Sciencemadness Discussion Board - lithium and what to do with it

lithium and what to do with it

d010060002 - 14-12-2010 at 18:48

I recently extracted some lithium from some AA energizer ultimate lithium batteries (there is a youtube of it from nurdrage. It's pretty good, except to get the cap off without shorting it too much I had to use a dremel). I was looking at one of my periodic tables, and below it there was a reactivity series for metals. What I found unusual was that Lithium was listed as the most reactive. I was wondering if:
A) this is true
B) if it is, can I use the lithium for a single replacement reaction for something like sodium?

crazedguy - 14-12-2010 at 19:18

Lithium is not the most reactive.

starch - 14-12-2010 at 20:12

lithium
Sodium
Potassium

there ya 3 main alkali metals, in order of reactivity as seen on periodic table, lithium being the least reactive

you can use the lithium for a single replacement reactions do some googling there is heaps of info out there, some reactions might call for a more reactive alkali, ie sodium or potassium

also have a look on youtube for the BBC chemistry volitile history video here is the link for the first part, its a great watch and will teach ya some things

http://www.youtube.com/watch?v=25lprEvoFJ8

have fun be safe

ThatchemistKid - 14-12-2010 at 20:43

Lithium is not the most reactive, in the sense that it does not react the fastest or most vigorously.
What is being referred to in that activity series of metals is the reduction potential of the metals. Lithium has the highest reduction potential (around 3 volts if I remember correctly)
Also it is hold its electron much more tightly than the other alkali metals so the amount of energy released when it reacts and donates its electron ( say if it were reacting with water) is actually higher then the amount of energy that cesium releases, but since its electron is held so tightly, it is just released in more controlled less vigorous manner.

hkparker - 14-12-2010 at 21:49

Thatchemistkid is right, it reacts the least vigorously of the alkali metals, due the lack of shielding effect of the electrons and the less distance between the nucleus and the valence electrons, but it is the most powerful reducing agent (with a reduction potential of around thee volts).

So in theory you could use it to displace sodium, but im not sure what kid of set-up or reaction conditions you would need. Molten lithium is practically pyrophoric in my experience, so the setup would be rather complex.

ThatchemistKid - 14-12-2010 at 22:05

I have previously, back a couple of years ago, been able to thermite sodium, potassium and calcium out of their respective chlorides ( with calcium actually i used the hypochlorite, it worked wonders I would end up with a marble sized, spherical ball of calcium) using lithium metal foil from batteries. I would just take some LI foil pack some potassium chloride in there until it looked sort of like a cigar. Then i would light it on fire with a torch and as it reacted I would let the molten potassium drip into a jar filled with mineral oil, If i was lucky I would end up with a gram or two of silvery white potassium chunks (after cleaning off the potassium that had dripped into the jar). I also found that inside the remains of the lithium cigar there would be small chunks of potassium that could be removed and cleaned.

but you should take care and only do small quantities at once, usually my cigars of lithium were an inch and a half long and filled with 2 gramsish of the salt.

wait it wasn't two years ago it was three, I was a freshman in college O_O wow where has the time gone.

[Edited on 15-12-2010 by ThatchemistKid]

blogfast25 - 15-12-2010 at 09:44

Quote: Originally posted by ThatchemistKid

I have previously, back a couple of years ago, been able to thermite sodium, potassium and calcium out of their respective chlorides ( with calcium actually i used the hypochlorite, it worked wonders I would end up with a marble sized, spherical ball of calcium) using lithium metal foil from batteries. I would just take some LI foil pack some potassium chloride in there until it looked sort of like a cigar. Then i would light it on fire with a torch and as it reacted I would let the molten potassium drip into a jar filled with mineral oil, If i was lucky I would end up with a gram or two of silvery white potassium chunks (after cleaning off the potassium that had dripped into the jar). I also found that inside the remains of the lithium cigar there would be small chunks of potassium that could be removed and cleaned.

but you should take care and only do small quantities at once, usually my cigars of lithium were an inch and a half long and filled with 2 gramsish of the salt.

wait it wasn't two years ago it was three, I was a freshman in college O_O wow where has the time gone.

[Edited on 15-12-2010 by ThatchemistKid]

You write in the conditional a lot, yet most of the reactions you describe are thermodynamically IMPOSSIBLE.

