Chlorothermal reactions (chloride thermites) for preparing AlCl3 and KAlCl4

WARNING: This is a hypothetical thread of untested and potentially highly dangerous thermite-like reactions. Use at your own risk!

Anhydrous aluminium chloride is a highly useful reagent, for example, as a lewis acid and for forming tetrachloroaluminates.

It sublimes relatively easily at 180°C (ref. Wiki) and this is one method of purifying it. While there are various potential routes to preparing AlCl₃, for example, reacting the metal with either anhydrous HCl or Cl₂, the idea occurred to me that there might be a simpler and purer way, by exploiting a chlorothermal reaction (chlorine thermite) that simultaneously sublimes the aluminium chloride formed and so generates a very pure product. The question is, what would be, hypothetically speaking, a suitable candidate?

The generic equation for such a reaction would be:

aluminium + metal 'x' chloride salt => aluminium chloride vapour + molten metal 'x'

Ideally, as anhydrous aluminium chloride is far more useful than hydrates, we would prefer an anhydrous metal 'x' chloride salt. This poses a problem as most solid transition metal chlorides are strongly hygroscopic and usually isolated as a hydrate salt of some form.

So, to start off, the best metal 'x' I could come up with is lead:

2Al(s) + 3PbCl₂(s) => 2AlCl₃(g) + 3Pb(l)

PbCl₂ is a white insoluble salt that doesn't form hydrates, so this is very attractive!

Furthermore, as we don't want the reaction to be too exothermic (since a gas is being produced), it is nice that both lead and lead (II) chloride have rather low melting points of 328°C and 501°C, respectively (ref. Wiki), so all we need now is for the reaction to be only moderately exothermic, ideally proceeding a little above 500°C, but not too much so for lead chloride is volatile, boiling at 950°C! (ref. Wiki)

So what does thermodynamics say?

I've calculated a hypothetical standard heat of reaction for the equation above (see attached) using NIST webbook data.

The hypothetical dH_rx at 20°C is -78kJ and dG_rx is -206kJ, so this looks very promising as it's exothermic, but not too much so.

Note: this was computed for molten lead and aluminium chloride vapours at 20°C which has no meaning physically, but can be used to calculate an adiabatic reaction temperature using the heat capacity functions for the components.

Lead (II) chloride, also known as cotunnite, is easily prepared by exploiting its low solubility in water, for example:

Pb(NO₃)₂(aq) + 2NaCl(aq) → PbCl₂(s) + 2NaNO₃(aq)

(Ref. Wiki)

Finally, should the exothermicity of the chlorothermal reaction prove a tad too low, one could potentially remedy it by the addition of a small amount of PbO well mixed with the PbCl₂, which would up the heat significantly via a standard thermite type reaction, but only if this should prove necessary!

****
For those interested in preparing anhydrous KAlCl₄ (that's you Zyklonb!), this reaction could be modified by adding KCl to the mix. This has a huge advantage of potentially acting as a flux and elliminating the sublimation. The only other product should be molten lead which would sink and solidify at the bottom and could be easily removed:

2Al(s) + 3PbCl₂ + 2KCl(s)=> 3Pb(l) + 2KAlCl₄(l)

Also, since KCl and AlCl₃ can form low melting salt mixtures (thanks blogfast for the phase diagram), KCl added in substoichiometric amounts could help the reaction proceed to form AlCl₃ if temperatures are not that high without the need for PbO to raise heats, however, there would be a slight loss in AlCl₃ yields due to the formation of some KAlCl₄.

?! Relatively direct route to anhydrous KAlCl₄ ?!

Attachment: dHrx and dGrx calculations for Al PbCl2 chlorothermal reaction.xlsx (9kB)
This file has been downloaded 679 times

* Credit to blogfast for finding the phase diagram image attached here



[Edited on 17-4-2014 by deltaH]

Quote: Originally posted by deltaH

I am very curious to see if you add stoichiometric amounts of KCl, if you could make molten Zn metal and KAlCl4 using anhydrous ZnCl2.

Agreed, though the NIST doesn't have liquid data for KAlCl₄, nor a heat of fusion

Nevertheless, I've been meaning to expand my EXCEL sheet with an adiabatic temperature calculation for the PbCl₂ case, the data there is complete (at least for AlCl₃(g) version).

Quote: Originally posted by deltaH

Agreed, though the NIST doesn't have liquid data for KAlCl4, nor a heat of fusion

Quote: Originally posted by blogfast25

I do think capturing the AlCl3 may be easier said than done: smoke implies the AlCl3 has already condensed and is now a suspension of fine solids in air. To capture that may be challenging...

Seems like this is also a most excellent way of producing masses of white smoke, no?

On a more serious note, I think the smoke problem would be solved if this was going through glassware and the AlCl₃ were condensing on existing crystals growing from the walls as opposed to very rapidly in cold air.

