Sciencemadness Discussion Board - Tetrachloroaluminates: ionic solvents?

Tetrachloroaluminates: ionic solvents?

blogfast25 - 24-3-2014 at 13:02

I just prepared some KAlCl4 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]

DraconicAcid - 24-3-2014 at 13:08

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.

Töilet Plünger - 24-3-2014 at 13:13

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?

DraconicAcid - 24-3-2014 at 13:15

Quote: Originally posted by Töilet Plünger

Can hexachloroaluminates and hexabromoaluminates exist as well?

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.

forgottenpassword - 24-3-2014 at 13:29

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.

blogfast25 - 24-3-2014 at 13:47

Quote: Originally posted by forgottenpassword

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.

Yes, it would be fairly easy way to prepare AlCl3 (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]

copperastic - 24-3-2014 at 13:55

How did you make it?

blogfast25 - 24-3-2014 at 14:02

@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) === > KAlCl4(aq) + 3/2 H2(aq)

[Edited on 24-3-2014 by blogfast25]

Metacelsus - 24-3-2014 at 18:02

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]

Texium - 24-3-2014 at 18:51

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!

Brain&Force - 24-3-2014 at 19:16

I wonder if adding aluminum chloride to sodium bromide will cause the melting point to be lowered further if a tetrabromoaluminate is not stable.

blogfast25 - 25-3-2014 at 05:20

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.

The really interesting further lowering of the MP is in eutectic points, e.g. with KCl. Know what I'm saying?

Brain&Force - 25-3-2014 at 05:47

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.

kmno4 - 25-3-2014 at 07:23

Quote: Originally posted by blogfast25

Al(s) + 3 HCl(aq) + KCl(aq) === > KAlCl4(aq) + 3/2 H2(aq)

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.

Bezaleel - 25-3-2014 at 07:44

Quote: Originally posted by blogfast25

The really interesting further lowering of the MP is in eutectic points, e.g. with KCl. Know what I'm saying?

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?

blogfast25 - 25-3-2014 at 11:39

@Bezaleel:

Yes, these salts don’t form hydrates, as far as I know.

I finished off the first batch of KAlCl4 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 NaAlCl4, which is boiling in on the right of the first photo. I expect its solubility behaviour to be different from KAlCl4.

According to this Google patent: http://www.google.st/patents/US6482381 the 70/30 (molar) NaAlCl4/ KAlCl4 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]

forgottenpassword - 25-3-2014 at 11:45

Have you confirmed that it has a low melting point?

DraconicAcid - 25-3-2014 at 12:07

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.

blogfast25 - 25-3-2014 at 13:38

Quote: Originally posted by DraconicAcid

The reduction of the tetrachloroaluminate to give aluminum metal, for one thing.

At the anode something has to be oxidised. In the presence of free chloride anions I would have thought that would be:

Cl- === > 1/2 Cl2 + e

Or perhaps AlCl4- === > AlCl3 + 1/2 Cl2 + e

I just don't see how Al could be reduced from such a solution.

If there is dissociation of the AlCl3 then reduction could occur at the cathode but it requires higher voltage than Na+, I think...

Quote: Originally posted by forgottenpassword

Have you confirmed that it has a low melting point?

Not yet. All in good time.

[Edited on 25-3-2014 by blogfast25]

DraconicAcid - 25-3-2014 at 13:49

Quote: Originally posted by blogfast25

Quote: Originally posted by DraconicAcid

The reduction of the tetrachloroaluminate to give aluminum metal, for one thing.

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]

forgottenpassword - 25-3-2014 at 13:53

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.

blogfast25 - 25-3-2014 at 14:06

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]

forgottenpassword - 27-3-2014 at 13:04

How did the melting points turn out?

blogfast25 - 27-3-2014 at 13:14

Quote: Originally posted by forgottenpassword

How did the melting points turn out?

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 AlCl3 in my product is too low (just below 0.5) . Acc. to the phase diagram for AlCl3/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 AlCl3 just below 0.5.

I've just finished a batch with an AlCl3 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]

forgottenpassword - 27-3-2014 at 13:27

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!

blogfast25 - 27-3-2014 at 13:51

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 KAlCl4 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]

Zyklon-A - 27-3-2014 at 14:04

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?

blogfast25 - 27-3-2014 at 14:07

Quote: Originally posted by Zyklonb

Well, could you test your product?

Yes, but it's not that simple, Zb. More results to follow...

Zyklon-A - 27-3-2014 at 14:25

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.

kmno4 - 28-3-2014 at 00:25

Quote: Originally posted by blogfast25

My Olde Holleman states that salts like KAlCl4 can indeed be obtained from water w/o hydrolysis

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:

books.bmp - 7kB

Quote: Originally posted by Zyklonb

I really need to do some more experiments, it's the best way to learn.

It is the worst way to learn, the best way is reading good literature.

[Edited on 28-3-2014 by kmno4]

blogfast25 - 28-3-2014 at 05:26

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.

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]

Zyklon-A - 28-3-2014 at 05:39

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]

blogfast25 - 28-3-2014 at 05:54

Quote: Originally posted by Zyklonb

[But if I do test it, (I'm not sure how,) then there is a likely potentiality for learning.

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.

kmno4 - 28-3-2014 at 08:58

Quote: Originally posted by blogfast25

This 'my literature vs. your literature' is a very poor argument and essentially an 'appeal to authority' fallacy.

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.

blogfast25 - 28-3-2014 at 13:41

Quote: Originally posted by kmno4

I can quote at least few articles about AlCl3-KCl-H2O system and about same KAlCl4.

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 Fe3+ 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 NH3 33 w%. The tube practically froze over due to so much Al(OH)3.

Right: addition of HCl 37 w%. Considerable precipitation resulted. Both the NaCl and AlCl3 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)3 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]

forgottenpassword - 4-4-2014 at 09:45

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]

blogfast25 - 4-4-2014 at 09:48

Quote: Originally posted by forgottenpassword

I stumbled upon this by chance, which might be of interest:

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...

Yep, that is interesting but I'm guessing no one here will be in a hurry to try it!