Here are NIST values of the Heat of Formation at 298 K for:

LiCl = - 408 kJ/mol
NaCl = - 411 kJ/mol
KCl = - 437 kJ/mol
CaCl2 = - 796 kJ/mol

For KCl + Li --- > K + LiCl the heat of reaction would be 437 – 408 = + 29 kJ/mol. Not possible thermodynamically, except in conditions where the K is distilled off.

For NaCl + Li --- > Na + LiCl the heat of reaction would be 411 – 408 = + 3 kJ/mol. Not possible thermodynamically, except in conditions where the Na is distilled off.

For CaCl2 + 2 Li --- > Ca + 2 LiCl the heat of reaction would be 796 – 2 x 408 = - 20 kJ/mol. Only JUST barely thermodynamically possible.

I think you’re more ThatBaloneyKid than ThatChemistKid, frankly…

From the wiki entry on Li:

“On a commercial scale, lithium is isolated electrolytically from a mixture of lithium chloride and potassium chloride.”

So there you have liquid Li metal in contact with liquid KCl: no reaction. Note also how the LiCl is reduced but not the KCl.

You were melting Li metal, nothing more.

[Edited on 15-12-2010 by blogfast25]

ThatchemistKid - 15-12-2010 at 11:24

No, I am quite certain I was getting chunks of potassium and sodium, Actually I have pictures stored somewhere, but I am not sure if I have a video of me doing this, the metal produced was definitely Na, K, Ca. Also it is possible to thermite sodium out of NaOH using magnesium powder in a method akin to this one.

there is a youtube video of this... here let me find it.

http://www.youtube.com/watch?v=908rjHQ5mmc

the safety on this is a little ehh but w.e.

honestly I can not find the pictures I have pictures of the sodium and potassium once it was produced and is in a jar but I can not find a picture of it while it was still in the process. Uhhmm I suggest that you try this yourself before pointing fingers etc etc...

maybe we can throw around somemore thermodynamics with the heat of formation of MgO or whatever.
But i have taken the chunks of metal that have formed from the reaction and dropped them in warm water and have had the characteristic yellow and purple (lilac) flames.

" 'On a commercial scale, lithium is isolated electrolytically from a mixture of lithium chloride and potassium chloride.'

So there you have liquid Li metal in contact with liquid KCl: no reaction. Note also how the LiCl is reduced but not the KCl.

You were melting Li metal, nothing more."

that is a little interesting though considering that lithium has a higher reduction potential than potassium.

[Edited on 15-12-2010 by ThatchemistKid]

ThatchemistKid - 15-12-2010 at 11:30

In fact there is a funny story to go along with this actually, I went back to visit my old highschool one day and brought along a chunk of potassium that I had made in this manner. As I was walking through the halls I met my old biology teacher, who was familiar with my antiks. So i had in hand ready a cup of water and also I had a chunk of potassium that I had made. Now my teacher was standing talking to another teacher I had never met before, so I assumed she was a substitute. I then dropped the potassium into the cup of water it fizzled then suddely popped with a lilac flame produced. I then just turned around and walked off. It turns out though, that the woman that my teacher was talking to was the dean of science for the district.. so apparently my little stunt caused a bit of a stir

blogfast25 - 15-12-2010 at 12:52

Quote: Originally posted by ThatchemistKid

No, I am quite certain I was getting chunks of potassium and sodium, Actually I have pictures stored somewhere, but I am not sure if I have a video of me doing this, the metal produced was definitely Na, K, Ca. Also it is possible to thermite sodium out of NaOH using magnesium powder in a method akin to this one.

there is a youtube video of this... here let me find it.

http://www.youtube.com/watch?v=908rjHQ5mmc

the safety on this is a little ehh but w.e.

honestly I can not find the pictures I have pictures of the sodium and potassium once it was produced and is in a jar but I can not find a picture of it while it was still in the process. Uhhmm I suggest that you try this yourself before pointing fingers etc etc...

maybe we can throw around somemore thermodynamics with the heat of formation of MgO or whatever.
But i have taken the chunks of metal that have formed from the reaction and dropped them in warm water and have had the characteristic yellow and purple (lilac) flames.

" 'On a commercial scale, lithium is isolated electrolytically from a mixture of lithium chloride and potassium chloride.'

So there you have liquid Li metal in contact with liquid KCl: no reaction. Note also how the LiCl is reduced but not the KCl.

