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blogfast25
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Tetrachloroaluminates: ionic solvents?
I just prepared some KAlCl<sub>4</sub> via wet route for some project and it was remarkably easy to do. The salt is now drying.
Searching for some properties I came across a MP of 245 C, which remarkably low for a salt. For the equivalent Na salt 154 C is listed and for the Li
salt 142 C!
With such low melting points and great thermal stability these sound like almost ideal low temperature ionic solvents. Has anyone here considered or
actually used them as such?
[Edited on 24-3-2014 by blogfast25]
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DraconicAcid
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I've considered them, but never actually used them. A great deal of work has been done with the tetrachloroaluminates of organic cations, to make the
melting points even lower, but these tend to be water-sensitive.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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Töilet Plünger
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What about the tetrabromoaluminates? I don't even know if they exist, but I suspect they may have lower melting points.
Can hexachloroaluminates and hexabromoaluminates exist as well?
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DraconicAcid
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I don't think so- the chloride ion is too large, and the aluminum ion is too small. The hexafluoroaluminate ion is stable, though; molten cryolite is
a standard high-temperature solvent for alumina.
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forgottenpassword
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Going by nothing more than the entry for melting point here: http://en.wikipedia.org/wiki/Lithium_tetrachloroaluminate the salts may decompose at the melting point. I'm sure it would be easy enough to find
out or try out. Perhaps they decompose releasing AlCl3 - which would be a useful result nonetheless.
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blogfast25
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Yes, it would be fairly easy way to prepare AlCl<sub>3</sub> (anh.) But for the potassium salt some quite high temperature (> 500 C)
applications have been cited (patents mainly).
I'll fire mine up tomorrow.
@TP: the bromine ion is likely to be too large for four to fit around an Al atom, too much electron repulsion.
[Edited on 24-3-2014 by blogfast25]
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copperastic
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How did you make it?
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blogfast25
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@coppertastic:
One pot method: 0.1 mol clean Al scrap + 20 % excess of HCl 37 % + stoichiometric amount of KCl. Heat slowly, reaction is highly exothermic when it
starts. To dissolve the last stubborn bits of Al I had to add a bit more HCl. Then simmered it down to about half the initial volume. Crystals
immediately appeared when I stopped heating. Cool, then chill and collect. Dry at 200 C for about 1 h - 2 h.
Al(s) + 3 HCl(aq) + KCl(aq) === > KAlCl<sub>4</sub>(aq) + 3/2 H<sub>2</sub>(aq)
[Edited on 24-3-2014 by blogfast25]
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Metacelsus
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I may try the equivalent reaction with sodium chloride this week. It's nice that the reactants are so cheap.
[Edited on 25-3-2014 by Cheddite Cheese]
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Texium
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Sounds interesting… something to add to the list of the many things I need to try soon. Nice to see something which I actually have all of the
reactants for in expendable amounts!
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Brain&Force
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I wonder if adding aluminum chloride to sodium bromide will cause the melting point to be lowered further if a tetrabromoaluminate is not stable.
At the end of the day, simulating atoms doesn't beat working with the real things...
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blogfast25
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Quote: Originally posted by Brain&Force  | | I wonder if adding aluminum chloride to sodium bromide will cause the melting point to be lowered further if a tetrabromoaluminate is not stable.
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The really interesting further lowering of the MP is in eutectic points, e.g. with KCl. Know what I'm saying?
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Brain&Force
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Uhh...not anymore.
I thought they were complexes, not eutectics. But I'm sure at this point I have no idea what I'm talking about.
At the end of the day, simulating atoms doesn't beat working with the real things...
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kmno4
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Do you really think that can get any tetrachloroaluminate in this way ????
Even traces of water give oxygen containing anions, you can forget about KAlCl4 made in water.
Слава Україні !
Героям слава !
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Bezaleel
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Correct. Would you also consider eutectics with LiCl or NaCl added to your KAlCl4 or KxAlCl3+x?
Your KAlCl4 crystals did not contain any crystal water, am I right?
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blogfast25
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@Bezaleel:
Yes, these salts don’t form hydrates, as far as I know.
I finished off the first batch of KAlCl<sub>4</sub> today by Buchnering off the supernatant liquid, which wasn’t easy because it was
viscous. I then washed the crystalline material with several aliquots of acetone and sucked them dry. These were then collected and dried at
progressively higher temperatures. Below left is the snow white salt drying:
Something strange happened though: the mixture of supernatant and acetone formed a two phase system in the vacuum flask:
I assume the ionic strength of the supernatant is too high to still be miscible with acetone.
