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blogfast25
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deltaH: I doubt that passivation is the problem here. My Ti dissolves easily in hot, concentrated H2SO4 and hot, 37 w% HCl. I don't really see what a
bit of Al would do but I might try it. Sure, Al does reduce TiO2 very slowly.
There is of course already a lab bench preparation of TiCl4: TiO2 + NaCl + potassium persulphate (see 'plante').
I think I'll test my AlCl3 as a Friedel Crafts catalysts, need to get some 'bits' together for that. I will also try and scale this up a bit with a
more suitable apparatus.
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deltaH
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Thanks blogfast for indulging me.
To elaborate my thinking: since ZnCl2 behaves as a mild HCl of sorts by my analogy, it is insufficiently strong an etchant to dissolve the
passivisation layer of titanium in the same way titanium does not react with dilute HCl. The purpose of adding some Al is twofold, as a reducing agent
and also the anhydrous AlCl3 generated, being a strong lewis acid, could help dissolve the passivation layer.
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]
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blogfast25
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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 
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blogfast25
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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]
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Zyklon-A
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Ah, very simple procedure! I like it. This way one can just keep the condensed AlCl3 in the reagent bottle - Although it's much bigger
than needed. Or perhaps the reaction could be done many times without emptying the bottle, until it's full, and can be stored as is.
Either way, this is about as easy as it gets to make anhydrous aluminum (III) chloride. No chlorine, no gas generators, no tubes or SS pipes,
condensing is so easy.
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blogfast25
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Zb:
I'd still like to try the chlorination with HCl with this set up. But I'd refrain from chlorine, if I was you. The HCl reaction, according to Holleman
('Inorganic Chemistry'), is self-sustaining once initiated. The Cl2 reaction will by definition be much more exothermic, thus harder to control I
think.
Brauer (Preparative Inorganic Chemistry, see library) gives a prep for Al + HCl but not for Al + Cl2, for instance...
[Edited on 4-5-2014 by blogfast25]
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deltaH
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Excellent work blogfast, I'm happy you got it to work!
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blogfast25
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Thanks. Can't believe I was so stupid to try this first with fairly narrow tubes and not expect blockages. That would only have been possible by
keeping them at 190 C or so. Hindsight is 20/20, I suppose...
I need to add that this set up can also be used to re-sublime the AlCl3, for instance to remove any ZnCl2 impurity that may have come over too.
[Edited on 5-5-2014 by blogfast25]
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deltaH
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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]
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blogfast25
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I know.
Yields could be improved fairly easily, IMHO. It's a solid-solid reaction, that explains low yields, I think.
[Edited on 19-2-2016 by blogfast25]
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JJay
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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.
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blogfast25
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| 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. |
Ethyl acetate won't work, I think. It's BP is only 77 C. The solid-solid reaction needs strong heating to initiate and the reaction front exceeds 180
C (otherwise the AlCl3 could not sublime off). Ethyl acetate is an aprotic solvent.
For Al + anh. HCl look in the library's Georg Brauer 'Hbk. of Preparative Inorganic Chemistry', it has a procedure + apparatus. Done correctly it will
be near-quantitative.
[Edited on 20-2-2016 by blogfast25]
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JJay
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Right, but in water, the reaction proceeds at room temperature.
IIRC, Brauer suggested using a "parting flask" at temperatures that looked dangerous for borosilicate glass.
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blogfast25
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Water is a very 'special' solvent, though. If the ZnCl2 fully dissociates in ethyl acetate, as it does in water, then the reaction may be
viable. Otherwise you'll have ZnCl2 particles sloshing against Al particles and no party.
Re. Al + anh. HCl, acc. my old A.F. Holleman ('Inorganic Chemistry'), after initially heating the Al, once the reaction starts it is self-sustaining.
That sounds plausible, as for:
Al(s) + 3 HCl(g) ===> AlCl3(g) + 3/2 H2(g)
... I estimate an Enthalpy of Reaction of about - 390 kJ/mol of AlCl3(g) (adjusted for the vapourisation Enthalpy of AlCl3(g)).
Compare it with the ZnCl2/Al reaction: only about - 75 kJ/mol.
So, quite a lot of heat!
[Edited on 20-2-2016 by blogfast25]
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JJay
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Oh for zinc chloride and aluminum? I just got some media bottles and a thermometerless distillation adapter in the mail today.
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blogfast25
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| 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. |
ChemPlayer's version was basically a simple distillation apparatus. Call it a 'dry distillation'.
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