Sciencemadness Discussion Board - Preparation of Lithium Ammonia complexes in EtO2.(Lithium Bronze)

Preparation of Lithium Ammonia complexes in EtO2.(Lithium Bronze)

Sedit - 21-12-2009 at 16:09

I think my memory may have failed me after a year and may have provided slightly inaccurate data to Len1 and UC. After the discussion VIA PM with UnintentionalChaos I think -38 degrees F was the temperature reaches almost everytime in my snow+HCl experimentation however with the addition of more snow and stirring the temperature could be push as low as -42 with ease. Sorry for the confusion but these experiments where quick and dirty and performed almost a year ago.

With that being said I decided to run the reaction again since I got about 16 inches of material to work with now but to my disappointment I found the battery of the thermometer is now so dead that you can not even get it to stay on let alone read the faint numbers present on the screen.

This was a blessing in disguise as it turns out. What I decided to do was setup a three test tube system in an attempt to liquefy Ammonia using this super cold mixture. Prior attempts last year provided a frost which sublimed on heating to "room temperature"(26 degrees outside at the time). This made me think that my thermometer was tweeking out since it was past its lower end of -32 and that quite possibly it was freezing the Ammonia.

Feeling that this was unlikely(although Nitric acid reportedly reaches like -97 IIRc) and more then likely due to trace H₂O Binding the NH3 into the frost I seen So I decided to add a drying tube in between the flask this time which consisted of EtO₂ and much Excess of Lithium in the Ether to ensure the exiting gas was sufficiently dry. This was more exciting and successful then I could have ever hoped for.

Setup:
Test Tube A: NH3 generator by adding powder NaOH to NH₄NO₃ Mixing completely and stoppering with a tube coming from the top of A leading into the bottom of B. When in operation a few drops of over the counter dilute aq Ammonia was added to initiate the reaction and this was placed in a warm water bath placed on a warming plate.

Test Tube B: Ether straight from the can, enough to cover Lithium from one half of a AA Lithium Advance battery. Lithium was dipped in other starter fluid to remove mineral oil it was stored in then patted dry before quickly chopping it up to the side of small BB's[note1]. A tube ran out of the top of B and down to the bottom of C

Test Tube C:
Lithium prepaired as stated above was placed in the bottom of the test tube with only a ml or so of EtO₂ to aid in thermal transfer.

Test Tube C and Test Tube B Where both placed in the super chilled HCl/Snow mixture.

Procedure
Test tube C was placed in an HCl/Snow slurry and the whole setup was sealed and initiated with a few drops of aq ammonia to start the NH₃ production. Ammonia gas bubbling into Vessel B commenced rapidly as the Lithium beads begun to quickly take on a noticeable copper color which proceded to bright gold as the reaction progressed.

Results:

Test tube A obviously produced Sodium Nitrate + H₂O + NH₃ this was not quantified. Test tube B Is where the unexpected happened. As the reaction progressed I expected that the ammonia would form a while powder due to its reaction with H₂O making LiOH and H+. This did indeed happen and you can see some Of the white powder on the bottom of the test tube. What is interesting though is the fact that the Ether itself begun to take on a slight blue color which quickly darken to an almost royal blue color similar to that of the toolbar at the bottom of a Windows OS system. Ill try to think of a better way to describe it since many may not understand what I'm talking about due to variances in computer systems and settings but as of now that's the best I can do. I have seen references to Crown ether's being used at low temperatures for solving electrons but that is normally well below -38f or -42 and somewhere in the range of -78degrees C If I recall correctly. I believe chilled EtO₂ with a little bit of solvat ed ammonia can accomplish this at a much higher temperature then previously reported. Further test will be performed tonight conditions pending

Test tube C Showed signs of condensate on the sides of the test tube eventually coalescing into a clear fluid liquid at the bottom of the test tube which in short time took on a dark blackish blue color due to the Lithium present "solvating" the electrons. In with this was the same Gold liquid noted in flask B Proving the excess Li that was added to the liquid Ammonia which condensed was also forming the Li[NH4]x complex also noted in literature as Lithium bronze.

