Sciencemadness Discussion Board - Apparatus for making anhydrous aluminium chloride

Apparatus for making anhydrous aluminium chloride

SuperOxide - 31-12-2021 at 21:33

I'm planning on making some anhydrous AlCl3 by reacting anhydrous ZnCl3 with Al powder. I know Chemplayer did a video on the synthesis which gives me a little hit as to what to expect more so than reading some articles about it. The main thing I noticed in their video was that they seemed to lose a lot of the product out from the gap between the elbow bend and GL45 neck, which I'm looking to avoid if possible.
It looks like Chemplayer had this setup for two reasons: 1) is to have the product end up in the reagent bottle, instead of having to move it from a flask to the bottle at the end, and; 2) to limit the exposure to moisture in the atmosphere.

I was thinking a setup/procedure like this would be a bit better.
1) Hook the elbow up to a 2 neck receiving flask, which would have an extra course frit inside of a vacuum port, which would stop the AlCl3 from leaving (I think EC would do that decently well, but I could be wrong. Maybe a finer grit would suffice)
2) Pull a vacuum (maybe even pull one during the entire procedure?)
3) After the reaction is complete, use very dry diethyl ether to dissolve the AlCl3 from the receiving flask, pour it into the storage bottle, then let the ether evaporate off under vacuum (or hooked up to a drying trap and let the ether vapour come through with gentile heating).

Here's a photo of the type of setup I'm thinking about :

Does anyone see something wrong with that type of setup?

Ubya - 1-1-2022 at 00:56

AlCl3 fumes would be very very fine, so i'm not really sure how much that grid would help, if it is too coarse you'll contaminate your vacuum pump, if it is too fine it will probably get clogged quite easily.
if it gets clogged, being under a vacuum, i don't think the pressure would increase much in the apparatus, at 151°C the vapour pressure of aluminium chloride is 100 torr (http://chemister.ru/Database/properties-en.php?dbid=1&id...), since you are cooling the receiving flask it will take a pretty good runaway reaction to go over one bar of pressure.

depending on how right i am you could just put everything under vacuum, seal the apparatus and do the reaction without having to worry about grits and loosing product.

SuperOxide - 1-1-2022 at 08:12

Quote: Originally posted by Ubya

depending on how right i am you could just put everything under vacuum, seal the apparatus and do the reaction without having to worry about grits and loosing product.

That's actually also what I was thinking. I could pull a vacuum while heating it until I see product coming over, then continue heating it but use a stopcock to close off the vacuum. Obviously I wouldn't seal one of the joints (probably the one going to the vacuum), just in case a positive pressure forms.

I think you might be right about it being too fine for the frit, and im not sure using a finer frit would do much other than clog, but I suppose I can experiment with it to find out.

Whats your take on using diethyl ether to get it out? Instead of scraping at it and pouring it into a separate container as a powder. I feel like I would never get everything.

[Edited on 1-1-2022 by SuperOxide]

Ubya - 1-1-2022 at 08:37

Quote: Originally posted by SuperOxide

Whats your take on using diethyl ether to get it out? Instead of scraping at it and pouring it into a separate container as a powder. I feel like I would never get everything.

[Edited on 1-1-2022 by SuperOxide]

that sounds a good idea, i can only think of 2 issues, i can't find a solubility value for aluminium chloride in diethyl ether so it may be great or you may need a few hundred milliliters of ether to disolve a few grams of the salt.
the other issue i can think of is that the ether really must be dry, like extremely dry, otherwise when you are going to boil off the ether your aluminium chloride will be damp already.

SuperOxide - 1-1-2022 at 09:02

Quote: Originally posted by Ubya

Yeah, I haven't been able to find an exact amount either, but luckily I have some on hand I can use. I also have 500g of ACS grade benzene if that's better. And I have some chloroform, or ethanol I dried out. In either case, I will do the first one on a smaller scale to give it a shot.

Quote: Originally posted by Ubya

the other issue i can think of is that the ether really must be dry, like extremely dry, otherwise when you are going to boil off the ether your aluminium chloride will be damp already.

I agree. I keep most of my solvents over some 3A sieves, but I suppose I could use some P2O5 to be 100% sure it's dry before moving forward (Im not sure if ether + phosphorus pentoxide reacts, if it does then I can use something else such as sodium metal).