I also started a batch of NaAlCl<sub>4</sub>, which is boiling in on the right of the first photo. I expect its solubility behaviour to be
different from KAlCl<sub>4</sub>.
According to this Google patent: http://www.google.st/patents/US6482381 the 70/30 (molar) NaAlCl<sub>4</sub>/ KAlCl<sub>4</sub> eutectic has an MP of 125
Celsius.
Now call me what you want (a dreamer?) but what would a priori be an obstacle to mixing in say 0.1 or 0.2 mol of NaCl (per mol of eutectic)
and electrolysing this melt at the right voltage to obtain liquid sodium metal?
[Edited on 25-3-2014 by blogfast25]
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forgottenpassword
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Have you confirmed that it has a low melting point?
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DraconicAcid
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| Quote: Originally posted by blogfast25 |
Now call me what you want (a dreamer?) but what would a priori be an obstacle to mixing in say 0.1 or 0.2 mol of NaCl (per mol of eutectic) and
electrolysing this melt at the right voltage to obtain liquid sodium metal? |
The reduction of the tetrachloroaluminate to give aluminum metal, for one thing.
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Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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blogfast25
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At the anode something has to be oxidised. In the presence of free chloride anions I would have thought that would be:
Cl<sup>-</sup> === > 1/2 Cl<sub>2</sub> + e
Or perhaps AlCl<sub>4</sub><sup>-</sup> === > AlCl<sub>3</sub> + 1/2 Cl<sub>2</sub> + e
I just don't see how Al could be reduced from such a solution.
If there is dissociation of the AlCl<sub>3</sub> then reduction could occur at the cathode but it requires higher voltage than
Na<sup>+</sup>, I think...
Not yet. All in good time.
[Edited on 25-3-2014 by blogfast25]
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DraconicAcid
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| Quote: Originally posted by blogfast25 |
At the anode something has to be oxidised. In the presence of free chloride anions I would have thought that would be:
Cl<sup>-</sup> === > 1/2 Cl<sub>2</sub> + e
Or perhaps AlCl<sub>4</sub><sup>-</sup> === > AlCl<sub>3</sub> + 1/2 Cl<sub>2</sub> + e
I just don't see how Al could be reduced from such a solution.
If there is dissociation of the AlCl<sub>3</sub> then reduction could occur at the cathode but it requires higher voltage than
Na<sup>+</sup>, I think... |
What's getting oxidized at the anode isn't really relevant- it's the reduction at the cathode that we're talking about. What's going to be easier to
reduce- the aluminum ions or the sodium ions? My money's on the aluminum, despite the coordination by chloride ions.
Citation: http://books.google.ca/books?id=0XmiTX6e_9YC&lpg=PA566&a...
[Edited on 25-3-2014 by DraconicAcid]
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forgottenpassword
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I think it would be worth checking out, myself. As kmno4 suggests, it is unusual to prepare aluminium compounds in water without water binding to the
aluminium. Simply heating a small amount in a test tube with a lighter would give you an easy way to determine a low melting point.
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blogfast25
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Actually DA, I'm gonna put my dunce cap on and stand in the corner for 15 minutes.
Everything points to Al being reduced at the cathode (not Na) and there are plenty of references to it, albeit using more fanciful, organic (even
lower melting) chloroaluminates.
It'd be worth a simple experiment in any case, this electrodeposition of Al at low temperature.
Forgotten: on my to do list for tomorrow. These compounds testify to the power of complexation.
[Edited on 25-3-2014 by blogfast25]
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forgottenpassword
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How did the melting points turn out?
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blogfast25
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I didn't get it to melt at all (propane Bunsen).
I've been running quite a few experiments to get to the bottom of this.
My tentative conclusion so far is that the molar fraction of AlCl<sub>3</sub> in my product is too low (just below 0.5) . Acc. to the
phase diagram for AlCl<sub>3</sub>/KCl, here:
http://www.crct.polymtl.ca/fact/phase_diagram.php?file=AlCl3...
... that would mean I'm in the steep part of the melt curve for molar fractions of AlCl<sub>3</sub> just below 0.5.
I've just finished a batch with an AlCl<sub>3</sub> target molar fraction of 0.65. That would be in the 'safe' zone, with MP below 200 C.
One thing is for sure: KAlCl4 can be obtained by boiling in a aqueous solution of it, without the substance undergoing hydrolysis: all samples I've
taken so far remain soluble in water.
Towards the end the liquid starts getting more and more viscous like syrup and somewhat hazy. Then a white material starts to build up.
[Edited on 27-3-2014 by blogfast25]
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forgottenpassword
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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!
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