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Author: Subject: lithium and what to do with it
ThatchemistKid
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[*] posted on 15-12-2010 at 19:07


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hkparker
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[*] posted on 15-12-2010 at 19:31


Yea, it sure sinks... I dont know though, something pretty weird would have to be going on for you to be producing potassium

KCl would sink in the oil im pretty sure (depending on the oil) so if theres enough KCl contamination, or Li2O contamination from the burning lithium it could cause it to sink, but that would probably take a lot of salts (again depending on the oil)
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ThatchemistKid
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[*] posted on 15-12-2010 at 20:04


this stuff is silvery white when freshly cut into and is playdough like, I dont know I am not trying to go against all of that thermodynamics, that seems like a huge task, I am just saying what I am seeing.
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[*] posted on 15-12-2010 at 20:06


I believe you that thats what your seeing, but I think there has to be an explanation besides that you made potassium, just based on the thermodynamic data. I dont have access to Li right now, anyone else want to try to reproduce his results?
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blogfast25
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[*] posted on 16-12-2010 at 08:27


@MrChemistKid:

You are continuing to delude yourself and are a classic example of confirmation bias (or ‘belief moves mountains’).

Photos and video are of very low quality, hard to make out anything but your rather funny struggles with lighting a blowtorch and putting on a glove (which later then you ditch and end up burning your finger).

The photo of ‘potassium’ is most likely lithium contaminated with the much heavier KCl (d = 2 g/ml) and that’s why it sinks.

The flame tests you claim to be lilac are mostly off-camera and I can’t see a thing…

Li and K aren’t that easy to distinguish by flame, certainly not in your rather confused conditions.

And since as an appeal to thermodynamics doesn’t seem to work with you (next up: “MrChemistKid’s perpetuum mobile! New! With crappy video!”), try a little common sense. Thanks to lithium batteries small amounts of lithium are now available to most determined home chemists (willing to waste rather a lot of money on new Li batteries), so how come the ‘Nurdrage sphere’ isn’t chocker bloc full of amateurs making K and Na from Li? If time travel worked many of us would be billionaires thanks to it: go back in time a thousand years, put a dollar in a bank, fly back and you’ve got a cool 1.03 to the power 1,000 dollars…

A word of advice: make good photos rather than crappy vids.

Secondly: you want some decent K or Na? Stick some decent electrodes in some molten KOH or NaOH and run a current through it (5 A should do the trick), the metals form quite purely at the cathode and can be scooped off with a gauze wire SS spoon and dunked in mineral oil for later use. Including REAL flame tests.

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blogfast25
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[*] posted on 16-12-2010 at 08:34


Thanks also to Fleaker for his addendum which merely confirms the thermodyn. side of things.
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hkparker
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[*] posted on 16-12-2010 at 09:13


No need to be harsh, he's just trying to understand what's going on, but I absolutly agree with blogfast25 about what's actually happening.
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[*] posted on 16-12-2010 at 09:37


Quote: Originally posted by hkparker  
No need to be harsh, he's just trying to understand what's going on, but I absolutly agree with blogfast25 about what's actually happening.


The problem is that such unsubstantiated reports can lead other naïve or inexperienced experimenters up the garden path for a long time, before they realise they’re chasing a dud. Open minds are to be encouraged but rigorous reasoning, experimentation and observation are the hallmarks of good science. Bad science I leave to l’Oreal and consorts…
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hkparker
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[*] posted on 16-12-2010 at 10:00


That makes sense, I understand
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ThatchemistKid
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[*] posted on 16-12-2010 at 10:17


no, no I definitely agree that I am maybe leading myself on... I am trying to avoid it. But getting yourself out of psychological loops like that is always hard.

Now is there something concrete that I may test that would give a positive for potassium but not for lithium?

I remember a warning to watch out for a red compound that may form on the surface of older potassium, but i definitely do not remember the formation of a red compound any any lithium I have seen.

ALSO. I had to throw that glove off it gave me no control, I would have just melted it into the mass or caught it on fire, so I sacrificed a little pain that I knew was coming, for science :D
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blogfast25
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[*] posted on 16-12-2010 at 10:35


First off, we need to understand that the flame colour (even the spectrum) of potassium and its salts are the SAME. Spectroscopy and flame tests rely on creating small amounts of isolated atoms which then absorb energy in the flame and then release it as spectral light.

