Edited on 17-12-2015 by blogfast25]

blogfast25 - 17-12-2015 at 08:50

Quote: Originally posted by JJay  


I don't know what the purity is... I don't think it is analytical grade, but I do suspect it is quite good.




Contaminants in your starting reagents will partly have been carried over.

One contamination you can bet on is CaSO4 because it has limited solubility, so your CuCl2 solution will have contained some. BaCl2 would have been better in that respect, because BaSO4 is far more insoluble.

You could use BaCl2 or Ba(NO3)2 to test your product for sulphates.

CaSO4 solubility at 20 C, about 21 g/L (water, Wikipedia).

[Edited on 17-12-2015 by blogfast25]

JJay - 17-12-2015 at 13:12

I am going to try dissolving the CuCl2 in acetone and reacting it with aluminum in an ice bath. I suspect that CaSO4 contamination is quite low due to its low solubility and the tendency for it to get salted out by more soluble compounds. I don't remember its solubility in water offhand, but 21 g/L is about 10x higher than I had thought.

DraconicAcid - 17-12-2015 at 13:16

Quote: Originally posted by JJay  
I suspect that CaSO4 contamination is quite low due to its low solubility and the tendency for it to get salted out by more soluble compounds.


You can't salt out an ionic compound from solution, unless the salt you're using is a calcium salt or a sulphate (common ion effect). If you increase the ionic strength of the solution without a common ion, you'll just make it more soluble.

JJay - 17-12-2015 at 13:21

Quote: Originally posted by DraconicAcid  
Quote: Originally posted by JJay  
I suspect that CaSO4 contamination is quite low due to its low solubility and the tendency for it to get salted out by more soluble compounds.


You can't salt out an ionic compound from solution, unless the salt you're using is a calcium salt or a sulphate (common ion effect). If you increase the ionic strength of the solution without a common ion, you'll just make it more soluble.


That is simply not the case

DraconicAcid - 17-12-2015 at 13:23

Quote: Originally posted by JJay  

That is simply not the case

Oh? Explain your reasoning.

mayko - 17-12-2015 at 13:41

Wikipedia lists the water solubility of CaSO4 as:

0.21g/100ml at 20 °C (anhydrous)[1]
0.24 g/100ml at 20 °C (dihydrate)[2]

By my math, that's ~2g per liter.

AJKOER - 17-12-2015 at 13:55

Per Atomistry on AlCl3 ( http://aluminium.atomistry.com/aluminium_trichloride.html ):

"Aluminium trichloride, AlCl3, was originally made by heating an intimate mixture of alumina and carbon to redness in a stream of chlorine (Oersted's method). It may be more readily prepared by heating aluminium in a wide glass tube in a rapid current of dry hydrogen chloride, or in a stream of chlorine. If it is required to prepare the chloride from the oxide, a neater method than Oersted's is to heat the oxide in a current of chlorine and sulphur chloride: -

4Al2O3 + 3S2Cl2 + 9Cl2 = 8AlCl3 + 6SO2.

Instead of chlorine and sulphur chloride, carbon tetrachloride vapour or carbonyl chloride may be used. A simple method of preparation is said to consist in heating crude alumina or clay to redness in a current of hydrogen chloride and carbon disulphide vapour, and purifying the aluminium chloride so obtained by sublimation over iron filings. "

Per Atomistry on COCl2 ( http://carbon.atomistry.com/carbon_oxychloride.html ):

"Carbonyl chloride may also be prepared by the oxidation of chloroform by chromic acid, when 20 parts of chloroform, 400 of sulphuric acid, and 50 of potassium dichromate are heated together on the water-bath:

2CHCl3 + K2Cr2O7 + 4H2SO4 = 2COCl2 + K2SO4 + Cr2(SO4)3 + 5H2O + Cl2"

blogfast25 - 17-12-2015 at 14:02

Quote: Originally posted by mayko  
Wikipedia lists the water solubility of CaSO4 as:

0.21g/100ml at 20 °C (anhydrous)[1]
0.24 g/100ml at 20 °C (dihydrate)[2]

By my math, that's ~2g per liter.