Pictures..... Because everyone loves picture I was only about to get the middle vessel after sitting over night since I did not have the camera but to my surprise it was still slightly blue and the Liquid Lithium bronze beads on top where still quite noticeable. I attempted to photograph them best as possible and the thou blueness of it is only present in one of the pictures that shows the brightest blue. The night before It was a bit darker shade more along the lines of light navy blue.

http://www.thevespiary.org/sedit/Private/Pictures/Lithium%20...
http://www.thevespiary.org/sedit/Private/Pictures/Lithium%20...
http://www.thevespiary.org/sedit/Private/Pictures/Lithium%20...
http://www.thevespiary.org/sedit/Private/Pictures/Lithium%20...
http://www.thevespiary.org/sedit/Private/Pictures/Lithium%20...
http://www.thevespiary.org/sedit/Private/Pictures/Lithium%20...

PS: When I get the pictures re sized I will embed them into this thread but as it stands there pretty huge and I don't have an image editor since my computer crashed. Anyone willing to crop them please do and let me know, Thank you.

I am scaling up in a little bit and I will try to document the entire thing for you folks this time and do it 100% right.

Peace and I hope you enjoy the pretty pictures... More to come because this will allow any chemist to perform the powerful birch reduction at room temperature without the need for Liquid ammonia.

[Edited on 22-12-2009 by Sedit]

Sedit - 21-12-2009 at 20:36

Now as far as discussion of this topic goes I would like some feed back please. With little interest in the condensation of Ammonia even though it has its own forms of potential in organic synthesis. My main goal at the moment is the Lithium Bronze complex AKA Li[NH3]4. This is the substance which is gaining notoriety for its ability to perform Birch like reductions at elevated temperatures. This is in a much more stable and condensed form of the common solvat ed Lithium in ammonia system which normally takes on the complex form of Li[NH3]x where x is any number of stable ammonium ion complex. This Li[NH3]4 system forms a biphasic mixture in a non polar solvent which can be stored at room temperature for several days before it begins to decompose.

This will allow Birch reductions to be performed at room temperatures in short amount of times without the hazards of gaseous Ammonia which carrys the risk of suffocation and/or severe chemical burns. I did have a little scare when performing this experiment because the text that lead to this experiment suggested performing a birch reduction in an atmosphere of NH3 in a balloon. I was a little Leary of doing so because if the balloon pops it could turn into a bad day which it did manage to do. Only a little pin hole in a balloon that was hardly inflated at all which blasted me in the face for a split second taking the breath out of me and scaring me more then anything. No harm was done but it emphasizes the dangers that could come from doing it in a balloon without proper testing of its chemical capabilities. There is also a possibility that I myself was the reason the pin hole formed but either way care should be taken when handling concentrated ammonia gas.

The main objective of all this work is to synthesis one of the so called Super Bases Lithium Amide with the chemical formula LiNH2. Since the Nitrogen atom is completely unhindered it can act as nucleophile aminating various alkylhalides to there amines thou an SN2 reaction in high yields If I recall correctly. It has the ability to deprotanate Dimethylsulfoxide when can then be used as an alkylating reagent. The methylation abilities of LiDMSO is one of my main interest right now. Normal synthesis of this compound involve the direct combination of its elements by adding Lithium Metal into Liquid ammonia at super chilled temperatures then distilling the Ammonia to yield the the LiNH2. This bypasses all of this using gaseous Ammonia forming the Lithium ammonia complex which can be decomposed by the formula,

Please correct me if I'm wrong because this seems unreasonable to me.

Li[NH3]4 ---Heat---> Li[NH3]4---> LiNH2 + 3NH3 + H+

Can anyone suggest alternative drying agents instead of wasted Lithium wash bottle. I hear some cause complexes but I can not remember which cause this.

[Edited on 22-12-2009 by Sedit]

not_important - 22-12-2009 at 01:15

Quote: Originally posted by Sedit

...
Can anyone suggest alternative drying agents instead of wasted Lithium wash bottle. I hear some cause complexes but I can not remember which cause this.

Obviously anything acidic is out, thus no H2SO4 or P2O5/P4O10. Calcium and magnesium salts often form complexes with ammonia, both MgCl2 and CaCl2 will form complexes that volumetrically contain almost as much ammonia as liquid NH3.

NaOH and KOH do well for preliminary drying, CaO can be used to finish the job. If you have ready access to CaO then just it alone; activate it by heating it to ~700 C for 15 minutes or so.

If the NH3 needs to be really dry, I'd use the lithium drying bottle for a final polish.