[Edited on 1-1-2022 by SuperOxide]

Triflic Acid - 2-1-2022 at 07:55

You should probably use some acidic dehydrating agent since any suspended NaOH or Na2O from the sodium drying will destroy your product. But P2O5 should be fine

AJKOER - 2-1-2022 at 08:25

My advice, keep it simple.

OK, for small amounts, a vessel with plastic lid is filled with Chlorine gas containing a small piece of say pure Al foil. The amount of the Al foil converted to moles should be, at least, 1/3 of the volume of the vessel in liters divided by say 22.5.

Place the vessel in a microwave for a few seconds treatment. The Al foil ignites and burns in the atmosphere of chlorine. Cool the vessel to collect the AlCl3 using ethanol in which AlCl3 is soluble.

Yes, a small amount of AlCl3, but as the equipment goes, just an appropriate vessel and a microwave. I have used a large thick glass vessel with a plastic cap to avoid any cracking or burning if the Al strip comes into contact with the vessel where the Al creates, in the microwave, a very hot plasma electric arc. I have just suspended the Al strip like a light bulb filament from the cap of the vessel which has a hole from being punctured with a nail. As a result, no cracking of the vessel (which has occurred with me using even a porcelain vessel).

Impurities from water vapor or oxygen presence, not much of an issue as the latter products are not likely soluble in the ethanol, but some lose in yield.

Simple and quite exciting display with the arc and subsequent burning. Small yield however, but can repeat the process to concentrate your ethanol collecting solution, however, make sure all alcohol fumes have been removed before reloading Cl2 and ignition (otherwise, explosion hazard and chloro-organic presence).

For safety, I would recommend safety equipment and also performing the chlorine generation and even the microwave treatment on a porch (outside) as potential Cl2 escape is less problematic on a windy porch.

[Edited on 2-1-2022 by AJKOER]

Ubya - 2-1-2022 at 09:04

Quote: Originally posted by AJKOER

My advice, keep it simple.

OK, for small amounts, a vessel with plastic lid is filled with Chlorine gas containing a small piece of say pure Al foil. The amount of the Al foil converted to moles should be, at least, 1/3 of the volume of the vessel in liters divided by say 22.5.

Place the vessel in a microwave for a few seconds treatment. The Al foil ignites and burns in the atmosphere of chlorine. Cool the vessel to collect the AlCl3 using ethanol in which AlCl3 is soluble.

Yes, a small amount of AlCl3, but as the equipment goes, just an appropriate vessel and a microwave. I have used a large thick glass vessel with a plastic cap to avoid any cracking or burning if the Al strip comes into contact with the vessel where the Al creates, in the microwave, a very hot plasma electric arc. I have just suspended the Al strip like a light bulb filament from the cap of the vessel which has a hole from being punctured with a nail. As a result, no cracking of the vessel (which has occurred with me using even a porcelain vessel).

Impurities from water vapor or oxygen presence, not much of an issue as the latter products are not likely soluble in the ethanol, but some lose in yield.

Simple and quite exciting display with the arc and subsequent burning. Small yield however, but can repeat the process to concentrate your ethanol collecting solution.

For safety, I would recommend safety equipment and also performing the chlorine generation and even the microwave treatment on a porch (outside) as potential Cl2 escape is less problematic on a windy porch.

[Edited on 2-1-2022 by AJKOER]

while surely being the best solution for big scale production i don't think this is a viable method for a small amount of pure product.
chlorine generation and drying is by itself a chore, you'll need a lot of chlorine to flush your vessel and keep it oxygen free, and in the end the product will still be pretty dirty from unreacted aluminium, aluminium oxide and hydroxide dust/particles, so you need another step, filtering a moisture sensitive solution before evaporating away the solvent.

the beauty of the zinc chloride/aluminium reaction is that you can get fairly pure aluminium chloride in anhydrous form by just heating 2 solids, and not having to deal with chlorine gas and combustion of aluminium metal.

in the end it is just a personal preference, making, purifying and drying zinc chloride and making the right equipment, or just burn some aluminium foil in a chlorine atmosphere like many of us have done, but then having more steps to clean the product.

Texium - 2-1-2022 at 09:49

Ethanol reacts with anhydrous aluminum chloride to form aluminum ethoxide.

AJKOER - 2-1-2022 at 09:59

I agree with Ubya comments, but I believe any Al2O3 presence, for example, is easily separated out by decanting the ethanol wash. Water may further add a small amount of another insoluble basic Aluminum salt.