So even if you did make, by some new process, potassium IN THE PRESENCE of lithium or one of its salts, a flame test or spectrum would show both and is as a test completely inconclusive.

One way (and quite tricky too) would be to purify the obtained metal. Recrystallising the metal (above MP and under inert solvent) repeatedly would eventually get rid of any non-metal occlusions of Li or K salts (but not really of any remaining Li metal alloyed with the K).

Then you could carry out a flame test, preferably with a decent spectroscope at hand. Relatively weak lines of lilac K should be distinguishable from the carmine red of Li (the lines are opposite sides of the VIS spectrum) but it will depend on apparatus and operator experience with spectra.

Is it possible that your metal contains small amounts of elemental K? Maybe: during rapid cooling of a heated mix of Li and KCl in non-equilibrium conditions small amounts of K may survive. But that’s hardly ‘producing K metal’, now is it?

Here’s a thought: despite having a decent spectroscope, Bunsen burner and plenty KCl, not once have I seen the K lines in a way that I found satisfactory and conclusive. Li on the other hand? Clear as a bell! It then becomes very seductive to confuse the carmine red from Li, possibly with some bright yellow from the irrepressible Na thrown in, with presence of K.



[Edited on 16-12-2010 by blogfast25]
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ThatchemistKid
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[*] posted on 16-12-2010 at 16:53


I do not have a spectrometer handy, well other then the ones at school. But Ill be damned! the metal melts well above the boiling point of water, which is to say it is definitely not potassium. Upon heating a test tube with the metal in it with the blow torch, the metal melted and caught fire. The fire was red in color, the metal also etched the glass of the test tube.

I know that I will from now on check the thermodynamics of a reaction before going about it, but for an experiment that I did years ago before I had even finished Gen chem I do not feel so bad.
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blogfast25
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[*] posted on 17-12-2010 at 12:55


There’s actually an easier method of distinguishing K and Li: by the solubility of their carbonates. Li2CO3 is only very poorly soluble in water, K2CO3 is highly soluble.

Dissolve your metal in water, acidulate slightly and filter. Add slowly concentrated Na2CO3 to the solution till fizzing stops, then add some more: a precipitate indicates Li rather than K.

No experiment is ever really a failure: yours simply confirmed the validity of chemical thermodynamics! :)


[Edited on 17-12-2010 by blogfast25]
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[*] posted on 17-12-2010 at 19:28


In the lab we did a "sodium fusion test" to find out if an unknown had nitrogen, sulphur or halide atoms. You just heat the unknown with some sodium until the organic decomposes and any heteroatoms bond to sodium forming Na salts. I wonder if you can you do the same thing with lithium.
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[*] posted on 17-12-2010 at 23:09


has anyone here reacted lithium metal with alkyl halides at home ?

this is something I wish to look into through the next year.




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hkparker
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[*] posted on 17-12-2010 at 23:19


Try posting this in organic
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[*] posted on 17-12-2010 at 23:28


ye fair enough but its just a question and this thread is what to do with lithium.

ill end up doing the tests in the end anyway hopefully I dont blow my self up.

much reading to do on this.


post 303 hehe like a shot from a gun.



[Edited on 18-12-2010 by Ephoton]




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[*] posted on 17-12-2010 at 23:32


Hehe good luck. I've always found organometallic chemistry fascinating
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[*] posted on 18-12-2010 at 01:33


OK, here we go. I have almost a kilo of lithium metal I don't have any use for. The metal is in lump form (rods and blocks), but I'm guessing I'll be able to reproduce some of the results.

I'm planning on using a metal box, closed, with lithium pieces (3-5 mm cubes) and the following:
- NaCl
- KCl
- CaCl2 anhydrous (I have some but I'll keep it melted for an hour to avoid possible Ca looses)

I read several times Li is a good reducer, and a Li fire is the absolute worst nightmare as it will burn it's way thru things such as concrete or azbestos. Only copper powder is supposed to be a good thing to put off a Li fire.



Experiments screduled for next week, probably Wednesday. I'll report back the results.