Sedit - 22-12-2009 at 01:54

I added an extra layer of NaOH atop the NH4NO3/NaOH mixture which I didn't do the second time(more on that attempt tommorow). I do believe that the extra layer aids in drying and a whole seperate drying tube with it may be prudent.

CaO sounds like an excellent Idea N_I. I used the middle drying bottle in the first attempt with Lithium grains in Ether because I could not think of anything else powerful enough and on hand considering last time I attempted MgSO4 as a drying agent in the past, which as you already stated, proved a huge mistake. When I did it I was getting NH3 exiting out of the system so I did not think much about it till I opened the drying flask that contained the MgSO4 some two to three weeks later and almost lost my breath to the amount of ammonia it released. I want to try to harness this in the furture as an ammonia storing system but thats a topic for a whole another thread.

I don't feel I need it really dry because I repeated the experiment just a little bit ago hence the reason im typing this 4:40 in the morning. I just feed NH3 from the basic generator right into an EtO2/ Lithium suspension and the formation of Lithium bronze formed rather nicely although there where a few complications and I will get into more detail about them tommorow and show the shitty pictures I took of the whole ordeal.

I want to find a solvate able to extract the Li[NH3]4 Aka lithium bronze without it reacting if such a solvent exist other then liquid ammonia.

Im going to experiment with this just a little bit more as we still have a good amount of snow on the ground but after that I am moving directly on to the synthesis of Lithium Nitride AkA Li3N which is a strong enough super base to deprotonate H2 itself synthesizing 1mol of Lithium amide and 2 mols of lithium hydride

thanks again for the suggestions and if anyone else has enything to share feel free. Keepem comming folks because sedits pretty excited over this one right now :cool:

Sedit - 22-12-2009 at 16:19

Attached is a photo of the the midway point or so of the Lithium Ammonia reaction. You can slightly see the tube which is full of pin holes going to the bottom of the flask. Atop of that is the Lithium Bronze already beginning to form pretty early. I will post more details of the setup later time permitting but I mush go now.
start.jpg - 27kB

Keep in mind that this attempt was not as successful as it may first appear due to a number of reasons I will explain later.

Ephoton - 28-12-2009 at 06:01

sedit if the amide is your goal then instead of heating the lithium bronze to obtain it
just add styrene or a simular alkene to the ether mix.

you will then reduce the alkene and precipitate lithium amide with out having to
heat the solution.

Sedit - 28-12-2009 at 11:53

My Main goal is to experiment with the stability and reactivity of the Lithium Bronze solution. I have obtained results showing a blue mixture which goes against what literature appears to state so my main goal at the moment is to figure out why there was a blue solution in the first place.

It appears that the bronze is more stable then I expected it to be. The addition of EtOH to the mixture precipitated a white powder and released a strong smell of ammonia from the flask.

I have alot more to add but been very busy due to the X-mas and what not so I will be riving this thread as soon as I find the time to post all the pictures and the write ups I have prepaired.

PS: Can anyone view the pictures posted in my first post? They do not appear to have linked correctly and I can relink them if needed.

stateofhack - 28-12-2009 at 16:16

I can't see them, i think hotlinking them is what is causing the problem

not_important - 28-12-2009 at 17:52

All the image URLs end with "Lithium%20Bronze/..." - something happened to the trailing section including the actual file name.

You might try running them through tinyurl to get a short URL.

If you are running MS Widows, download IrfanView for an image edotor. Free, should handle most if not all your resizing, cropping, contrast and brightness adjustments, and simple annotation needs just fine.

Sedit - 28-12-2009 at 18:31

Im just downloading Adobe photoshop as we speak so I can edit the size and place them inline like I wanted to. Before they where so large I did not want to :IMG: them in like I normaly would. Give me a few hours and I will update it all.

Just eliminate the ... at the end of the file and it will take you to the directory to view the files.

not_important - 28-12-2009 at 18:54

Lordy, using Photoshop to resize images is like using a nuclear aircraft carrier to cross a stream. Takes forever to load, eats memory like Microsoft had coded it; on top of that it likes to phone home to Adobe and leave lots of footprints in files it edits.

Trust me, IrfanView will do what you need with a lot smaller footprint, and is free-as-in-beer. You can even batch process a group of files, say resizing and changing the file type for them all in a few seconds.