This is analogous to steam from a pot, boil it open and you get masses of steam (smoke) as the vapour mixes with cold air quickly, but put it through a condenser and you just form liquid water nicely.

Anyhow, the safety of the reaction would have to be considered... the alarming acceleration could be problematic.

As for the PbCl₂, I can understand that it's not very interesting to you because you have plenty of anhydrous ZnCl₂, but for those who don't, I've gone ahead and calculated an adiabatic reaction temperature for the reaction assuming AlCl₃ vapours are produced.

I've assumed that the Cp of aluminium chloride vapour and molten lead will be constant over small enough temperature ranges (I'm lazy/am an engineer, take your pick) and calculated a temperature of about 320°C. It's a tad low, but I suppose this way it's a little safer in terms of slowing down the rate of the reaction, though I think heating a mass of PbCl₂ an Al in its entirety (as in a crucible), is a very bad idea, I would say you have to start with a cold heap that you ignite from the top with burning magnesium ribbon or whatever.

For the amateur chemist, reacting lead with nitric acid to form a solution of lead nitrate and then precipitating that solution with sodium chloride is trivial. After the chlorothermal reaction is done, you would have regenerated a lump of lead... so it's pretty neat.

I think this could work well for making potassium tetrachloroaluminate when using stoichiometric KCl and in that case you don't worry so much of the danger of producing large volumes of vapours rapidly. It's probably too fast for safely making AlCl₃, where the method with ZnCl₂ might be better suited.

[Edited on 17-4-2014 by deltaH]

Photos please if you can ... and be safe!

Quote: Originally posted by deltaH

Do fishing sinkers have the same problem? [Edited on 17-4-2014 by deltaH]

Hi bismuthate, long time no see, glad to hear from you again

I'm a bit concerned about oxidising salts, like copper(II), iron(III), etc. Also, these can form hydrates, although they do dry up nicely by conventional means. Remember, this reaction produces large volumes of gas, so very worried that if any but the strongly reducing metals like zinc, etc. are used, it's going to be way too violent!

That said, it might be worth thinking about lower oxidation state transition metals, e.g. copper (I) chloride or manganese (II) chloride, but maybe these are purely academic (for practicality reasons?).

Tin (II) chloride is maybe also an option, but I still don't think it beats zinc chloride (if you have the anhydrous kind) or lead (II) chloride (if you have to go the DIY route).

Can't wait for blogfast's write up. I've got a test to take today but this should be a nice spoil afterwards

******

This line of thinking does lead me down another related train of thought, that of preparing things like titanium tetrachloride and tin tetrachloride from titanium powder and anhydrous zinc chloride with heating, and tin metal and something mildly oxidising like anhydrous copper (II) chloride, respectively?

I think this might make blogfast's day if it could work? I myself have worked with titanium tetrachloride and found it a very usefull reagent for many things!

For those who don't know, be warned, titanium tetrachloride is a volatile liquid that smokes like crazy in air because of the fast reaction with ANY moisture forming TiO₂ and irritating HCl fumes!

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Okay, looking at the thermos of

Ti(s) + 2ZnCl₂(s) => TiCl₄(l) + 2Zn(s)

I'm considering the reaction with TiCl₄ as a liquid because making a gas from it is just a question of heating as it boils at 136.4°C (ref. Wiki).

The NIST webbook unfortunately doesn't have data for zinc chloride , but I managed to find a dHf,std from Wikipedia for ZnCl₂, some small error because the NIST uses 293.15K as its standard state while the Wiki data is reported for 298K.

Anyhow, for my reaction as written, the dHrx, standard is 26kJ/(mol Ti), this puts it in the 'maybe' catagory with heating depending on what dGrx is. Again, the fact that we make a volatile product plays in our favour as it allows us to pull a moderate equilibrium to the right.

Does anyone have a dGf, std or entropy of formation value for ZnCl₂ on hand?

[Edited on 18-4-2014 by deltaH]

ΔH⁰_f for ZnCl2 = - 415.1 kJ/mol
ΔG⁰_f for ZnCl2 = - 369.4 kJ/mol
(CRC ’86)

Here are the results of the second test.

Apparatus:



Left: large, thick walled test tube (‘boiling tube’), middle: connecting glass tubes and stopper, right: condenser on ice bath.

1.2 g of Al powder and 7.6 g anh. ZnCl2 were ground up and loaded into the reactor and everything closed off. Heating then commenced, somewhat more cautiously than before. Reaction started almost immediately:



I removed the Bunsen and the reaction continued to proceed! The surface of the charge appeared to be bubbling. So we have a bit of a Goldilocks situation here: this reaction is just energetic enough to be self-sustaining without being violent.

Once the apparatus had heated enough the smoke clears for clear AlCl3 vapours which condensed without problems in the condenser (I’ve yet to have a closer look at the solid AlCl3).