BTW, Li metal ignited burns much like Mg (white flame); so does Sr and Ba. Only the ions will give the characteristic spectral colors, which is not true for K and Na (the pure metals burn with their corespondant colors).
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[*] posted on 18-12-2010 at 01:36


I once burnt lithium when it was held by a steel retort stand and clamp.

it melted the clamp :D and the window it was a foot a way from :o




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[*] posted on 18-12-2010 at 05:59


Last year, I prepared some LiOH by dropping very small chunks of lithium from Li batteries in a beaker of distilled water. Little did I know that I had to vigorously and constantly stir the lithium in until dissolved to prevent it from catching on fire... Actually I knew it would ignite, but I didn't know what would happen next! :mad:

My very first chunk was maybe a bit too big and I didn't stir, so after 2 or 3 seconds it did catch on fire and emitted a beautiful magenta flame... but then the chunk of molten metal slowly made its way to the side of the beaker, and upon contact, the still flaming Lithium proceeded to melt an nice hole on the side of the pyrex beaker, which cracked in several places. Thankfully it stayed together and didn't spill. Aargh! It was my only good 500 ml beaker!

So from that point on, I used an inexpensive mason jar and dropped smaller chunks of Li while stirring the contents of the jar. I dropped Li until the last bits took an eternity to dissolve, and at that point, I had a fairly concentrated solution of LiOH.

But retrieving Li from batteries is fun stuff, you just have to do this outside or in a fume hood because some gases are emitted when you expose the batterie's guts to the air. Anyone knows if that Lithium metal is high purity?

Robert
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[*] posted on 18-12-2010 at 06:55


Quote: Originally posted by a_bab  
OK, here we go. I have almost a kilo of lithium metal I don't have any use for. The metal is in lump form (rods and blocks), but I'm guessing I'll be able to reproduce some of the results.

I'm planning on using a metal box, closed, with lithium pieces (3-5 mm cubes) and the following:
- NaCl
- KCl
- CaCl2 anhydrous (I have some but I'll keep it melted for an hour to avoid possible Ca looses)

I read several times Li is a good reducer, and a Li fire is the absolute worst nightmare as it will burn it's way thru things such as concrete or azbestos. Only copper powder is supposed to be a good thing to put off a Li fire.



Experiments screduled for next week, probably Wednesday. I'll report back the results.

BTW, Li metal ignited burns much like Mg (white flame); so does Sr and Ba. Only the ions will give the characteristic spectral colors, which is not true for K and Na (the pure metals burn with their corespondant colors).


This gets sillier and sillier. A_bab, READ THIS THREAD: reductions of NaCl, KCl and CaCl2 with lithium ARE NOT POSSIBLE, they’re thermodynamically VERBOTEN. As outlined not only in great detail by me but also proved by the fact that what MrChemistKid believed to be potassium metal turned out not to be such, as firmly (and easily) predicted by thermodynamical theory.

Why waste an expensive and dangerous chemical like lithium to prove a negative????

On top of that, Li is very capable of reducing compounds (in particular fluorides) but due to its affinity for oxygen, these reductions are best carried out in the absence of air or else Li + ¼ O2 --- > ½ Li2O. It doesn’t take much oxygen to oxidise quite a lot of lithium.

Also, if you wanted to carry out a successful reduction with Li it will by its very nature be highly exothermic: a ‘metal box’ as crucible won’t do.


[Edited on 18-12-2010 by blogfast25]
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[*] posted on 18-12-2010 at 08:51


I belive it is very pure mr dent I think this was discussed by some of the people using
birch reductions.

though I am not sure of the impurities you would get from exposure to air when
extracting it from the battery or from the other parts of the battery contacting it.




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[*] posted on 18-12-2010 at 13:14


Are you guys sure about the reaction between NaOH and Mg? The equation used to "prove" the thermodynamics seems rather incomplete.

The correct equation would be:

2NaOH + 2 Mg --> 2Na + 2MgO + H2

and essentially be a combination of the following 3 reactions:

2NaOH --> Na2O + H2O
Mg + H2O --> MgO + H2
Na2O + Mg --> 2Na + MgO

Also there seems to be NO mention of the fact that a reaction that is thermodynamically unfavourable cán proceed at elevated temperatures. After all, delta(G) = delta(H) - T(delta)S

delta(H) = 349.08 kJ/mole and delta(S) = 76.5 J/K

If delta(G) is calculated for room temperature it is 326.7 kJ/mole...
It does not become negative untill a temperature of 4365 K is reached. Evidently a different mechanism is driving this reaction.

[Edited on 18-12-2010 by Nerro]




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[*] posted on 18-12-2010 at 13:36


It must be, because that reaction has worked quite well for me in the past
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