Sedit - 28-12-2009 at 22:33

I know... I know....but its what im use to and since I do a good share of graphical work I try to get the best I can. Since my computer crash I have so many unexpected issues its unreal because the back up version was corrupted and missed many DLLs and what not. I installed Photoshop and now I keep getting an svchost.exe - No Disk error blah blah.... ::Sigh:: Either way the more important pictures should be avalible tommorow. Assuming my computer does not learn how to fly and decide to take a leap out the window......

Should I

A:Cancel
B:Tryagain
Or C: Continue

Im chosing option D at this point which normaly involves a sledge hammer

Sedit - 12-1-2010 at 20:13

Ok so im considering returning to the experiments very soon. Caught up in Xmas New years and my birthday but now I want to review the synthesis and perform it again on roughly a two gram lithium scale. I have trouble sleeping at night because of the test tube with 1.7 grams of Li thats been in Et2O since Xmas. If that burst from pressure build up I have a serious issue on my hands and Im not happy about it and would like to make use of it. I have aired it a few times and it appears to be no longer building pressure but it still disturbs me.

The reason im bringing this up with no experimental data like I normaly would is because I want to get some thoughts on the complexe between Copper and Ammonia that forms.

If powdered Cu was added here what would happen?

I know it complexes with NH3 but would it do the same in this case? What would be the product? A complexe with Lithium, Copper and Ammonia or what?

Sedit - 17-1-2010 at 01:44

I spent much time making sure everything was on the up and up.

I had my drying tube of NaOH flakes in place, suck back tube in place. Everything was dryed in the Microwave to be sure it was H2O free. I mixxed a slight excess of NaOH with ammonium nitrate to be sure to add that little extra drying ability. All tubes feed into 1.7 grams of Lithium in Et2O.

It all was well and I checked and double checked to make sure all was sealed and air tight which they where but the problems were yet to come and rather rapidly at that.

See See it took me about 2 hours to setup and photograph to be sure I went down my checklist of alls well.

Once I was sure it should go correctly I mixed CaCl2 hydrate with Crushed Ice and to my suprize I did manage to quickly get the temperature down to an easy -35/-40 degrees F i believe(maybe C but in the confusion i forget)

After mixing an amount of ammonium nitrate with an slightly larger then eq molar amount of NaOH I could smell ammonia as expected. This was taking its good old time till it ran away from the liquid NH3 solution stated below. I had an icebath ready but this proved futile. To initiate such as I normaly do I drip a few drops of ammonia hydroxide solution to get it ready to go.

Perhaps big mistake..... or just giving a boost to whats gonna happen anyway.

Well withing about 1 minute I started to hear a" loud" fizzing sound proving the NH3 was being created(way to fast at that) but somewhere on the line it was either pluged or there was to much pressure thru such little lines(doubt it I believe the NaOH drying tube did it). First instinct was to attempt to cool the reaction vessle to quell the runaway but it was to late and of the seals between the lines busted blastingme in the face with 100% pure NH3 giving minor chemical burns. I made sure I was ok then quickly got a cup of HCl and placed the leaking tube in it filling the area with relatively harmless ammonium chloride. This helped alot and I quickly diluted the NH3 generator with excess H2O. I blew HCL fumed across the entire work area until ammonia chloride stopped presenting itself and made sure to wash completely even thought I keep myself cover almost entirely when working.

uggg I got alot more to say but il write tommorow.

Peace and be safe
I have pics of the setup but when my hands stop hurting tommorw ill post them.

~Sedit

[Edited on 17-1-2010 by Sedit]

Sedit - 17-1-2010 at 11:50

Sorry for the double post but since almost no one appears to have interest in a highly useful and reactive alloy between a metal and a gas that gives a very powerful and even more useful superbase upon decomposition im left to pretty much talk to myself and record my notes here like a good chemist should. On a brighter note atlest someone heres doing instead of just talking which is a plus right.

Well anyway the reasoning for this post is not to rant but to say that the cause of the seal blow out has been determined and it was exactly what I thought it was. A prill of NaOH from the drying bottle managed to clog the inlet comming from the Ammonia generator. This is what lead to the quick catastrophic failure of the tube that connected the drying test tube and the Ammonia generator. When it popped off it turned right towards me like a high presure fire hose would have and I had to quickly grab it to avoid being completely overwhelmed with NH3 gas. In the process I suffered very minor chemical burns on my right hand but nothing more then a bit of redness on my hand(And I almost ALWAYS where gloves of some kind but my Ego got in the way this time and I didn't see the need Save the scolding for someone that doesn't know better). As stated the tube was quickly lead into a small vessle of Muratic I had on hand right there just for this sort of thing so I guess being a boy scout and the always be prepaired motto paid off for a change eh

.