I also got to witness melting AlCl3: near the top of the tube sheets of solid AlCl3 had accumulated but on heating these promptly melted and collected further down. By more or less heating the entire tube and connecting tubing I collected most of the AlCl3 where I wanted it to be.

The condenser shape is not ideal and will be replaced by something more suitable this afternoon, to do a decent yield determination.

At the end of the reaction: dirty zinc powder (presumed):



I want to do a few more tests but all in all this is very promising as a simple lab bench preparation of AlCl3. It sure beats messing with Cl or HCl generators, IMHO!


[Edited on 18-4-2014 by blogfast25]

deltaH:

Interesting what you said about the TiCl4. I’ll probably try that soon, just to see. Quite excited, really!

The latest test for AlCl3 was done using a modified apparatus (photo of dress rehearsal):




Apart from the condenser, the tubing had been changed over to a wider ID.

I doubled the charge also to 2.4 g Al and 15.6 g ZnCl2 anh.

I can confirm that once the reaction is initiated it proceeds by itself. All went well until about half way, when a blockage of solid AlCl3 in the tube running through the left hand side rubber stopper caused it to pop off. I’d clamped and stoppered everything rather loosely to keep damage low in case of pressure build up, so no problems occurred, apart for a lot of smoke escaping. I unblocked it and carried on but a few minutes later the same thing happened and I had to abort.

I got about 0.3 g of snow white solid AlCl3 for my trouble, low yield due to the interrupting blockage:




In short, this is the right type of reaction but the wrong geometry. I now envisage a fairly wide combustion tube, stoppered at the left end, then some dry sand (to protect the stopper), then glass wool (to block the sand), then the reagent charge. At the right end of the tube would be mounted an ice cooled recipient but without any thin connecting tubing, where AlCl3 could condense and solidify. This set up would avoid blockages caused by condensing AlCl3.

Please be aware that those wanting to use Cl or HCl generators as chlorinating agents may also be plagued by blocking caused by prematurely solidifying AlCl3! With Cl or HCl gas such a blockage springing a leak could potentially be seriously dangerous.

With regards to a CuCl2 based reaction, someone already reported that here: it behaves like a tamish thermite. For CuCl2 the Standard Heat of Formation is only - 205.85 kJ/mol, so that reduction reaction is exothermic by another additional 200 kJ/mol, with respect to the ZnCl2 reduction!

Edit:

@deltaH:

How do you get to your values for the Ti chlorination?

For TiCl₄ from CRC ’86 I get:

ΔH⁰_f = - 804.2 kJ/mol
ΔG⁰_f = - 737.2 kJ/mol
Both for liquid TiCl4.

For 2 ZnCl₂(s) + Ti(s) →TiCl₄(l) + 2 Zn(s) that gives me:

ΔH⁰_reaction = + 26 kJ/mol
ΔG⁰_reaction = + 1.6 kJ/mol



[Edited on 18-4-2014 by blogfast25]

[Edited on 18-4-2014 by blogfast25]

Maybe you might want to try insulating the parts of your apparatus where you don't want condensation to occur. Even wrapped cloth turnings (bandage) should do the job.

Your dGrx looks far more believable than mine, exactly what I would expect, something slightly positive. It's weird, NIST webbook quotes dHf for TiCl₄ as -804.16kJ/mol (same as yours) and a dSf of 221.93J/mol.K, so my dGf TiCl₄ is: dGf = dHf -T.dSf = -804.16kJ/mol - (293.15K)(221.93J/mol.K)(1 kJ/1000J) = -869.2kJ/mol ... and there's the problem, it's much lower than your CRC sourced one of -737.2kJ/mol.

Unless I made some silly algebra error, I don't believe the NIST one, it looks highly suspect, let's work with yours which looks realistic (slightly unfavourable equilibrium).

Out of interest, a dGrx = +1.6kJ/mol gives a Keq = exp(-1600/8.314/298) = 0.5 @298K (I assume the CRC data is at 298K ?), so this this should tick over nicely with heating to remove the TiCl₄ formed.

To be honest, if this can be shown to work, I would find it remarkable.

Anyhow, best of luck... break a beaker

*****
I'd like to briefly go off on a somewhat philosophical tangent of sorts to this topic. We have touched on it already, that of these reactions being replacements for less practical HCl or Cl₂ routes.

Firstly, let's consider anhydrous HCl. In many respects, this can be seen as a 'salt' of H⁺ and Cl^- where the hydrogen behaves as a weakly reducing metal would and has a corresponding standard reduction potential of 0V. Zinc is similar, but being more strongly reducing with a reduction potential of -0.76V.

The result is that anhydrous zinc chloride behaves as a concentrated, solid sluggish HCl. It is thus well suited for inorganic reactions where one would want a practical alternative to anhydrous HCl and at the same time, where the metal you are reacting it with is highly electropositive, in which case the 'sluggishness' comes in handy.