Any water in the ZnCl2 again leads to reduce yield.

Note: to industrially prepare very high purity HCl, for example, one burns H2 in Cl2, so the burning process is more of a high purity path.

One problem I have with Zinc metal as the source of ZnCl2, is that the underlying Zn may not be very pure. Case in point, adding pure HCl to Zn may NOT result in just odorless and non-toxic H2, but some rather somewhat noticeable and highly toxic, albeit in small amounts, but still problematic, fumes, as noted here https://beta-static.fishersci.com/content/dam/fishersci/en_U... .

Per another source: 'On the Impurities of Commercial Zinc, with Special Reference to the Residue Insoluble in Dilute Acids, to Sulphur, and to Arsenic', here https://www.jstor.org/stable/25057965?seq=1#metadata_info_ta... to quote:

"The common zinc of commerce generally contains a portion of lead, copper, iron, traces of arsenic and manganese, and a little plumbago..."

With respect to known impurities in ZnCl2, an interesting comment in this patent https://patents.google.com/patent/US3148944 to quote:

"It is known, by way of more specific delineation, that manganese as an impurity in zinc chloride solutions used for the preparation of acrylonitrile polymer products seriously affects the acrylonitrile polymerization rate and the control of polymer molecular Weight."

where manganese was a previously cited as a common impurity.

Further, chemistry may accentuate the impurity problem as, for example, the action of HCl on Zinc forms active atomic hydrogen (or the hydrogen radical, ...), where the action of H with an impurity like Pb or As can form very problematic toxic gas, for example, arsine gas (AsH3).

It makes me believe that starting with problematic chlorine gas (even containing water vapor and oxygen) and 100% pure Al foil may be worth the effort if product purity/safety is an issue.

[Edited on 2-1-2022 by AJKOER]

Ubya - 2-1-2022 at 10:35

Quote: Originally posted by AJKOER

I agree with Ubya comments, but I believe any Al2O3 presence, for example, is easily separated out by decanting the ethanol wash. Water may further add a small amount of another insoluble basic Aluminum salt.

Note: to industrially prepare very high purity HCl, one burns H2 in Cl2, so the burning process is more of a high purity path.

Any water in the ZnCl2 again leads to reduce yield.

One problem I have with Zinc metal as the source of ZnCl2, is that the underlying Zn may not be very pure. Case in point, adding pure HCl to Zn may NOT result in just odorless and non-toxic H2, but some rather somewhat noticeable and highly toxic, albeit in small amounts, but still problematic, fumes, as noted here https://beta-static.fishersci.com/content/dam/fishersci/en_U... .

Per another source: 'On the Impurities of Commercial Zinc, with Special Reference to the Residue Insoluble in Dilute Acids, to Sulphur, and to Arsenic', here https://www.jstor.org/stable/25057965?seq=1#metadata_info_ta... to quote:

"Wittstein says that the metals with which common zinc is contaminated are iron, cadmium, tin, and lead"

It makes me believe that starting with chlorine gas (even containing water vapor and oxygen) and 100% pure Al foil is worth the effort if product purity is an issue.

Yes, the zinc chloride must be as dry as possible, if one wants to be as diligent as possible you could put the moist/not very dry zinc chloride direcly in the reaction apparatus, pull a vacuum, and heat the salt to 150-200°C, let it cool a bit under vacuum, then add the dry aluminium powder (yeah i'd also quickly dry it) and start the reaction.

I wouldn't worry much about metal impurities in the zinc, their chlorides have all pretty high melting/boiling points compared to aluminium chloride
(CdCl2 964 °C, PbCl2 950 °C, SnCl2 623 °C, FeCl2 1,023 °C, ZnCl2 732 °C, AlCl3 180 °C)

Any sulphur compounds in the impure zinc metal get turned into volatile gases (probably H2S) when reacting with HCl, so while the hydrogen produced isn't very pure, the metal salt solution should get "purer".

Without proper analysis of the aluminium chloride produced it's all just speculation, maybe @Neptunium could help us out xD

AJKOER - 2-1-2022 at 10:52

True but I have since added "With respect to known impurities in ZnCl2, an interesting comment in this patent https://patents.google.com/patent/US3148944 to quote:

"It is known, by way of more specific delineation, that manganese as an impurity in zinc chloride solutions used for the preparation of acrylonitrile polymer products seriously affects the acrylonitrile polymerization rate and the control of polymer molecular Weight."

where manganese was a previously cited as a common impurity found in Zinc metal.