Well Im going to revamp my drying tube and possibly attempt to reproduce this experiment tonight since the Lithium vial was open and surely more decomposition is to be expected although after stoppering the vial there was no pressure build up so thats a very good sign showing that H2 was not being produced due to decomposition. The fact that its Sunday and I may not be able to find time to perform it tonight since I have to get up early and perform most of my work at night when the young budding chemist are all asleep.

:Sigh:

I would like some kind of feed back here folks even if its to tell me im a complete idiot. If not I will more then happily take my research elseware. But I can almost assure you that aside from the synthesis of the Lithium ammonia complexe that there is some form of contamination that can solvate the electron in ether.

This effect has already been noted in various ethers as it stands at extremly low temperatures so the idea of finding something to promote this effect at relativly higher temperatures is noteworthy indeed if not publishable if you ask me.

Also I would like to know Ethers aside what Non reactive non polar solvent has the highest solubility for NH3.

Thanks for your time for those that felt the need to show some interest!

watson.fawkes - 17-1-2010 at 14:08

Sedit, I can affirm that you are a good person.

Also, I am reading this thread with interest but nothing to say.

aonomus - 17-1-2010 at 18:58

I'm also reading this thread with interest, but I don't really have anything to add or comment on. This is something that is interesting, but I can't replicate it myself, nor contribute, so I can only cheer from the sidelines. Keep at it Sedit

Sedit - 17-1-2010 at 21:21

thanks for the support guys. Anonmus trust me if I can do it you can do it

. There is no doubt in my mind about that. I am just now getting prepaired for another take at it and I made a new drying bottle clog proof and in about 10 minutes alls a go.

Honestly the drying bottle is "useless" other then increasing yeilds because H2O turns to Hydrogen and LiOH while the rest of the lithium complexeswith the Ammonia.

Now can someone please explain to me the REAL structure of this compound? Is there a delocalized electron in there somewhere thats screwing with my attempts to draw this structure? They state Li(NH3)4 but this makes little sence to my chemicaly declined mind.

When Im finished this I will attack many other complexes with NH3 just out of curiousity. Copper seems like it could be interesting honestly.

Anyway I will try to report and post pictures tommorow but as many know my computer is crap and just getting pictures posted and resized is a PITA so we will see. When all gets worked outI will have a mess load of eye candy for you folks.

aonomus - 17-1-2010 at 21:52

Well its not so much the resources/chemicals, but the location. From the way you describe it, you went outdoors and did it on a porch or something. I live in a dense townhouse block with no completely fenced in backyard. Someone taking their dog for a walk could make my life complicated to say the least.

I'll see if I can look up some references for you Sedit over my lunch break, perhaps something in the literature can indicate what the Li(NH3)n complex is.

Also, its already known that Cu*(NH3)n complexes exist, probably a few videos on youtube. I would maybe like to see fine powdered Cu attacked by gaseous NH3 out of curiosity.
You could probably go
CuSO4 + Na2CO3 -> CuCO3 -heat-> CuO
CuO + C -heat-> Cu(s)

Although one tip I might have for you, take a trip to the pet store and look for plastic airstones, they should be resistant enough to solid NaOH prills, and could solve the clogging/exploding/owowhand problem for you.

Oh, and you can do [ i m g=WxH ] url [ / i m g ], but it only resizes in the browser, so the people viewing the thread still download all that data.

Sedit - 18-1-2010 at 02:30

Give me tommorow and you will see it all. I will update some pictures and inline them.

I just reproduced but got some of the strangest results yet.

No blue color.... No clear with Gold biphasic mixture,.......

Brown

and I have no idea why the hell the Et2O turned brown. All im left to assume is that its some sort of colloide with the Lithium bronze or someshit.
Ill let you see the pictures tommorow and see if we can make sence of them.

Gotta run tired and in a hurry.

watson.fawkes - 18-1-2010 at 07:20

Li(NH3)4 is an excellent lesson on why chemistry is applied quantum mechanics. The standard rules are not fundamentally true in any way, but represent approximations of the standard central-force solutions to the Schrodinger equation. Sometimes the full picture doesn't follow those standard rules. This article on the compound in question
A Molecular Perspective on Lithium-Ammonia Solutions has full text available at that link. It's a review paper and is more readable than a research paper. Here's the abstract. (Some formatting errors are present.)