You were melting Li metal, nothing more."

that is a little interesting though considering that lithium has a higher reduction potential than potassium.

[Edited on 15-12-2010 by ThatchemistKid]

Quote: Originally posted by ThatchemistKid

Deep sigh… when in a hole, stop digging.

The Mg/NaOH reaction proceeds PRECISELY because it is THERMODYNAMICALLY favourable:

HoF MgO = - 602 kJ/mol
HoF NaOH = -429 kJ/mol

NaOH + Mg --- > Na + MgO + ½ H2, Heat of Reaction = 429 – 602 = - 173 kJ/mol, highly exothermic. Also: one of the reaction products leaves the reaction (H2), thereby driving the equilibrium further to the right (‘mass balance effect’). Now you can also work out what the estimated end temperature of the reaction products mix would be, in adiabatic conditions (closed system, no heat losses), using available heat capacity constants of the reaction products. At a mere glance the end temperature would probably exceed 600 – 700C. Consistent with observations.

In your thermodynamically VERBOTEN reactions there wouldn’t be any heat evolved AT ALL.

The methods are therefore NOT AKIN: yours don’t generate any lattice energy and simply don’t proceed. In a closed off system with no air, a mixture of NaCl and Li, heated to high temperature WOULD contain some Na because the equilibrium constant K = ([Na] x [LiCl]) / ([Li] x [NaCl]) (brackets indicate activities) and because ΔG = - R T ln K and with ΔG ≈ 0, then K ≈ 1.

On cooling such a mixture would revert back to NaCl + Li because that’s the highest Free Energy for the system.

Regards reduction potentials and ‘reactivity’: the reduction potential is obtained in watery solution and only correlates (not equates) with reduction potentials obtained in different conditions. ‘Reactivity’ is a subjective term that doesn’t mean much: what’s the unit of measurement of ‘reactivity’? There is none.

When reactions take place involving highly ionic solid substances and giving rise to other highly ionic solid, then the highest lattice energy side of the equation determines the outcome.

Hence it’s possible to reduce NaOH with Al (has been done on this forum) but also to reduce AlCl3 with Na (one of the first methods to isolate small amounts of Al – look it up). This seeming contradiction stems merely from a misinterpretation of the reduction potential.

Here’s what happened: armed with your +3 V reduction potential for Li+ + e --- > Li and resulting conviction that Li reduces the world and everything in it, you melted some Li foil mixed with salts and believed you made K, Na and Ca. YOU DID NOT.

ThatchemistKid - 15-12-2010 at 13:28

NaOH + Mg --- > Na + MgO + ½ H2, Heat of Reaction = 429 – 602 = - 173 kJ/mol

I knew the reaction was thermodynamioucally favorable...
and that this discussion was going to lead to talking about all of this.
But, when I tried out the reaction, I will admit that I did not know much about that stuff, so I just tried it.
and for some reason it did work... or atleast I think it did, but in view of all of this evidence, I will go try it again and see what happens. Maybe one of us will learn something we did not know.

blogfast25 - 15-12-2010 at 13:53

There exist many other such ‘anomalies’ that are inexplicable by means of the limited concept of reduction potential. There existed once a process for making Mg metal from magnesia and cokes: MgO + C --- > Mg + CO at very high temperatures. In these conditions both the Mg and the CO escape as gas but can be separated due to their different boiling points. Again driven by Mass Balance Effect.

And you wouldn’t expect the more ‘reactive’ rubidium to be reducible by calcium but RbCl + ½ Ca --- > Rb + ½ CaCl2 proceeds and was once (perhaps still?) used for Rb production. Presumably the HoF of ½ CaCl2 is higher than that of 1 RbCl. That seems plausible based on the lattice energies of both: the twice charged, smaller Ca2+ ions provide much Coulombic energy to the CaCl2 lattice. The higher melting point of CaCl2 (772C) v. RbCl (718C) slightly suggests that too...

But none of these things apply for the reactions you cited…

[Edited on 15-12-2010 by blogfast25]

hkparker - 15-12-2010 at 16:21

Yea I had my doubts on that synth of potassium... otherwise I would have done it by now

I have done the magnesium/sodium hydroxide reaction before with much success but I highly doubt potassium was produced. And calcium hypochlorite is a dangerous oxidizer, it would probably just oxidize the lithium and reduce to calcium chloride or oxide.