So, depending on the application intended, a more pure path (with respect to various metals) may be warranted.

However, your melting point data does suggests that the ZnCl2 path may, indeed, result in a pure product.
============================

Did find a lab experiment for feeding generated Cl2 into a heated Al source here https://edu.rsc.org/experiments/reactions-of-chlorine-bromin...

Also, a Youtube on the Cl2 and Al reaction here https://www.youtube.com/watch?v=aK85PZX2xNE . Interesting, preheat a ball of Al foil with a blowtorch and just drop it in the open reaction vessel. No need for microwave ignition. Suggests to me adding a long exit tube to capture generated AlCl3. Appears to be rather simple but highly exothermic.

[Edited on 2-1-2022 by AJKOER]

Ubya - 2-1-2022 at 11:01

but again even if manganese is present it shouldn't be an issue in the production of pure anhydrous aluminium chloride by its reaction with aluminium powder.
Aluminium chloride sublimes from the reaction RBF and condenses in the receiving flask, unreacted zinc chloride, and all other metal salts impurities will be left behind in the reaction vessel. Impurities in the zinc chloride would have been an issue if the reaction was in acqueous solution (something i had to deal with when making ultrapure zinc chloride to then convert it to zinc sulfide, an attempt to make phosphorescent powder), but in this case, even quite dirty zinc choride should be fine, as long as it doesn't have impurities with low boiling points or high vapour pressures

AJKOER - 3-1-2022 at 07:27

Just noticed a cited reaction between Al and NH4Cl quoted as proceeding as follows:

2 Al + 6 NH4Cl → 2 AlCl3 + 3 H2 + 6 NH3

So, I guess one could mix Al powder with dry NH4Cl and heat, and 'collect' the AlCl3 powder. Note, some explosion danger if the hydrogen gas is allow to accumulate in a room, while the ammonia is passed into water.

The above reaction parallels the below reaction noted at https://en.wikipedia.org/wiki/Copper(II)_chloride

CuO + 2NH4Cl → CuCl2 + 2NH3 + H2O

Along with the reaction involving Cu in place of Al:

Cu(s) + Cl2(g) → CuCl2(l)

So, NH4Cl basically replaces problematic Cl2 with unfriendly NH3 (albeit, the latter is easily pass into water/dilute HCl) and introduces a potential H2 explosion hazard (unless the experiment is performed in a fume hood or outdoors).

[Edited on 3-1-2022 by AJKOER]

clearly_not_atara - 3-1-2022 at 09:59

AlCl3 will not dissociate from NH3. You'll get a complex that is no less tenacious than AlCl3*6H2O. In fact this is a pathway to preparation of aluminium nitride:

https://pubs.acs.org/doi/pdf/10.1021/acsomega.9b01140

However, I'm pretty sure that ZnCl2 in this rxn can be replaced by MnCl2, with the advantage that the latter can be dried by mere heating rather than complex methods (also involving ammonia in this case).

AJKOER - 3-1-2022 at 11:14

Clearly_not_atara:

For the record, the correct reaction thus appears to be:

2 Al + 6 NH4Cl → Al2Cl6·2NH3 + 3 H2 + 4 NH3

for temperatures above NH4Cl sublimation temperature of 338 °C and under 400 °C, otherwise at higher temperatures, more including AlCl3 (g) and possibly AlN.

Note, in the cited reference, the author only presents a series of equations and does not attempt to present a net reaction. So, why avoiding technical speciation and temperature dependent product issues, it unfortunately allows the promulgation of totally misleading net reactions.

[Edited on 3-1-2022 by AJKOER]

SuperOxide - 4-1-2022 at 07:46

Quote: Originally posted by Triflic Acid

You should probably use some acidic dehydrating agent since any suspended NaOH or Na2O from the sodium drying will destroy your product. But P2O5 should be fine

How about sodium/lead mixture as a drying reagent for the ether? I hear that works very well. I do have like... 600g of P2O5, but I cherish it and try not to use it if I don't need to, lol. It's much easier for me to get sodium than it is to get P2O5.

Quote: Originally posted by Ubya

Quote: Originally posted by AJKOER

I agree. I think heating the zinc chloride and Al powder is actually pretty simple, and seems to make a decently pure product. It seems like the difficult part is keeping the product from escaping the apparatus.