Quote:

A detailed molecular orbital (MO) analysis of the structure and electronic properties of the great variety of species in lithium-ammonia solutions is provided. In the odd-electron, doublet states we have considered: e-@(NH3)n (the solvated electron, likely to be a dynamic ensemble of molecules), the Li(NH3)4 monomer, and the [Li(NH3)4+ e-@(NH3)n] ion-pairs, the Li 2s electron enters a diffuse orbital built up largely from the lowest unoccupied MOs of the ammonia molecules. The singly occupied MOs are bonding between the hydrogen atoms; we call this stabilizing interaction HH bonding. In e-@(NH3)n the odd electron is not located in the center of the cavities formed by the ammonia molecules. Possible species with two or more weakly interacting electrons also exhibit HH bonding. For these, we find that the singlet (S=0) states are slightly lower in energy than those with unpaired (S=1, 2) spins. TD-DFT calculations on various ion-pairs show that the three most intense electronic excitations arise from the transition between the SOMO (of s pseudosymmetry) into the lowest lying p-like levels. The optical absorption spectra are relatively metal-independent, and account for the absorption tail which extends into the visible. This is the source of Sir Humphry Davys fine blue colour first observed just over 200 years ago.

Sedit - 18-1-2010 at 10:11

Sorry if its a little large I am having a bitch of a time with this computer and did manage to get one decent picture resized so show whats going on. This is completely different results then I had before and the only thing I can think of is that it has something to do with the fact that the Lithium has been sitting in the Et2O since Christmas

This is the very early into it and next to no Lithium bronze has started forming yet although if you look very closely you will see very small specks of it in there that sort of resemble bubbles

This morning I woke up to find the Ether is a greenish color with all sorts of precipitate and unreacted Lithium out the wahzoo. I had major issues with the NH3 generator which highly complicated things to say the lest keeping me up until about 6 in the morning which is something im not happy about working around this stuff while half asleep.

Im going to hunt down the Kipp generator I made sometime back and use that from now on since that should be able to give me the nice steady supply of Ammonia I need for the experiment. I did manage to get a good amount of the bronze complex and if I can find out WTF I did with my camera 6 in the morning I will take pictures and post so that you all can see the now Green Ether.

I will post more later but its early and the kids are off school so I have a mess load of chours to do(God I sound like a woman

).

PS: I also did away with the drying tube which I think I forgot to mention before I believe. So as it stands you DO NOT need to dry the ammonia to form this complex.

watson.fawkes - 18-1-2010 at 20:22

Quote: Originally posted by Sedit

Now can someone please explain to me the REAL structure of this compound? Is there a delocalized electron in there somewhere thats screwing with my attempts to draw this structure?

When I posted the review paper above, I had only scanned it briefly. The answer to your question, specifically, is yes. It's described in section 5.1. "The Singly Occupied Molecular Orbital (SOMO)". That's where your electron went. There's a density diagram for it. It's somewhat like a Rydberg state (high principal quantum number around a center), but with a deformation where the electron density is concentrated around the four triangular faces of a cuboctahedron, each face corresponding to an ammonia unit.

Sedit - 18-1-2010 at 22:53

LOL thanks for the responce watson i mean it.... but its late and ill have to get back to you on that one tommorow because there was so many big words in there

.

I kind of get the point though ust a bit confused.

watson.fawkes - 19-1-2010 at 07:55

Quote: Originally posted by Sedit

[...] because there was so many big words in there

. [...] I kind of get the point though ust a bit confused.

You think you're confused now, just wait to you get to the article. I'd highly recommend reading it, even if it takes you a month or two plow through it while you learn the background material. It's the most direct way to getting smarter I know of.

Apropos the biphasic mixture of blue solution with bronze solution on top, it only occurs at lithium concentrations and temperatures. There's a phase diagram in the paper you'll find useful in that regard.

Plus, I didn't know that the lithium bronze solution isn't just metallic-looking, it's a proper liquid metallic phase! And, if that weren't enough, it's a better electrical conductor than mercury. That paper is full of great stuff, such as a computation of the density of solvated free electrons (they become part of the solvating cage, more or less), interesting discussions about the transition to metallic state (TMS) (a phase transition that's just not all that common), a novel kind of H-H bond (not the typical hydrogen bond, but enough to stabilize complexes).

