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BromicAcid
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I have done electrolysis of lithium chloride in DMSO dozens of times. But you do not get lithium metal that way... you only get ... a surprise.
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Zyklon-A
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Great, it looks like potassium perchlorate, nitrate and iodide are quite soluble.
I don't know, my dad ordered it on his paypal account, so I can't track it.
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BobD1001
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Well, I just ordered 1Kg of Propylene carbonate from Alfa Aesar, hopefully Ill receive it shortly, and get some time to experiment with it!
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deltaH
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BromicAcid, I wasn't suggesting one use DMSO, I was using the data as a comparison to try to infer something with regards to
solubilities of inorganic salts in propylene carbonate. I would guess that if you made any strongly reducing metal in DMSO, you would quickly reduce
it to dimethyl sulfide... a surprise indeed! Let me guess, it stank like &%#@
Zyklonb, I would strongly suggest you abstain from any oxidising anions which would probably react violantly with alkali metals. I'd
say nitrates are definately out!
Halides should be fine and hydroxide are a maybe.
Remember the trick here is that you not only want the electrolyte to be soluble, you simultaneosly want the magnesium salt of that anion to
be poorly soluble. I'd stick to the simple halides and possibly play with hydroxide just out of curiosity.
What you planning for a power supply?
The thought also occured to me that if you manage to dissolve some hydroxide in propylene carbonate, maybe you may want to try an aluminium anode as
well, hypothetically forming sodium aluminate precipitate (if using sodium hydroxide off course). If that works, it would be even cheaper and easier
than magnesium. However, I think halide electrolytes are out with aluminium as aluminium halides are probably very soluble in these kinds of solvents.
[Edited on 5-4-2014 by deltaH]
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Zyklon-A
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| Quote: |
Zyklonb, I would strongly suggest you abstain from any oxidizing anions which would probably react violently with alkali metals. I'd
say nitrates are definitely out! |
I guess, many of my potassium salts are oxidizing, so I thought that would be cheapest.
| Quote: |
What you planning for a power supply? |
Hopefully I will buy this soon.
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Oscilllator
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I really dont think sodium hydroxide is a good idea, because it ALWAYS has some water in it - a few percent at least. Therefore you will probably have
problems forming sodium metal.
@DeltaH - Dont you think using sodium hydroxide and an aluminium anode is a bad idea? Sodium hydroxide reacts spontaneously with aluminium as I'm sure
you know. Or does the reaction require water?
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deltaH
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Nice current capability, but I'm concerned about the single out 5V. I think for Mg anodes this might be overkill, in short, a lot of money to spend on
a gamble. Why not start with a few D cells and experiment, then later you can source the right kind of power supply for your needs with greater
confidence, knowing better how your system behaves?
Oscillator, I thought at first it would, but then wrote out the equation:
Al + NaOH + 3H2O => NaAl(OH)4 + 1.5H2
So yeah, it's water that's actually getting reduced, the sodium hydroxide is just a spectator, but an important one, as without it the aluminium is
kinetically unreactive because of it's passivating oxide layer. In fact, maybe a nice way to dry the solvent with electrolyte in it, provided you have
some means for the hydrogen to escape. (PS. google ALICE rocket when you have some time for a cool extreme of this  )
As for the water in the hydroxide with magnesium, all that should happen is that the cathode will bubble off hydrogen in the beginning, as
blogfast mentioned, until the water is consumed, a bit of a waste of magnesium, but not end of the world.
***
BTW, the E standard calculated from Wiki's table of standard electrode potentials is a mere 20mV for the cell:
Mg(s)|Mg(OH)2(s)||Na+(aq)|Na(s), while this will obviously not be the same for an organic system, it is suggestive that these
kinds of cells will be current beasts, requiring very little work to make large currents flow. In such a situation, all the secondary resistances
become important, e.g. the need for good electrolytic conduction which implies decent solubility and mobility of the ions, also good electrical
conductivity on the external circuit and plates, etc. All these can introduce significant losses in this system because the voltage drop due to the
REDOX electrochemistry per say won't be very large at all.
[Edited on 13-4-2014 by deltaH]
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Zyklon-A
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Yes, the 5 volts will be overkill, but I want that PSU anyway for other projects - specifically for large(ish) scale chlorate production, I could
lower the volts to ~ 3 or so.
I've got plenty of Mg, so no problem there. I also have quite a lot of MMO, so if that's useful for anything....
I found a PDF which says K was isolated from potassium tetrachloroaluminate (KAlCl4), in propylene carbonate. He used either potassium or
aluminum anodes, both of which worked, at least in this particular case.
This Document has lots of information on electro-depositing reactive metals, and is very useful in this discussion.
