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blogfast25
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Quote: Originally posted by forgottenpassword  | | That's unfortunate. Perhaps kmno4 was right. I had high hopes after reading your SnCl4 prepation that this would work similarly to give AlCl3. Good
luck with your troubleshooting! |
To be honest, I didn't even see kmno4's contribution until now.
My Olde Holleman states that salts like KAlCl<sub>4</sub> can indeed be obtained from water w/o hydrolysis. Of course just because a book
says so doesn't make it true but so far these substances obtained from watery solutions do redissolve easily in water, confirming what Holleman
claims.
But my purpose here wasn't to prepare AlCl3, analogous to SnCl4.
[Edited on 27-3-2014 by blogfast25]
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Zyklon-A
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Well, could you test your product? If this procedure indeed does work, I would like to try it myself.
Can anybody find the solubility data on this? Or does it react with water?
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blogfast25
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Yes, but it's not that simple, Zb. More results to follow...
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Zyklon-A
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I'll try it anyway, when I get home. I really need to do some more experiments, it's the best way to learn.
I'll post my results too.
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kmno4
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Clearly, you read some false literature.
I recommend "Introduction to Advanced Inorganic Chemistry" by Durrant and Durrant. One of the few really good books of IC.
Google gives appropriate preview from that book:
It is the worst way to learn, the best way is reading good literature.
[Edited on 28-3-2014 by kmno4]
Слава Україні !
Героям слава !
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blogfast25
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| Quote: Originally posted by kmno4 | Clearly, you read some false literature.
I recommend "Introduction to Advanced Inorganic Chemistry" by Durrant and Durrant. One of the few really good books of IC.
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This 'my literature vs. your literature' is a very poor argument and essentially an 'appeal to authority' fallacy.
I believe the salt clearly exists in water up to very high concentrations but getting the last bit of water out may introduce (hydroxyl)aluminates.
It's not a practical way to produce anhydrous chloroaluminates (although I've still got one ace (or Deuce?  ) up my sleeve) for use as ionic liquids.
It's not as simple as you make it out to be.
Experimentation is a great way to learn and see things for yourself. It not an 'either or' situation: you can't be a chemist by being a bookworm only.
[Edited on 28-3-2014 by blogfast25]
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Zyklon-A
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| Quote: | Originally posted by kmno4 ,
It is the worst way to learn, the best way is reading good literature. |
In some ways, yes. For instance, if I do try it, and I get the same product that blogfast25 got, and I put it in a vial and label it Potassium
Tetrachloroaluminate, then really I've learned nothing. I haven't tested it, I just followed some instructions and assumed they
where accurate.
But if I do test it, (I'm not sure how,) then there is a likely potentiality for learning.
[EDIT] fixed quote.
[Edited on 28-3-2014 by Zyklonb]
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blogfast25
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The almost definitive test would be to determine molecular weight. This could be done by quantitatively determining Cl and Al content. Not easy at the
hobby level but not impossible either.
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kmno4
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I can quote at least few articles about AlCl3-KCl-H2O system and about same KAlCl4.
I can write what happens when you add conc. KCl solution to conc. AlCl3 solution... I can, but I am not going to - adequate literature can be found in
the web (of course I do not mean patents).
Zyclonb - that is the point.
Synthesis are the essence of chemistry, but only if you can test (somehow) the product(s).
There are also 'bad' experiments: doing sometning that cannot succed, because experimentator overlooked (or did not consider) some important matters.
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blogfast25
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Please do. I'd be very interested. A cursory search yielded nothing for me.
Tonight I concentrated on the batch prepared with NaCl instead of KCl. It looks different because the supernatant is less viscous and the crystals are
much better formed. Here they are after a few days in the fridge (the yellow tint is due to Fe<sup>3+</sup> contamination:
I decanted off the supernatant into three test tubes and added things to them (sorry it's so out of focus: too close):
Left: with addition of methanol. Nothing happened.
Middle: with addition of NH<sub>3</sub> 33 w%. The tube practically froze over due to so much Al(OH)<sub>3</sub>.
Right: addition of HCl 37 w%. Considerable precipitation resulted. Both the NaCl and AlCl<sub>3</sub> would be forced out of solution by
conc. HCl (the latter as the hexahydrate).
The crystals were then washed firstly with methanol (several times), then acetone (several times), then dried. The product was easily soluble in water
(albeit with minor residual turbidity) and a fairly strong solution was prepared and tested in the same way as the supernatant liquid:
Left: MeOH has no effect.
Middle: ammonia has no effect.
Right (one before last): HCl causes minor precipitation
Right (last tube): this was to test the dry product’s solubility in MeOH, in which it is sparingly soluble.
In particular the second result is revealing: the negative test for Al(OH)<sub>3</sub> with ammonia strongly suggests that all aluminium
was found in the supernatant liquid.
That would mean the crystals are mainly NaCl, as the third and fourth tubes seems to suggest.
[Edited on 28-3-2014 by blogfast25]
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forgottenpassword
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I stumbled upon this by chance, which might be of interest (or not):
ELECTRODEPOSITION OF ALKALI METALS FROM NONAQUEOUS SOLVENTS
Sodium was electrodeposited from a 1.0 molar solu
tion of sodium tetrachloroaluminate (NaAlCl4) in liquid
sulfur dioxide maintained at —15° C. to —25° C.
http://patentimages.storage.googleapis.com/pdfs/US3493433.pd...
[Edited on 4-4-2014 by forgottenpassword]
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blogfast25
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Yep, that is interesting but I'm guessing no one here will be in a hurry to try it!
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