iceax3000 Posted July 13, 2009 Share Posted July 13, 2009 i have heard that you can extract the rather good quality lithium from unused lithium batteries, and if so i was wondering if i could use a simple single replacement reaction to extract sodium from sodium chloride. it seem pretty simple but i need some help in the actual way to have the reaction happen.should the sodium chloride be in solution?should the lithium be powdered? any help will be appreciated. Link to comment Share on other sites More sharing options...
Mumbles Posted July 13, 2009 Share Posted July 13, 2009 Well, the fact that this has to get ran in water really kills it. The metals, both lithium and sodium will instantly turn to hydroxide from the water, and potentially explode. Link to comment Share on other sites More sharing options...
a_bab Posted July 13, 2009 Share Posted July 13, 2009 Basically the reaction should be conducted at high temp, in a closed very resistant pressure vessel. The shape of the lithium has no importance. The amount of lithium you'd get from the batteries, and your lack of knowledge makes me think you may get hurt even when trying to extract lithium (there are things in these batteries that hiss and splash and they tend to blind ). Link to comment Share on other sites More sharing options...
Ignes Posted July 13, 2009 Share Posted July 13, 2009 It is possible to produce sodium via electrolysis of molten sodium hydroxide, but if you're not a very experienced (and a bit mentally retarded) chemist you should not be even thinking about it. Sodium hydroxide is pretty corrosive, but molten sodium hydroxide even etches glass! You'd need electrodes of a nickel-alloy as graphite just disintegrates and most metals obviously just dissolve, even platinum. Link to comment Share on other sites More sharing options...
inonickname Posted July 13, 2009 Share Posted July 13, 2009 You can extract sodium and chlorine gas from sodium chloride. It needs to be moltern, then a current is passed through two electrodes. Graphite can be used for the chlorine electrode, I'm not sure what for the sodium one. The chlorine gas can then be bottled or released (bad), or taken by a chlorine scavenger. It will be extremely difficult to replace lithium as it is the highest in the activity series. Water obviously needs to be removed completely. Link to comment Share on other sites More sharing options...
a_bab Posted July 13, 2009 Share Posted July 13, 2009 (edited) Lithium is actually the lowest in the activity series among the alkali metals, and your statement makes no sense here. Lithium has the lowest vapor pressure in the series, so low that the purification via metal vaporisation is difficult, and takes several times more than for sodium for instance. That's why the reaction: Li + NaCl-->Na + LiCl is going to progress to the right, as the lithium is consumed, and sodium released as vapors. Lithium is a good reducer, even for Cs. But all of these are just dry theory, as the user is unlikely to attempt any of these methods just becase he saw on youtube how "cool" the sodium is. Edited July 13, 2009 by a_bab Link to comment Share on other sites More sharing options...
Swede Posted July 13, 2009 Share Posted July 13, 2009 All good answers. Isolation of elemental Sodium by any number of methods has been a bit of a holy grail for the guys at science madness. After reading about their efforts, I decided it wasn't worth the trouble... and it IS a huge and challenging thing for a home lab to do. Link to comment Share on other sites More sharing options...
NightHawkInLight Posted July 13, 2009 Share Posted July 13, 2009 (edited) If YouTube ever gets around to approving my video for revenue sharing I'll link over to my most recent video. It shows rather convincingly that the NaOH/Mg method of synthesizing sodium does indeed produce true metallic sodium. Though I am still foggy on the chemistry. Whatever occurs is EXTREMELY exothermic, and requires no atmospheric O2 to progress. If only NaOH was a more tame substance in its alkalinity and hygroscopic nature it would make for a beautiful yellow/orange star comp. Edited July 13, 2009 by NightHawkInLight Link to comment Share on other sites More sharing options...
iceax3000 Posted July 13, 2009 Author Share Posted July 13, 2009 thanks guys. i know about the synthesis involving molten sodium chloride but i was just wondering if extracting it with lithium would be easier. just a thought. anyways thanks for the help is there anything useful to do with lithium (any good comps that include it) Link to comment Share on other sites More sharing options...
