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[*] posted on 14-12-2010 at 18:48
lithium and what to do with it


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?
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[*] posted on 14-12-2010 at 19:18


Lithium is not the most reactive.



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[*] posted on 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 :D

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[*] posted on 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.
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[*] posted on 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.
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[*] posted on 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]
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[*] posted on 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]
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[*] posted on 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]
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biggrin.gif posted on 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 :D
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[*] posted on 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  
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

{big snip}

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

[Edited on 15-12-2010 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.

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[*] posted on 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.
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[*] posted on 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]
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[*] posted on 15-12-2010 at 16:21


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

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.
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[*] posted on 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

chemicals.JPG - 165kB



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

sodium.JPG - 153kB


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.

crude stuff in dish.JPG - 152kB
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[*] posted on 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]




Neither flask nor beaker.


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[*] posted on 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.



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[*] posted on 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.
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[*] posted on 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?
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[*] posted on 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.
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[*] posted on 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.




Neither flask nor beaker.


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[*] posted on 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.
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[*] posted on 15-12-2010 at 18:39


hard to distinguish them???

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

flame tests for lithium, sodium, and potassium.
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[*] posted on 15-12-2010 at 18:43


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



Neither flask nor beaker.


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[*] posted on 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.
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[*] posted on 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.


P1010866.JPG - 86kB
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