So far we have considered and blogfast has kindly volunteered his time to investigate practically:

(i) 2Al(s) + 3ZnCl₂(s) => 2AlCl₃(g) + 3Zn(s) + heat
(ii) Ti(s) + 2ZnCl₂(s) <=> TiCl₄(l) + 2Zn(s)

The second reaction is an endothermic equilibrium at STP with a Keq of about 0.5 based on dG calculations. It thus requires heating to volatilise the low boiling TiCl₄ and so drive the reaction to the right and to completion:

(ii)* Ti(s) + 2ZnCl₂(s) + heat => TiCl₄(g) + 2Zn(s)

Now for cases like this, one may want something more closely related to anhydrous HCl to make it less sluggish. Lead (II) chloride is suited to this because lead has a standard reduction potential of -0.13V, quiet close to that of hydrogen. In this way, it may be considered a close replacement to anhydrous HCl or even a solid alternative to it.

This may better suite reactions where a 'sluggish' HCl alternative won't suffice, for example in preparing TiCl₄ from Ti.

But this opens scope up for other chlorides too, particularly the volatile ones. The extra 'umph' of PbCl₂ coming in handy to volatilise the higher boiling transition metal chlorides, for example VCl₄ and perhaps even ZrCl₄ and NbCl₅.

The we can progress further and look toward things that may be considered a practical replacement for Cl₂, this would require a oxidising metal. In this way, anhydrous CuCl₂ can be considered a 'sluggish' replacement for Cl₂.

This could be useful, for example, in generating exotic chlorides like chromium (IV) chloride which is a gas and only stable at high temperature, reverting to CrCl₃ at room temperature (ref. Wikipedia). In this way, the reaction may be used to produce a gaseous product that can be boiled off then condensed, reverting to high purity anhydrous CrCl₃ and Cl₂.

One quickly realised that there is MUCH scope for experimentation and preparative inorganic chemistry here!

[Edited on 19-4-2014 by deltaH]

Does titanium passivate in chlorine? I thought it only does that in oxygen.
It is too bad that making titanium chloride wasn't working out. I might want to try some at some point, I lack Ti though.

Quote: Originally posted by deltaH

Since the thermo's say this should work, I'm inclined to think that something is preventing it from starting... hence I turn to the usual suspect in such cases. [Edited on 20-4-2014 by deltaH]

The thermos say nothing about kinetics, of course. Maybe it's just dead slow or needs a catalyst of sorts.

On the plus side, if you need something that resists boiling ZnCl2, look no further

I ran another AlCl3 preparation, this time with a simple and effective geometry:



A large borosilicate boiling tube (OD about 25 mm, length 145 mm) is horizontally clamped up, in such a way that its neck penetrates into a 250 ml Simax wide necked reagent bottle, which will act as a condenser flask.

A freezer bag containing some ice and a bit of water is then draped over the reagent bottle, the reagent mix loaded and ignited.

The AlCl3 vapours collect in the condenser flask and condense (despite the rather large gap, with minimal losses). The flow of AlCl3 can be regulated somewhat by adding gentle heat from a medium hot Bunsen burner or by withdrawing it.

I would recommend using a similar set up to anyone preparing AlCl3 (using whatever chlorinating agent) because the risk of blockage by solid condensed AlCl3 is basically zero.

[Edited on 4-5-2014 by blogfast25]

Quote: Originally posted by Zyklonb

No chlorine, no gas generators, no tubes or SS pipes, condensing is so easy.

Quote: Originally posted by deltaH

Excellent work blogfast, I'm happy you got it to work!

Chemplayer made a video of this reaction

https://www.youtube.com/watch?v=g7sS69fQMsk

Yields are crap though, weird.

[Edited on 19-2-2016 by deltaH]

Quote: Originally posted by deltaH

Chemplayer made a video of this reaction https://www.youtube.com/watch?v=g7sS69fQMsk Yields are crap though, weird.

Quote: Originally posted by JJay

Zinc chloride and aluminum definitely do react to produce aluminum chloride, but capturing the product is tricky. I have tried copper (II) chloride/aluminum thermite, and it definitely produces AlCl3, but it is tricky to control the rate of the reaction. It may be possible to run it in ethyl acetate (or not, and that might be dangerous... not sure). I am getting a 125 cm quartz tube in a few days and plan to try out the classic method with anhydrous HCl. I think by heating one end of the tube with nichrome wire and cooling the other with copper tubing, I should be able to convert aluminum to AlCl3 almost quantitatively.

Quote: Originally posted by JJay

Right, but in water, the reaction proceeds at room temperature.

Quote: Originally posted by JJay

Oh for zinc chloride and aluminum? I just got some media bottles and a thermometerless distillation adapter in the mail today.