ThatchemistKid - 15-12-2010 at 16:23

I am familiar with the those methods that you mentioned (well actually I had never seen the rubidium one, that is pretty interesting) and I have seen a few other similar industrial processes.

anyways I have a couple of videos uploading to youtube that I will post. one of them shows some sodium that I just made, but it isnt too obvious. it was kind of hard to show it that well with this camera. So since the sodium was not obvious I then made some potassium so the color would be more obvious.
I didnt catch the potassium on camera in the first trial but in the second trial there is a lovely purple pop! when i hold a match up to the beaker for the first time... so that should do it.

this one is sodium that I made
http://www.youtube.com/watch?v=nASPY9vB4XU

and this one is potassium that I made
http://www.youtube.com/watch?v=XLp5c9wU8nM
if you pause it at 1:03 you will see the purple flame from the
potassium

http://www.youtube.com/watch?v=6u9LokoiGDE
here is the method that i mentioned using the lithium cigar

also I will upload some picture of the method and some chucks of sodium that were produced.

this is just a picture of some of the stuff used, just the nacl, lithium and mineral oil, the stuff in the background is just some benzene, some chromic acid, some chloroform, and some formic acid that just happened to be there

some sodium after i cleaned it up and pressed it using a glass tube

and what the crude sodium looked like before i pressed it into the ribbion, the lithium burns when I do this, it does not melt, i end up with a hard lithium oxide/nitride/chloride other ide crust on top of the sodium that i have to remove.

Fleaker - 15-12-2010 at 16:42

Quote: Originally posted by blogfast25

There exist many other such ‘anomalies’ that are inexplicable by means of the limited concept of reduction potential. There existed once a process for making Mg metal from magnesia and cokes: MgO + C --- > Mg + CO at very high temperatures. In these conditions both the Mg and the CO escape as gas but can be separated due to their different boiling points. Again driven by Mass Balance Effect.

And you wouldn’t expect the more ‘reactive’ rubidium to be reducible by calcium but RbCl + ½ Ca --- > Rb + ½ CaCl2 proceeds and was once (perhaps still?) used for Rb production. Presumably the HoF of ½ CaCl2 is higher than that of 1 RbCl. That seems plausible based on the lattice energies of both: the twice charged, smaller Ca2+ ions provide much Coulombic energy to the CaCl2 lattice. The higher melting point of CaCl2 (772C) v. RbCl (718C) slightly suggests that too...

But none of these things apply for the reactions you cited…

[Edited on 15-12-2010 by blogfast25]

Just as an addendum to your tutelage on thermodynamics:

The reduction of Rb and Cs chlorides (and bromides and chromates) by calcium, zirconium and maybe magnesium works mostly because those metals can be vacuum distilled. This pushes the reaction forward, much as you mentioned in an earlier post addressing this hogwash. Preparing Cs and Rb this way is via a distillation--it is detailed in Brauer and I can promise you that the prep works as advertised if one uses calcium turnings or powder (granules work only marginally for distilling Cs). The method is so amenable to scaling, I imagine that all of the Cs and Rb are made from their chlorides via calcium reduction. Afterward, I wager the the metals are carefully redistilled in vacuo to rid them of the carried over salts which sometimes distill if careful temperature control is not maintained. The reaction is not exothermic, and indeed must be heated in a vertically oriented tube furnace.
I think any reaction must be thought of first in thermodynamic terms (really quite simple) and then kinetically before anything is ever put in a flask.

[Edited on 16-12-2010 by Fleaker]

BromicAcid - 15-12-2010 at 18:10

I remember reading that if you melt lithium in a Pyrex beaker it will burn a hole through it quite spectacularly.

hkparker - 15-12-2010 at 18:14

Oh that's very true, it accidentally happened to be one time when I was trying to melt it and I discovered it was practically pyrophoric when molten, very easy to light, and destroyed a beaker very energetically.

ThatchemistKid - 15-12-2010 at 18:21

So I am not sure exactly what to think, apparently all the thermodynamics mentioned says that I cannot reduce potassium chloride with Li metal, if I had read that data beforehand I would have never tried this experiment.
But, Its quite obvious to me that I made some potassium, I have it sitting in a jar next to me, it burns with a purple flame as shown in the video i posted earlier. I am not sure quite what to think?

hkparker - 15-12-2010 at 18:26

After watching your video it doesn't look nearly reactive enough to be potassium, more like lithium. So what probably happened it your lithium melted and got an impurity of potassium chloride, so that's the flame color you saw.

