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Author: Subject: Chlorothermal reactions (chloride thermites) for preparing AlCl3 and KAlCl4
deltaH
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[*] posted on 17-4-2014 at 02:04
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 dHrx at 20°C is -78kJ and dGrx 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 282 times

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

AlCl3-KCl.jpg - 69kB

[Edited on 17-4-2014 by deltaH]




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[*] posted on 17-4-2014 at 04:17


Also worth considering is:

ZnCl<sub>2</sub> + 2/3 Al === > Zn + 2/3 AlCl<sub>3</sub>

because ΔH<sup>0</sup> and ΔG<sup>0</sup> are about - 50 kJ/mol (based on CRC '86 data), so thermodynamically favourable but only barely exothermic.

Ideally, you want something with ΔH<sup>0</sup> ≈ 0 because then by heating reagent mix, the AlCl<sub>3</sub> reaction product would distil off, pulling the equilibrium from left to right.

Adding KCl to the mix might have a further 'tempering' effect because it will act as a heat sink. But I doubt that KAlCl<sub>4</sub> will form, as the KCl won't melt ad the AlCl<sub>3</sub> will tend to evaporate out. All preparations of chloroaluminates point to closed reactors.

For PbCl<sub>2</sub> it's probably better to use HCl as the precipitant, because it will lead to less occlusion than NaCl and can be driven off more easily.


[Edited on 17-4-2014 by blogfast25]




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[*] posted on 17-4-2014 at 04:35


Thanks blogfast.

I was reading through the Wiki article on ZnCl2 and encountered this under preparation:

Quote:
Anhydrous samples can be purified by sublimation in a stream of hydrogen chloride gas, followed by heating the sublimate to 400 °C in a stream of dry nitrogen gas. Finally, the simplest method relies on treating the zinc chloride with thionyl chloride.[12]


That's a little troubling, do you know if zinc chloride hydrate is one of those salts that cannot be dehydrated by conventional means (heating/desiccator)?




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[*] posted on 17-4-2014 at 04:46


You can buy anh. ZnCl2 quite cheaply. From me among other people.

I plan to run a quick test with this soon, so hold your fire.

Was the ΔH value for the PbCl2 reaction calculated per mol or for 2 mols?

[Edited on 17-4-2014 by blogfast25]




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[*] posted on 17-4-2014 at 04:51


I look forward to your results!

The dH is per my equation written, so per 2 mols.

****
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.

It would make Zyklon's day I'm sure!

[Edited on 17-4-2014 by deltaH]




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[*] posted on 17-4-2014 at 07:39


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.



That can be calculated, if you start from an adiabatic enclosure. Use NIST Shomate equations to calculate adiabatic end temperature (assume that AlCl3 and KCl won't 'merge' because there won't be any thermochemical data on KAlCl4. The formation of the latter may yield a little extra enthalpy)

420 C (MP of Zn) 'sounds' a little high for 50 kJ/mol reaction enthalpy but only calculation or experiment can really tell.




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[*] posted on 17-4-2014 at 07:49


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).




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[*] posted on 17-4-2014 at 08:05


Quote: Originally posted by deltaH  
Agreed, though the NIST doesn't have liquid data for KAlCl4, nor a heat of fusion :(



Like I said, for now just ignore KAlCl4: treat it as a physical mixture of KCl and AlCl3. It's a relatively small error.




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[*] posted on 17-4-2014 at 08:50


Well, well. A first test for ZnCl2 + Al was both interesting and encouraging.

3.8 g of anh. ZnCl2 powder was mixed with 0.5 g of superfine Al powder (stoichiometric ratio) and placed in a nickel crucible without lid. I applied max heat (propane Bunsen with full open air inlet) and almost immediately white fumes started coming off. This then intensified to a somewhat alarming formation of clouds, then smoke formation gradually subsided. Never did the mixture show any tendency to 'light up', glow or deflagrate. At the end of the test a greyish powder was left in the crucible which I'll investigate.

At first sight this would allow preparing AlCl3 with nothing more than a test or boiling tube, by gradually heating the mixture (perhaps with extra Al) and condensing the smoke using some connected condenser or other.

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...


[Edited on 17-4-2014 by blogfast25]




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[*] posted on 17-4-2014 at 09:20



This sounds very promising indeed! Well done and thanks for trying it out.

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]




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[*] posted on 17-4-2014 at 10:21


I think the PbCl2 idea is well worth pursuing but you won't get KAlCl4 without some containment, of that I'm pretty certain. That means a bomb reactor but you'd also need that to prepare the KCl/AlCl3 eutectic. Calculated adiabatic end temperature might not even be that important: just heat to about 450 C for 20 - 30 minutes.

"For the amateur chemist, reacting lead with nitric acid to form a solution of lead nitrate [...]"

Only if you've got good quality lead. Forget battery lead: the reaction 'stalls' early on. Still, lead nitrate or acetate aren't difficult to buy.

Time allowing, I'll run another test with ZnCl2 tonight, this one with a cold trap of sorts.

[Edited on 17-4-2014 by blogfast25]




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[*] posted on 17-4-2014 at 10:28


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

Quote:
Only if you've got good quality lead. Forget battery lead: the reaction 'stalls' early on.

I did not know that. Do fishing sinkers have the same problem?

[Edited on 17-4-2014 by deltaH]




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[*] posted on 17-4-2014 at 13:13


Quote: Originally posted by deltaH  
Do fishing sinkers have the same problem?

[Edited on 17-4-2014 by deltaH]


Most lead contains hardeners (alloying agents), these seem to be causing the problems.