What about keeping some solvent in the receiving flask? That way the aluminium chloride will actually go right into solution. Not sure if that would help much, but thought I'd ask.

Ubya - 4-1-2022 at 12:04

Quote: Originally posted by SuperOxide

What about keeping some solvent in the receiving flask? That way the aluminium chloride will actually go right into solution. Not sure if that would help much, but thought I'd ask.

not if the flask is uner vacuum xD

SuperOxide - 6-1-2022 at 15:34

Quote: Originally posted by Ubya

Quote: Originally posted by SuperOxide

What about keeping some solvent in the receiving flask? That way the aluminium chloride will actually go right into solution. Not sure if that would help much, but thought I'd ask.

not if the flask is uner vacuum xD

Well yeah, but as discussed earlier - I could pull the vacuum, seal it, then start the reaction. I know some would evaporate off while the vacuum is being pulled (which could maybe be prevented to a certain extent by cooling it down as much as possible), but then once it's pulled and the valve is closed, any of the aluminium chloride that comes over should go into solution.

Idk if it would be beneficial at all, I'm just thinking out loud

Thanks.

neptunium - 13-1-2022 at 12:00

Quote: Originally posted by Ubya

Without proper analysis of the aluminium chloride produced it's all just speculation, maybe @Neptunium could help us out xD

Hummm..... sounds like a job for the ICP-OES... interferences from Chloride will ruin the signal on MS ..

Unfortunately I do not own one.....yet

AJKOER - 14-1-2022 at 08:44

Interestingly, did find a commercial path in this reference, "Oxidative Dissolution of Metals in Organic Solvents", found here https://pubs.acs.org/doi/pdf/10.1021/acs.chemrev.0c00917 for the general dissolution of metals in an appropriate organic solvents empolying a dissolved halogen, which may be of use.

For example, in the case of aluminum metal and an organic solvent like methanol (MeOH), per Table 1 containing say Chlorine can apparently be used to dissolve Aluminum. Also, this statement:

"Metal processing in organic solvents or mixed with small amounts of water (<50 vol %) is termed “solvometallurgy”, which is an emerging branch of extractive metallurgy and complementary to hydrometallurgy and pyrometallurgy. [11]"

And further:

"Addition of water can enhance the oxidative dissolution of metals in chlorine-DMF. [27] For example, when chlorine in DMF was used as a leaching system, only small amounts of rhenium and tungsten (<5%) could be oxidatively dissolved. However, an increase in the water content in the leaching system resulted in a higher amount of metals dissolved in the solution. Further increase in the water content resulted in a decrease of the dissolution of molybdenum and tungsten (Figure 4). "

where, I would speculate, in the case of Al with a protective Al2O3 coating, the presence of chloride (from wet chlorine or drops of water) may be beneficial in starting the reaction. Other paths involve the use of Hg, as is cited in a referenced thread provided below.

Also, cited Reaction (1) with a metal M creating a metal chloride as, for example, with elemental Bromine:

M + x/2 Br2 = MBrx (1)

On Page 4510, it is further noted that "It is known halogens can react with organic solvents, [18],[36] ..", which was confirmed below.

Further research produced this link to a former SM thread on precisely this topic here https://www.sciencemadness.org/whisper/viewthread.php?tid=28... and interesting comments. Note, the referenced video in the opening thread still works, albeit, on bleaching/restoring faded plastic, and after many minutes, the prep of a Cl2/MeOH solution, as the active agent. The bleaching agent here is Cl2 and, I suspect, with a further application of violet and especially ultra violet light (per the classic chain reaction demo of H2 and Cl2 in light https://www.youtube.com/watch?v=tJhVy1x9X2c), the active chlorine radical itself (see, related research here https://www.researchgate.net/publication/347808849_Reactivit...). Note, more generally, a photo-assisted path suggesting the use of radicals has been cited in the leaching literature (see, for example, Table I in this 2021 work: "The Role of Solar Energy (UV-VIS-NIR) as an Assistant for Sulfide Minerals Leaching and Its Potential Application for Metal Extraction", in Minerals 2021, 11, 828, https://doi.org/10.3390/min11080828 ).

I recommend attempting only when appropriate safety measures are in place.

[Edited on 14-1-2022 by AJKOER]