Fleaker - 15-12-2010 at 18:36

Molten lithium is phenomenally corrosive to glass and ceramics.

@Hkparker, that is most likely what happened. Potassium has much stronger flame colour than lithium does anyway, and it's hard to distinguish them.

ThatchemistKid - 15-12-2010 at 18:36

I thought about that, I tried this over and over trying to disprove it to myself, the lithium is only on the outside. I get a nice lump of metal on the inside of the cigar with a hard heavy crust on the outside, I can remove the crust and if i take the time to completely clean up the metal formed, when dropped in water or acid it bursts into flames with the characteristic color, if it were lithium fizzling and burning in that video there would have been a very distinct crimson that would have over powered the lilac of the potassium, I have seen lithium burn that is not the right color for it.

I am not trying to argue I am honestly trying to understand.

maybe there is another way to test if this is potassium.

ThatchemistKid - 15-12-2010 at 18:39

hard to distinguish them???

http://www.youtube.com/watch?v=QNojS6ZZ4og

flame tests for lithium, sodium, and potassium.

Fleaker - 15-12-2010 at 18:43

Simplest test is density. Potassium is much denser than lithium is.

ThatchemistKid - 15-12-2010 at 18:43

also you have a good point about the reactivity. I thought that it was reacting a bit slowly, but the one in the video was not cleaned up it was still in that crusty shell of oxide, nitride and chloride.

ThatchemistKid - 15-12-2010 at 18:58

the metal that I have produced sinks in mineral oil. while the lithium foil floats.

the picture is taken from the bottom the lithium is floating at the top and the potassium(?) is at the bottom.

ThatchemistKid - 15-12-2010 at 19:07

clearer version

hkparker - 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)

ThatchemistKid - 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.

hkparker - 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?

blogfast25 - 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.

blogfast25 - 16-12-2010 at 08:34

Thanks also to Fleaker for his addendum which merely confirms the thermodyn. side of things.

hkparker - 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.

blogfast25 - 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…

hkparker - 16-12-2010 at 10:00

That makes sense, I understand

ThatchemistKid - 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

blogfast25 - 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]

ThatchemistKid - 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.

blogfast25 - 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]

cnidocyte - 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.

Ephoton - 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.

hkparker - 17-12-2010 at 23:19

Try posting this in organic

Ephoton - 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]

hkparker - 17-12-2010 at 23:32

Hehe good luck. I've always found organometallic chemistry fascinating

a_bab - 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).

Ephoton - 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

and the window it was a foot a way from

Arthur Dent - 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

blogfast25 - 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]

Ephoton - 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.

Nerro - 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]

hkparker - 18-12-2010 at 13:36

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

blogfast25 - 18-12-2010 at 13:56

Quote: Originally posted by Nerro

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]

NaOH: HoF 298 K = - 426 kJ/mol
MgO: HoF 298 K = - 601 kJ/mol

Hess Law: reaction enthalpy independent of reaction path:

For NaOH + Mg ===> Na + MgO + ½ H2:

Heat of reaction at 298 K: ΔH = + 426 – 601 = - 175 kJ/mol

(break it down as NaOH === > Na + ½ O2 + ½ H2 (ΔH = 426 kJ/mol)
And Mg + ½ O2 ===> MgO (ΔH = - 601 kJ/mol) and use Hess)

Because there is gas evolved, ΔG is probably even more negative, because of Delta S. But - 175 kJ/mol is good enough...

This reaction proceeds with dangerous levels of heat produced.

[Edited on 18-12-2010 by blogfast25]

blogfast25 - 20-12-2010 at 07:31

For those still interested in using lithium in thermite-like reductions, you really need to think fluorides. For example, the very first bench scale productions of plutonium were achieved by PuF4 + 4 Li == > Pu + 4 Li.

Slightly more realistically

, Li should be capable of reducing the trifluorides of the Rare Earths. Here’s an example: the HoF of NdF3 (at 298 K) is – 1,661 kJ/mol:

http://www.sciencedirect.com/science?_ob=ArticleURL&_udi...

The HoF of LiF is – 617 kJ/mol (NIST value).

NdF3 + 3 Li == > Nd + 3 LiF: enthalpy of reaction = 1661 – 3 x 617 = - 190 kJ/mol, very exothermic.