It's also yielded the first real insight I've ever had about the quantum origin of the liquid phase. They do computations for coupled radicals, starting with the dimer. The dimer has a number of geometric configurations which are all very slightly exothermic, a little over an eV/molecule, which is down at thermal energies, and which differ from each other by about an order of magnitude less. So one can pretty easily imagine how the liquid phase arises when bonding can happen promiscuously upon collision, given the high geometric density of permissible configurations, and yet also break because of fluctuations from thermal motion.

Sedit - 19-1-2010 at 09:43

Im trying to get the paper Watson but I keep receiving this error

Session Cookie Error
An error has occured because we were unable to send a cookie to your web browser.

I set it up to accept all cookies but the problem is still there. Would you be kind enough to upload the paper I really want to read it.

[EDIT]

Quote:

Plus, I didn't know that the lithium bronze solution isn't just metallic-looking, it's a proper liquid metallic phase!

Yes there is no doubt about this at all. If you ever decide to prepair this(its pretty easy honestly just hard to do it right

) you will see that it is indeed a liquid without a doubt. It mostly likes to form small globs that clump up and disperse with shaking just like you would expect from a thick oil on water basicly. The only difference is its metallic in appearance with a bright gold color. The unreacted lithium is now Royal blue in color just as I would expect a normal birch to look. It is reminisent of (aq)Cu salts +ammonia complex colors

[Edited on 19-1-2010 by Sedit]

Sedit - 21-1-2010 at 07:50

All attempt to isolate the upper layer have failed. This substance appears to wet everything it touches. It is very sticky to say the lest.

However I did notice something very interesting and I wish I had my camera at the time but when attempting to pipette what I thought was going to be a mixture of Et2O and the Lithium complex the pipette was filled with nothing but the thick gold material. I attempted to dropper it into another flask for storage but it wouldn't come out. As I watched it the gold began to change color over the course of a few minutes and went from Gold to deep royal blue color. It was really cool to watch but annoying that I could not get it out of the pipette.

In the morning the pipette was full of off white powder with a little bit of blue and gold specks deeper in it. The fact that any of the gold or blue color remained overnight is a testiment to the stabilty of this substance. It has remained in a less then dry enviroment for a while now stoppered in a test tube poorly yet the upper layer still remaines.

Im having problems with the wasted Ammonia that annoys me. Everytime I have calculated excess NH3 to be generated by far yet when it begins to run out there is still more often then not excess Lithium present. I am considering rigging up something to trap the ammonia into a balloon after it exits the Et2O in an attempt to conserve it and allow it more time to react. I also plan to attempt this at a much larger scale of several grams at a time which would allow me to rapidly stir the mixture as the Ammonia is feed in. This will more then likely solve many issues right away.

As it stands though I think if one where able to synthesis this, isolate and store in a vial in a freezer it would keep for many weeks if not months at a time. One or two more attempts will be made with Lithium and then I will move on to Copper since im curious as to what an anhydrous complex of it and Ammonia would appear like. First I have to get some Ammonia sulfate because its cheeper then the Ammonia nitrate I have been using but I have been saving the Sodium nitrate so its still not to bad a deal.

PS: That paper you provided is awsome. It will no doubt take me alot of time to fully read thru and understand but non the lest I love it thanks. The images of various colors look so familiar it isn't even funny. I have seen this substance as blue, gold, and red so far and the paper allows me to gage ammonia concentration thru that alone

Sedit - 4-3-2010 at 10:12

This is a little off topic but my recent attempt at making the complex was a fail due to messed up NH3 production. I can only think that ammonia nitrate come in a ydrated form and since I was using recrystalized ammonia nitrate before that some H2O was carried along and allowing the reaction with the NaOH to take place. This time I just powdered NH3NO3 right from its cold pack source and it did nothing but produce very little NH3 and just clumped together wasting alot of reagents.

Anyway can someone clear up something for me. Is there anyreason why a birch reduction can no be performed in an alcoholic solution of Ammonia in the presence of Lithium. I would think that the reaction mechanics would stay the same yet I am unsure about the reacion mechanisms although some of what I have read have suggested it would be possible.

Anythoughts?