TheSidewinder Posted July 13, 2009 Share Posted July 13, 2009 If YouTube ever gets around to approving my video for revenue sharing I'll link over to my most recent video. It shows rather convincingly that the NaOH/Mg method of synthesizing sodium does indeed produce true metallic sodium. Though I am still foggy on the chemistry. Whatever occurs is EXTREMELY exothermic, and requires no atmospheric O2 to progress. If only NaOH was a more tame substance in its alkalinity and hygroscopic nature it would make for a beautiful yellow/orange star comp. Nighthawk, why don't you upload the video here? It would be a welcome addition. Link to comment Share on other sites More sharing options...
NightHawkInLight Posted July 13, 2009 Share Posted July 13, 2009 Nighthawk, why don't you upload the video here? It would be a welcome addition.I was going to if YouTube gave me much more trouble. However I just got home from work and see it has been accepted. Here it is: Keep in mind I made this video for entertainment value in addition to experimentation purposes, so it takes a moment to get to the point. Link to comment Share on other sites More sharing options...
FrankRizzo Posted July 14, 2009 Share Posted July 14, 2009 (edited) I was going to if YouTube gave me much more trouble. However I just got home from work and see it has been accepted. Here it is: Keep in mind I made this video for entertainment value in addition to experimentation purposes, so it takes a moment to get to the point. Very cool. I'd say that's pretty conclusive evidence that it's a valid technique! Good work, sir. A better recovery technique would be to heat the mineral oil containing the slag to ~200F. The sodium will melt, and hopefully agglomerate into a larger mass, with the slag sitting on top. A fellow named "len1" over on Sciencemadness used paraffin. http://www.sciencemadness.org/talk/viewthr...103&page=14 Edited July 14, 2009 by FrankRizzo Link to comment Share on other sites More sharing options...
NightHawkInLight Posted July 14, 2009 Share Posted July 14, 2009 Very cool. I'd say that's pretty conclusive evidence that it's a valid technique! Good work, sir. A better recovery technique would be to heat the mineral oil containing the slag to ~200F. The sodium will melt, and hopefully agglomerate into a larger mass, with the slag sitting on top. A fellow named "len1" over on Sciencemadness used paraffin. http://www.sciencemadness.org/talk/viewthr...103&page=14Thank you! An interesting discussion on that board. The massive yield speaks for itself on the electrolysis process, that is if you want sodium badly enough to make such a superb cell as was used. I attempted to isolate the sodium by melting, but there was just so much slag in the way that I never could tell if it was working after maintaining high temperature for 30 minutes. I tried it on the very batch in the video, though I ended up just isolating the Na through the method I show. There seems to be a large amount more slag than there is sodium produced, which leads me to believe that the initial reaction is what needs improving. The large orange flames expelled from the crucible also show that much is wasted. How to make the reaction progress more slowly at a cooler temperature yet still be hot enough to leave as little as possible unreacted... Perhaps a bit of mineral oil added prior to the reaction would do the job. Just enough to lightly coat the reactants. Link to comment Share on other sites More sharing options...
NightHawkInLight Posted July 14, 2009 Share Posted July 14, 2009 Sorry to double post, but I believe this warrants a new one. I just asked a chemist I've been in contact with how this reaction progresses, here is his response: "Ah, its nothing hard. The magnesium reacts with the sodium hydroxide to form magnesium oxide, sodium and hydrogen gas. 2Mg + 2NaOH --> 2MgO + 2Na + H2 The enthalpy of formation of magnesium oxide exceeds the enthalpy of formation of sodium hydroxide. It's certainly an odd reaction to the inexperienced chemist but one must remember that at high temperatures thermodynamics plays a more signficant role than kinetics. The extremely high formation constant of MgO is lattice energy driven." Link to comment Share on other sites More sharing options...
inonickname Posted July 17, 2009 Share Posted July 17, 2009 Lithium is actually the lowest in the activity series among the alkali metals Oops. That's what I meant.. Typo. Link to comment Share on other sites More sharing options...
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