Here is the PDF:
Attachment: eScholarship UC item 8sz1w229.pdf (1.8MB)
This file has been downloaded 861 times
BTW, it also gives a procedure overview of making Potassium Tetrachloroaluminate, if you're reading this blogfast25 - if you're still interested in making it.
[Edited on 13-4-2014 by Zyklonb]
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deltaH
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Thanks Zyklon for the paper.
Now we know that potassium tetrachloroaluminates works for these kinds of cells and they got decent amounts dissolved into PC (at least 0.5M).
In your context, it could mean that you can enhance the solubility of whatever alkyl halide salt you with to employ in your cell by adding anhydrous
AlCl3.
Chlorine producing cells don't seem like a good idea because of their reported degredation of the solvent, especially in the pressence of water,
however, using a sacrificial anode probably solves this problem as no chlorine would be evolved and one would also be able to run at a low potential.
I'm still uncertain whether PC will tolerate hydroxides dissolved in them. My gut feeling says that the hydroxide ion, being a strong nucleophile, may
attack PC quickly and in the pressence of small amounts of water, forming carbonate salts and ethylene glycol?? But maybe this stops as soon as you
consume the water present?
Anyhow, as things stand, alkyl chlorides, possible with added AlCl3 to enhance solubility, would seem the electrolyte of choice using a magnesium
anode so as to not produce free chlorine and drop the powder requirements of the cell drastically.
[Edited on 13-4-2014 by deltaH]
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Zyklon-A
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Yeah, So when it is deposited on the electrode, how will it attach to it? Will it stick to it, or crumble off?
Either way, I think NaCl will be easiest, although I might try a K salt first. Hydroxides aren't known to work AFAIK, so I won't try that first.
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deltaH
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| Quote: | | Hydroxides aren't known to work AFAIK, so I won't try that first. |
Agreed! According to the article, the
potassium didn't attach on very strongly and did flake off. Who knows what will happen with other metals. But I think it's important that you
establish the solubility of your electrolyte. Doesn't help you use NaCl if it won't dissolve! LiCl may be very soluble though if we infer a similar
behaviour to the DMSO data I dug up. We also now know that anhydrous AlCl3 can help you to dissolve more electrolyte if need be.
[Edited on 13-4-2014 by deltaH]
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blogfast25
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Quote: Originally posted by deltaH  | Thanks Zyklon for the paper.
Now we know that potassium tetrachloroaluminates works for these kinds of cells and they got decent amounts dissolved into PC (at least 0.5M).
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It also means you need to prepare KAlCl<sub>4</sub>, no sinecure at the hobby level.
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Zyklon-A
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Yes, but it does give a procedure to make it, it's not easy, but possible.
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deltaH
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I was thinking of 'simply' co dissolving anhydrous AlCl3 and KCl in reasonably dry propylene carbonate and not so much preparing incredibly
anhydrous KAlCl4 by fusing salts under anhydrous conditions as the authors have done.
In this respect, I'm viewing AlCl3 as a solvant aid when added to the solvent, enhancing the amount of KCl that would otherwise have
dissolved. Am I missing something obvious here?
I know hydrolysis can be a problem for such salts, but surely not when the water present is much less than KCl present?
[Edited on 13-4-2014 by deltaH]
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blogfast25
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| Quote: Originally posted by deltaH | Am I missing something obvious here?
I know hydrolysis can be a problem for such salts, but surely not when the water present is much less than KCl present?
[Edited on 13-4-2014 by deltaH] |
KAlCl<sub>4</sub> is in fact a eutectic mixture:
http://www.crct.polymtl.ca/fact/phase_diagram.php?file=AlCl3...
... with an astonishingly low MP (about 250 C), for a salt. The point at molar fraction AlCl<sub>3</sub> = 0.5 is where the composition
corresponds to KAlCl<sub>4</sub>.
I think your proposed procedure wouldn't work but I'm not sure why. For starters, if it can be done that easily, why bother preparing the
KAlCl<sub>4</sub> separately as they did in that paper?
Water is of course a great enemy of AlCl<sub>3</sub>, which causes it to hydrolyse. Minute amounts of water can never be entirely
avoided, of course. I would certainly go to every bit of reasonable trouble to eliminate water as much as possible.
[Edited on 13-4-2014 by blogfast25]
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Zyklon-A
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| Quote: |
KAlCl4 is in fact a eutectic mixture: |
In that case, it could be made by heating stoichiometric amounts of KCl and AlCl3 until they melt right? That's how they did it in the PDF. The great thing is, iy only has to be bought once. The cell only
consumes KCl, leaving AlCl3 in solution. So, one can just replenish the used KCl, by measuring the chorine produced and adding KCl
stoichiometricly:
| Quote: | Originally posted in said PDF
The process used to study the stability of the solvent with respect to chlorine evolution was the electrolysis of KA1Cl4 in propylene carbonate,
Chlorine was produced and evolved at the anode.