Without having made any detailed thermochemistry calcs, I predict that that’s enough heat to heat the reaction products to above their respective MPs and thus would be obtained in the molten state, allowing the heavier metal to coalesce out, neatly protected from oxygen by the LiF ‘blanket’.

Nd, for experimenters that want to see an Li reduction in action, also has the advantages of:

1. availability: see neodymium magnets (they sell them by the tonne on eBay nowadays),
2. NdF3 is insoluble in water and thus relatively easy to synthesise.

This does mean 'messing' with the inevitably dangerous fluorides. Read up before you start!

Magnesium, BTW, should aslo do the trick...

[Edited on 20-12-2010 by blogfast25]

[Edited on 20-12-2010 by blogfast25]

Arthur Dent - 31-12-2010 at 11:31

Hi guys,

This morning, I noticed that a small quantity of metallic lithium I had put aside in a jar filled with mineral oil looked kinda funny, so I decided to get rid of it because it started to get very dark and foamy.

Since Lithium is still quite valuable, I decided to turn it into LiOH by dropping little chunks in distilled water. The metallic lithium was dabbed with a paper towel to remove some of the mineral oil before being dropped in the water in small 1 cm chunks, but I guess it still had a fair amount of oil on it.

After dissolving the entire stock, I was left with a fairly concentrated solution of Lithium Hydroxide, but unfortunately it seems that the oil was somehow emulsified along in the solution. 2 hours after the procedure, the solution is quite turbid and there is a very thin layer of even milkier-looking oil at the surface. Shouldn't aqueous LiOH be water clear?

My next step would be to filter off the solution. So my idea was to let the solution settle and maybe it would separate a bit more somehow, then use a fine Whatman paper filter to filter off the solution, but i'm afraid that the oil will clog the filter paper. What would be my best way to acquire a clean solution of LiOH with this hideous brew?

Robert
==========
UPDATE: I went ahead and chucked the whole shabang in a funnel with a plain coffee filter paper. Looked like it worked. All the oliy white gunk stayed in the filter and the solution, still turbid, looks free of contaminants.

I have stored the solution in a HDPE bottle, it should keep well? Are there risks that the LiOH absorbs CO2 and slowly precipitates to a carbonate?

Would the chloride be more stable for storage?

Robert

[Edited on 31-12-2010 by Arthur Dent]

plante1999 - 31-12-2010 at 16:21

Quote: Originally posted by Arthur Dent

Hi guys,

This morning, I noticed that a small quantity of metallic lithium I had put aside in a jar filled with mineral oil looked kinda funny, so I decided to get rid of it because it started to get very dark and foamy.

Since Lithium is still quite valuable, I decided to turn it into LiOH by dropping little chunks in distilled water. The metallic lithium was dabbed with a paper towel to remove some of the mineral oil before being dropped in the water in small 1 cm chunks, but I guess it still had a fair amount of oil on it.

After dissolving the entire stock, I was left with a fairly concentrated solution of Lithium Hydroxide, but unfortunately it seems that the oil was somehow emulsified along in the solution. 2 hours after the procedure, the solution is quite turbid and there is a very thin layer of even milkier-looking oil at the surface. Shouldn't aqueous LiOH be water clear?

My next step would be to filter off the solution. So my idea was to let the solution settle and maybe it would separate a bit more somehow, then use a fine Whatman paper filter to filter off the solution, but i'm afraid that the oil will clog the filter paper. What would be my best way to acquire a clean solution of LiOH with this hideous brew?

Robert
==========
UPDATE: I went ahead and chucked the whole shabang in a funnel with a plain coffee filter paper. Looked like it worked. All the oliy white gunk stayed in the filter and the solution, still turbid, looks free of contaminants.

I have stored the solution in a HDPE bottle, it should keep well? Are there risks that the LiOH absorbs CO2 and slowly precipitates to a carbonate?

Would the chloride be more stable for storage?

Robert

[Edited on 31-12-2010 by Arthur Dent]

I ave already made LiOH 4month ago from very fresh lithium from battery. the solution ave some cloudiness but my solution is colorless. I keep it in dry HDPE bottle.

garage chemist - 1-1-2011 at 05:20

LiOH solutions always get cloudy due to uptake of atmpspheric CO2, forming sparingly soluble Li2CO3.
This problem is much more severe with e.g. Ba(OH)2 solution, which is generally always cloudy except immediately after filtration.