Ooopsie. Also by my math. :(

blogfast25 - 17-12-2015 at 14:05

Quote: Originally posted by JJay  


That is simply not the case


DA is completely correct, JJay.

But the contamination by CaSO4 will be small and won't interfere here.

JJay - 17-12-2015 at 15:53

Quote: Originally posted by blogfast25  
Quote: Originally posted by JJay  


That is simply not the case


DA is completely correct, JJay.

But the contamination by CaSO4 will be small and won't interfere here.


No, he is not. But the contamination by CaSO4 will be very small indeed.

DraconicAcid - 17-12-2015 at 16:01

Quote: Originally posted by JJay  
No, he is not.

Citation? Or reasoning?

If you increase the ionic strength of a solution, you will decrease the activity of the ions, so they basically act as if they have a lower concentration. This makes ionic compounds *more* soluble in solutions of high ionic strength.

It is known.

JJay - 17-12-2015 at 16:38

Let me put it to you this way: you can't look at ionic strength alone to determine solubility. You have to look at activity (which is sometimes expressed as a function of ionic strength) and concentration. Ordinarily, the effect of increased ionic strength decreases as concentration increases.


Also, copper (II) forms complexes with water.

blogfast25 - 17-12-2015 at 17:06

Quote: Originally posted by JJay  
Let me put it to you this way: you can't look at ionic strength alone to determine solubility. You have to look at activity (which is sometimes expressed as a function of ionic strength) and concentration. Ordinarily, the effect of increased ionic strength decreases as concentration increases.


Also, copper (II) forms complexes with water.


You sound confused, TBH. No one said 'alone' but increased ionic strength increases solubility (except in the case of common ion effects).

See e.g.:
Quote:
The theory of activity versus concentration is important in industrial, environmental, and biochemistry. The increase in solubility of an electrolyte in a solution of a second electrolyte with no common ions compared with pure water is not an easy concept to grasp because it seems to be counterintuitive. The simple experiment described here illustrates this principle visually and dramatically. Students attempt to dissolve CaSO4•2H2O (gypsum) in pure water and in 0.25 M NaCl. The gypsum dissolves almost completely in the sodium chloride solution, but not in pure water. Students then measure the calcium concentrations in filtered aliquots of both solutions to quantify the solubility difference they observed. Students calculate mean activity coefficients using their measured concentrations and also from the Davies Equation, an extension of Debye–Hückel theory. The basic principle is there are ionic interactions between the solute ions and the solvent ions, which allow for more dissolution because only free ions enter into the expression for the solubility product equilibrium constant. From a simple mathematical point of view, in higher ionic strength solutions, activity coefficients for calcium and sulfate become smaller, and hence the concentrations must be larger to maintain a constant solubility product at equilibrium.


Source.

Yes, Cu(+2) forms hexaaqua coordination complexes. What has that got to do with anything here, though? It also forms weak tetrachloro coordination complexes but even in pure CuCl2 solution not so much.

[Edited on 18-12-2015 by blogfast25]

JJay - 17-12-2015 at 17:20

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.

blogfast25 - 17-12-2015 at 17:25

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.


And this is your theory? Or based on what mainstream chemistry theory?

blogfast25 - 17-12-2015 at 17:30

Basic theory of activities, ionic strength and solubilities

Why believe your system would behave differently?

JJay - 17-12-2015 at 17:32

That is how those activity-concentration equilibrium equations work. It's not *my* theory.

blogfast25 - 17-12-2015 at 17:48

Quote: Originally posted by JJay  
That is how those activity-concentration equilibrium equations work. It's not *my* theory.


With ions that are attracted to the opposite sex? Whatever makes you happy, I guess...

