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agorot
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Potassium from lithium in an organic solvent
I had an interesting idea. Is there an anhydrous solvent that will dissolve potassium chloride (ionize it as well) so that when lithium metal is
added, a single-replacement reaction will occur?
Possibles(?):
-DMSO
-Acetone
-Nitrobenzene
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mewrox99
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How 'bout anhydrous diethyl ether.
KCl is soluble in it.
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Nicodem
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KCl soluble in diethyl ether? Well, I guess soluble or not soluble is a relative term.
I guess you forgot that a solvent needs to be inert under the reaction conditions. All these react very violently to elemental alkali metals!
Try with diglyme and some more soluble potassium salt.
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agorot
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I unfortunately do not have any diglyme. Other options?
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Eclectic
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Propylene Carbonate....but I think Potassium is more reactive than Lithium...it would work as hot molten salt because you could distill off the
potassium vapor...but then you might just as well use activated carbon and K2CO3.
Whimsy: http://www.dilbert.com/2010-08-10/
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agorot
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If you look at a standard reduction potentials table, you will see that lithium is a more powerful reducing agent than potassium.
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Eclectic
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Per gram, maybe, but for something like reaction with water, slow fizz with lithium vs. detonation with cesium...
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blogfast25
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A similar thing's been tried on this forum using a K-salt and magnesium power, based on a 'patent' and without any success.
While Li [0] is probably just about capable of reducing K [+I], this isn't going to happen at room temperature, at least not without a catalyst of
sorts.
At elevated temperature, KCl dissociates a bit: KCl ---> K + 1/2 Cl2. The lithium would then mop up the chlorine, form LiCl and heat and a
self-propagating reaction could be created, depending on the heats of formation of LiCl and KCl. Of course if you could distill the K off, you drive
the equilibrium to LiCl + K.
At RT there isn't enough energy in dissolved KCl +Li to allow reactive collosions, IMHO.
And finding a solvent for KCl that not reactive with lithium metal? Good luck with that...
This is not much more than a pipe dream...
Quote: Originally posted by Eclectic  | | Per gram, maybe, but for something like reaction with water, slow fizz with lithium vs. detonation with cesium... |
You may be relying on the British Science Abuse experiments, which turned out to be a banal hoax and a self-confessed one too: Cs does react
quite violently with water but not explosively.
Edit:
On Brainac's (Science Abuse) fake Cesium/water 'explosions':
http://www.badscience.net/2006/07/brainiac-fake-experiments-...
[Edited on 10-8-2010 by blogfast25]
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agorot
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| Quote: Originally posted by blogfast25 |
While Li [0] is probably just about capable of reducing K [+I], this isn't going to happen at room temperature, at least not without a catalyst of
sorts.
At elevated temperature, KCl dissociates a bit: KCl ---> K + 1/2 Cl2. The lithium would then mop up the chlorine, form LiCl and heat and a
self-propagating reaction could be created, depending on the heats of formation of LiCl and KCl. Of course if you could distill the K off, you drive
the equilibrium to LiCl + K.
At RT there isn't enough energy in dissolved KCl +Li to allow reactive collosions, IMHO.
And finding a solvent for KCl that not reactive with lithium metal? Good luck with that...
This is not much more than a pipe dream...
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Well, I really don't want to try to melt lithium... Liquid lithium scares me. 
Oh well. It was worth a try. Maybe I'll work on a furnace for a KCl electrolysis apparatus.
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blogfast25
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agorot:
Wise words. The dream of making alkali metals at RT is akin to the alchemist dream of developing the 'Elixir of life' or turning lead into gold.
Precisely what can reduce what depends enormously on circumstances: the reduction potentials (in solution) are a poor guide for determining this. For
instance, Al can reduce NaOH (to Na and H2) but Na can reduce AlCl3 (to Al).
Small amounts of potassium can be made quite safely by electrolysing molten KOH, although an argon blanket would help greatly. You can 'fish out' the
floating potassium with a steel gauze. But starting this line of experimentation would be safer with NaOH.
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agorot
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Small amounts of potassium can be made quite safely by electrolysing molten KOH, although an argon blanket would help greatly. You can 'fish out' the
floating potassium with a steel gauze. But starting this line of experimentation would be safer with NaOH.
[/rquote]
KOH is such a strong dissicant though, I'm sure my supply has water in it. I would have to heat it at red heat for a while. KCl doesn't have that
problem.
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densest
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Hot chlorine from electrolysing molten KCl would probably etch every piece of metal in your lab & give you a 6-week cough unless you have a way to
catch it & dispose of it. Been there....
By the time KOH is molten I suspect all the water has been driven off. In any case, the water would be destroyed electrically.
You do have a spatter shield & eye protection? Hydroxide burns to the cornea are very fast and often permanent. A 30cm square piece of glass or
clear plastic on a stand is pretty easy & cheap.
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agorot
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| Quote: Originally posted by densest | Hot chlorine from electrolysing molten KCl would probably etch every piece of metal in your lab & give you a 6-week cough unless you have a way to
catch it & dispose of it. Been there....
By the time KOH is molten I suspect all the water has been driven off. In any case, the water would be destroyed electrically.
You do have a spatter shield & eye protection? Hydroxide burns to the cornea are very fast and often permanent. A 30cm square piece of glass or
clear plastic on a stand is pretty easy & cheap.
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IIRC nickel undergoes passivation with the halogens and is quite robust. Molten hydroxides are pretty intense...I think ill stick with chlorine.
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ScienceSquirrel
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Wikipedia references the preparation of potassium from potassium formate.
http://en.wikipedia.org/wiki/Potassium_formate
I do not have access to the book that they use as a reference and all my web searches have run in to dead ends so far.