Another run with ZnCl2/Al was surprisingly successful. Recovery of the AlCl3 proved to be a doddle. Story and photos tomorrow...

[Edited on 17-4-2014 by blogfast25]




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[*] posted on 17-4-2014 at 15:26


Excellent topic, I get my lead from pellets (for a pellet gun), It seems to be quite pure. The best place is to buy it online.
What about copper (II) chloride + aluminum? That should be quite energetic, enough to keep the reaction going once it has been initiated. Copper (II) chloride is quite easy to make, copper sulfate + calcium chloride comes to mind - among many other ways.




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[*] posted on 17-4-2014 at 15:43


CuCl2/Al is disscussed a bit here: http://www.sciencemadness.org/talk/viewthread.php?tid=10249
Also I would just go with Cu+HCl+H2O2 or HCl+CuCO3 for making CuCl2; I find it easier.




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[*] posted on 18-4-2014 at 01:13


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

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!

****
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 :mad:, 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]




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[*] posted on 18-4-2014 at 04:18


ΔH<sup>0</sup><sub>f</sub> for ZnCl2 = - 415.1 kJ/mol
ΔG<sup>0</sup><sub>f</sub> 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]




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[*] posted on 18-4-2014 at 06:24


Wow, beautiful blogfast, I'm very excited about the scope this opens up.

Quote:
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.

361gwd.jpeg - 29kB
Quote:
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!

This thought also occured to me today, that anhydrous ZnCl2 could be a convenient preparative alternative to anhydrous HCl in certain cases.

Using your figures for dG supplied (thanks), I determine a dGrx of -130kJ/(mol Ti) for the preparation of TiCl4! So while this would require constant heating, it should progress readily.

I must say that I'm surprised that the dG is so largly negative when the dH is positive, normally the difference is not so large on endothermic reactions.

Anyhow, this is encouraging news... could I trouble you to try this out as well please. I know titanium chemistry is second hat for you!

Thank you again for the wonderful photos of your AlCl3 preparation.

[Edited on 18-4-2014 by deltaH]




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[*] posted on 18-4-2014 at 06:46


Yes, this is great news! Not a cheap method on the large scale, but much more convenient for small quantities.
I suppose copper chloride and aluminum is too fast for a glass setup, but should be easy in an SS apparatus, making it won't be easy though. I hope to get a welder, which will be a great tool for many applications.
Titanium chloride probably isn't very cheap right? It's the main precursor to Ti metal, but it's hard to make...




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[*] posted on 18-4-2014 at 08:26


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 TiCl<sub>4</sub> from CRC ’86 I get:

ΔH<sup>0</sup><sub>f</sub> = - 804.2 kJ/mol
ΔG<sup>0</sup><sub>f</sub> = - 737.2 kJ/mol
Both for liquid TiCl4.

For 2 ZnCl<sub>2</sub>(s) + Ti(s) →TiCl<sub>4</sub>(l) + 2 Zn(s) that gives me:

ΔH<sup>0</sup><sub>reaction</sub> = + 26 kJ/mol
ΔG<sup>0</sup><sub>reaction</sub> = + 1.6 kJ/mol



[Edited on 18-4-2014 by blogfast25]

[Edited on 18-4-2014 by blogfast25]




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[*] posted on 18-4-2014 at 11:48


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 ;)

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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]




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blogfast25
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[*] posted on 19-4-2014 at 07:36


I tried the ZnCl2/Ti reaction today but no joy. I used the same apparatus as in my previous post but shortened the left hand side of the connecting glass tube a lot.

7.6 g of ZnCl2 anh. and 2.0 g Ti coarse powder (a 50 % excess) where loaded into the reactor and heated to maximum heat (full open air inlet on propane Bunsen). Bubbling started immediately and appeared to stop each time the heating source was removed.

But no fumes of TiCl4 came over and no liquid TiCl4 was found in the condenser.

I assume the bubbling was simply caused by the ZnCl2 boiling (BP = 754 C). The volume of liquid ZnCl2 also didn’t appear to diminish at all.

After the test, I washed out the ZnCl2 with water and subjected the remaining black, coarse powder to cold dilute HCl: no reaction was obtained, a clear sign no elemental zinc was present (Ti only reacts quite slowly with hot, conc. HCl but not dilute).

CuCl2 could be tried for this reaction, as it is much less stable than ZnCl2.

I’m running another ZnCl2/Al reaction tonight.

Edit:

Using a slightly modified apparatus and pre-heating some parts of it I now got a blockage in the condenser tube. This may call for a finger type cold trap. I collected another 0.5 g of product before the tube froze up. I'll be investing in a small piece of glass to get the desired geometry.

The hot AlCl3 also wreaks havoc on the rubber. Ground glass joints at the hot points are really a must.

[Edited on 19-4-2014 by blogfast25]




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deltaH
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[*] posted on 19-4-2014 at 14:14


That's a bummer! I think titanium's passivation layer is the problem here, IMHO. May I suggest mixing some aluminium powder together with your titanium powder and see if it works then? If so, your product would be contaminated with some AlCl3, but not a train smash as it could be distilled if need be or simply tolerated.



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[*] posted on 19-4-2014 at 17:51


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:(.




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[*] posted on 19-4-2014 at 23:43


"Does titanium passivate in chlorine? I thought it only does that in oxygen. "

I was thinking of the oxide passivation layer already present on the surface of the titanium particles, i.e. TiO2. It could be that ZnCl2 fails to etch through it. I was hoping that aluminium powder would help solve this problem by reducing the passivation layer, also with the help of AlCl3 formed.

[Edited on 20-4-2014 by deltaH]




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