JJay - 17-12-2015 at 17:53

Quote: Originally posted by blogfast25  
Quote: Originally posted by JJay  
It looks like perhaps the easiest way to react CuCl2 with aluminum while capturing the product in an anhydrous state is to dissolve the copper (II) chloride in dry ethanol and then drip it onto the aluminum. I think aluminum chloride should go into solution with the copper precipitating out. Once the reaction is complete, the reaction mixture could be filtered and the aluminum chloride precipitated out of the ethanol with chloroform. Right?

[Edited on 16-12-2015 by JJay]


I'm wondering, just wondering, if AlCl3 with ethanol will not perhaps give Al ethoxide? Or at least partly? Hmmm...


It does, or at least it gives a strong complex that behaves very similarly to Al ethoxide... so chloroform probably won't precipitate the AlCl3... It looks like acetone might work instead of alcohol... I suspect that AlCl3 will form complexes with acetone too, but likely they can be cleaved with a boiling water bath under vacuum (rather than in a stream of HCl gas at 300 C).

gdflp - 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.

blogfast25 - 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]

JJay - 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]

JJay - 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.

DraconicAcid - 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.

JJay - 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.

blogfast25 - 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?

Bezaleel - 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).

JJay - 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.

JJay - 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.

blogfast25 - 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]

JJay - 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.

blogfast25 - 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]

moominjuice - 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]

blogfast25 - 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.

moominjuice - 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.

JJay - 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]

moominjuice - 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]

JJay - 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.

JJay - 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.

blogfast25 - 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]

blogfast25 - 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]

moominjuice - 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]

blogfast25 - 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]

clearly_not_atara - 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]

blogfast25 - 8-1-2016 at 15:29

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


I. 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...

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



I. Why: because 'peach was peach' (real life: John)? I never believed that project would work and still don't believe it now. Not with benzophenone either.

Not at RT, not without a protonic oxidiser.

You'd probably be in trouble if it did:

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

... is worth about - 400 kJ/mol at STP, seriously exothermic!

His notes are still on the site, search and yee shall find! :)

II. And yet AlCl3 and NH3 don't seem to form a Lewis adduct in these here conditions, they're funny like that these pesky Lewis acids and bases ;). If you have references to the contrary, I'd be very interested.

[Edited on 8-1-2016 by blogfast25]

Magpie - 8-1-2016 at 19:13

I tried to make PCl3(l) by gassing white P with Cl2. With a standard enthalpy of formation of -320 KJ/mol the reaction was getting out of control and I decided not to try that again. It can be done however, as demonstrated by garage chemist at versuchschemie.de. This file may be hard to find due to the administrative problems at versuchschemie.de.

The standard enthalpy of formation for PCl5(s) as calculated from that of its gas and its heat of vaporization is ~ -375 + 65 = -310 KJ/mol. I did accomplish this with no heat problem whatsoever by dissolving the white P in chloroform then gassing it with Cl2.

blogfast25 - 8-1-2016 at 19:24

@Magpie:

As far as I know Al doesn't combine with Cl2 when starting at RT (unlike with Br2 or I2). The same holds true for Al + HCl: it needs activation energy.

That's why I never believed 'peach' would pull it off (Al + HCl in DCM at RT).

P is notorious for its reactivity in direct union reactions.

aga - 9-1-2016 at 11:16

Quote: Originally posted by blogfast25  
As far as I know Al doesn't combine with Cl2 when starting at RT

So .... if one were to be considering the direct route, the Al would need to be Hot then ... hmm ...

blogfast25 - 9-1-2016 at 12:45

Quote: Originally posted by aga  
Quote: Originally posted by blogfast25  
As far as I know Al doesn't combine with Cl2 when starting at RT

So .... if one were to be considering the direct route, the Al would need to be Hot then ... hmm ...


Yes. Heat gently till reaction starts, after initiation it should be self-sustaining...