Does anyone know anything about this process?
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Nicodem
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The so called formate-potash process is just a process for producing K2CO3 from K2SO4. Potassium formate is an intermediate product obtained after the
first stage where K2SO4, Ca(OH)2 and CO are reacted under high pressure. The potassium formate is then heated in air to give K2CO3 (hence the name
"formate-potash").
Why not trying to electrolyse some potassium salt in diglyme? Obviously a semi-permeable membrane is needed to separate the electrodes, but as long
the salt is at least slightly dissociated in the solvent it should give some potassium at the cathode. Diglyme and other polyglymes are a bit
expensive, but some salts could be soluble enough in THF as well.
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len1
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Since you are not doing it I think we know the answer to that one
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blogfast25
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A bit mean, isn't it, len1?
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ScienceSquirrel
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| Quote: Originally posted by Nicodem | The so called formate-potash process is just a process for producing K2CO3 from K2SO4. Potassium formate is an intermediate product obtained after the
first stage where K2SO4, Ca(OH)2 and CO are reacted under high pressure. The potassium formate is then heated in air to give K2CO3 (hence the name
"formate-potash").
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Thanks for that, it is a reaction I have never heard of.
It might be interesting to see if boiling an aqueous solution of formic acid, potassium sulphate and calcium hydroxide would have the same effect.
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len1
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Maybe. But is it not mean to suggest people try something that you have no experience in. One should have more respect for other peoples time and
people in general, thats my suggestion to nicodem.
If potassium could really be obtained this way dont you think it would be well known by now? A high school experiment? Kolbe did report some
electrogeneration of alkali metals in pyridine, but noone has been able to repeat that
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ScienceSquirrel
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To be honest, I think that it is very unlikely to be possible.
Getting a solution of a potassium salt in a sufficiently unreactive solvent seems like a tough order to me.
If it was possible than it would have been done by now and would almost certainly be noted in advanced text books.
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Nicodem
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| Quote: Originally posted by len1 | | Maybe. But is it not mean to suggest people try something that you have no experience in. One should have more respect for other peoples time and
people in general, thats my suggestion to nicodem. |
Yes, I know, but this is such a contagious behaviour and I learned from the best. You could be proud and I should be ashamed.
| Quote: | | If potassium could really be obtained this way dont you think it would be well known by now? |
This is not a good argument so I don't agree. If it was done it would certainly not be well known. Potassium has no practical use in rechargeable
batteries due to its high molecular weight, so you can not say it would be well known like it is for lithium with its optimally low MW and high redox
potential. For potassium it would be known only to those who are willing to dig in the literature, which is very few (and at the moment I'm not going
to be one of these few).
How about using some more convincing argument. I'm all open for a scientific explanation. Let's suppose it is not possible, which is indeed likely,
but I would like to know why not. So now you are given an excellent opportunity to explain why lithium can be electrolysed in glymes while potassium
supposedly can not be. That would surely be something interesting to learn about (I mean this seriously and without sarcasm!).
| Quote: | | Kolbe did report some electrogeneration of alkali metals in pyridine, but noone has been able to repeat that |
But that does not count as an argument against what I said! Already in my first reply here I explicitly mentioned that even before planing an
experiment a "solvent needs to be inert under the reaction conditions". So pyridine is out of play with potassium metal.
Actually that is not the main problem. Some potassium salts are highly soluble in polyglymes. For example, already potassium iodide, which would at
the same time also be the optimal choice for such an experiment, has good enough solubilities in diglyme and excellent in tri- and tetraglymes
(g/100g):
glyme: 0.5 g; diglyme: 2.5 g; triglyme: 26 g; tetraglyme: 30 g
…there is a human touch of the cultist “believer” in every theorist that he must struggle against as being
unworthy of the scientist. Some of the greatest men of science have publicly repudiated a theory which earlier they hotly defended. In this lies their
scientific temper, not in the scientific defense of the theory. - Weston La Barre (Ghost Dance, 1972)
Read the The ScienceMadness Guidelines!
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watson.fawkes
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| Quote: Originally posted by Nicodem | | Let's suppose it is not possible, which is indeed likely, but I would like to know why not. So now you are given an excellent opportunity to explain
why lithium can be electrolysed in glymes while potassium supposedly can not be. |
Well, one non-candidate for
an explanation is ionization energies, as that of the 2s electron in lithium (520.2 kJ/mol) is significantly more than the 4s one in potassium (418.8
kJ/mol). If anything, there's something to look at with respect to ion size and solvation energies, as the Li(+) ion is significantly smaller than the
K(+) one. Perhaps one of the dimethoxybenzenes would be suitable for potassium if the smaller dimethoxyethane isn't.
Edit: The ionization energy is more for Li, not less. Stupid sign error.
[Edited on 12-8-2010 by watson.fawkes]
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blogfast25
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Completely agree with Nicodem. In fact based on his and watson's input, I'd say there are probably good reasons to try this... And it won't be me
because I haven't any glymes... :-(
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watson.fawkes
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On second thought, I can't tell what I was
thinking, because i got it backwards. The ionization equations are thus:
520 kJ/mol + Li <--> Li(+) + e- (at infinity)
418 kJ/mol + K <--> K(+) + e- (at infinity)
The net equation is thus:
102 kJ/mol + Li(0) + K(+) <--> Li(+) + K(0)
These equation don't take into account either solvation energies or lattice energies, which might tip the balance.
Sorry.
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ScienceSquirrel
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I willl subsidise Nicodem half a litre of any glyme he likes within reason if he agrees to do the experiment.
Not that he will!
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