Potassium was reduced or deposited at the cathode. The overall reaction involved
is:
KAlCl4 ↔ K++ Cl- + AlCl3
The half reactions at the cathode and anode are, respectively:
K+ + e- → K, and Cl- + Cl- → Cl2 + 2e-.
Assuming a 100% current efficiency, two moles of potassium were deposited
for every mole of chlorine evolved.This proposed chlorine evolution anodic reaction was quite attractive
from the point of view of raw material considerations. Theoretically only
KCl is consumed in the electrolysis, It can be replenished by simply
adding KCl into the system. Hence this process is capable of producing
valuable chemicals, potassium and chlorine, from a relatively cheap
source, KCl. The most important consideration of all, however, is that
the system can be operated at room temperature.
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This is exactly what I've been looking for!
[Edited on 14-4-2014 by Zyklonb]
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12AX7
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Willfully stepping into the middle of a thread without reading;
Can you get propylene carbonate from lithium batteries?
Nevermind that they already have elemental lithium in them to begin with...
LiPo batteries are normally charged to 4.20V absolute maximum, with the claimed reason that, any higher and lithium metal is deposited, which is a
shorting hazard, bla bla, fire explosion bla bla. Might it be feasible to, say, remove the lithium anode and separator membrane, and use that as the
cathode in a solution with LiCl, and graphite (or something) anode? The added separation between electrodes, of course, helping prevent shorts, and
the added volume providing for more lithium content.
In response to the posts I see above; I'm surprised KAlCl4 doesn't make aluminum!
Tim
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blogfast25
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It would appear so. How many batteries will you have to break into to get a ml of this solvent though? Which then will need purification, of course.
A week ago I extracted some Li from an Energizer AA, it reminded me just how hard and time consuming cracking open these little fortresses actually
is.
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Zyklon-A
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Extracting Li from batteries is a pain it the @ss. I assume that getting PC from batteries is even more of a time waist. But you could try.
This weekend I plan on making anhydrous aluminum (III) chloride, which will be used to make potassium tetrachloroaluminate. I might have to buy a
nafion membrane, which as seen here, is likly going to be rather expensive. Do I really need that membrane? Or will a simple membrane do the job?
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blogfast25
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Not strictly speaking, no. But it does prevent chlorine moving from the anode towards the cathode.
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Zyklon-A
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Hmm, I wonder what membrane is used in lithium batteries. Perhaps it could be used, I don't have any of the remains of the Li batteries that I've
taken apart, but I still have two or three fresh batteries... Seems like big companies can afford to put a good membrane in a battery, considering how
expensive they are in relation to the amount of Li in them.
Here's some more information useful to this discussion:
http://epub.uni-regensburg.de/22953/1/ubr11901_ocr.pdf
You have to pay to view this whole paper, but here's some of it:http://pubs.acs.org/doi/abs/10.1021/j100639a008
Here's something else I found:Potassium Carbonate as a Salt for Deep Eutectic Solvents.
It has to do with Room temperature ionic liquids. I haven't read the whole thing yet.
Here's a short paper on topic: Investigation of the Electrochemical Windows of Aprotic Alkali Metal (Li, Na, K) Salt Solutions
Here's a patent that shows a method for preparing high-purity propylene carbonate and for simultaneously making passivated electrodes.
For $278.00, you can buy four liters of PC
[Edited on 15-4-2014 by Zyklonb]
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blogfast25
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| Quote: Originally posted by Zyklonb | | This weekend I plan on making anhydrous aluminum (III) chloride, which will be used to make potassium tetrachloroaluminate. |
Don't forget to report that here now. I'm planning Al + 3 HCl for the next, next week end.
Interesting links, BTW...
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Zyklon-A
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OK, I plan on doing 2 Al + 3 Cl2 (g), rather than HCl (g). Chlorine is slightly cheaper to make than gaseous HCl, at least for me.
H2SO4 + 2 NaCl → 2 HCl(g) + Na2SO4 is more expensive for me, as I only have < 1L 98% sulfuric
acid which cost me ~ $20, while I have a gallon of con. HCl(aq) which cost me only ~ $6.
4 HCl(aq) + Ca(ClO)2 → CaCl2 + 2 Cl2 + 2 H2O. How are you going to make your HCl(g)?
[Edited on 15-4-2014 by Zyklonb]
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plante1999
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Don't forget to dry the chlorine, sulphuric acid works great for that.
I never asked for this.
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Zyklon-A
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Yes, I might bubble the chlorine through sulfuric acid, or just use a CaCl2 drying tube, (which I already have). If sulfuric acid is used,
it will only be diluted right? With the possibility of dissolved chlorine, which can be gassed off with heating?
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