DJF90 - 4-3-2010 at 10:26

t-Butanol is often used as a proton source in dissolving metal reductions using Li/NH3(l), for example the birch reduction of anisole. When reducing an enone with Na/NH3(l)/t-BuOH, inclusion of just 1eq. of alcohol allows the isolation of the saturated ketone, as opposed to the alcohol, which is the further reduction product. With no proton source, and using Lithium as the metal, then the Li-enolate is obtained prior to work up, which can be reacted further.

crystalXclear - 7-1-2011 at 10:40

Quote: Originally posted by aonomus

Although one tip I might have for you, take a trip to the pet store and look for plastic airstones, they should be resistant enough to solid NaOH prills, and could solve the clogging/exploding/owowhand problem for you.

Verry interesting thread sedit, keep up the great work.
Sorry Xtal can't be more help on the subject matter, but one way to get around the clogging (Xtal has one of the plastic air stones from the states, and has offten thaught about it's potential, great idea aonomus. But Xtal remembers dreaming of the aquarium tube being firstly
filled with an inch of cotton wool, This cotton plug is then pushed up the tube, leaving just over an inch of hollow tube. Then scealed with heated long nose pliers.
and finaly, a hot pin is used to pierce the hollow end repeatedly, creating a verry good alternative to air stones.
The tube can also have a small quantity of granular drying agent added, but may not be practical for some aplications.

Hope this is of some help. As Xtal has only read the thread to this post thus far, someone may have allready sugested it. In which case, Xtal is sorry for any repatition.

May the alchemy smile upon you in your quest. Xtal

Sedit - 7-1-2011 at 11:17

Thank you, in essence what you talk about is pretty much the method I ended up taking for feeding the Ammonia into the reaction. I just thinned out the end of the tube and poked many holes into the tube to dispense the Ammonia. I still feel given the option I would go with Airstones if you have one.

Its cold enough to start experimenting with this again but I don't really have the funds or resources on hand right now. If you look there is another fellow here that also reproduced the complex as well.

Since you brought this up again I have noticed that most of the pictures are gone due to my site being taken off line so I will see what I can do about re-uploading them.

plastics - 26-2-2011 at 06:07

Excellent thread. Have been looking for a way to make lithium amide in order to use it to synthesise 9, 10-bis(phenylethynyl)anthracene (one of the fluorophors used in glow sticks) according to this patent:

http://www.freepatentsonline.com/3911038.html

The sequence to lithium bronze and onwards to lithium amide appears relatively straightforwards according to the beginning of this thread

Anyhow I set up an ammonia generator using solid NaOH and (NH4)2SO4 initiated by a few drops of water. THe ammonia was led through a 'U' tube containing CaO onto a Dreschel bottle containing lithium in dry hexane. The exhaust from this went via a rubber tube to an inverted filter funnel just dipping into a bowl of water to absorb any excess NH3. These are the results after 1 hour and 10 hours:

bbartlog - 26-2-2011 at 06:43

Nicely done plastics. That's quite an attractive layer of bronzeish liquid there, looks almost like a machined piston.

aonomus - 26-2-2011 at 08:05

Impressive and beautiful. Liquid bronze!

Sedit - 26-2-2011 at 15:21

Wow, im am so jelouse you don't even know. I have never produced such a large amount that is if I understand the scale of what im looking at.

Tell me this isn't some of the coolest looking stuff you have ever seen. Its like Gold and Mercury had a baby and this is what was produced. Its awsome to suck it up into a plastic pipett and watch it change from gold, to red to purple to blue and in the end a white or grey powder.

Its obvious to me that your Lithium has come from a much better source then mine which originated from batteries.

I really want to know the experimental detail if you have them avalible.

What was the weight of the Lithium at the start of the experiment?
How much Ammonia Sulfate and NaOH was used in the NH3 generator. I always used Ammonia Nitrate/NaOH and the generation of Ammonia gave me problems on a few occassions messing up experiments before they even begun causing me to lose materials.
Did you by any chance measure the weight gain compaired to the start of the experiment since this would give a pretty good idea of exactly how this complex is contained as long as the NH3 in solution of hexane is accounted for.

Did you stir, I never did and Im almost 100% positive that this would dramaticly cut down the time needed to take this to completion.

Did the ammonia generator make way to much ammonia at first then die off quickly wasting alot of ammonia or was it relatively steady. If it was steady how did you go about mixing the materials? To fine a power always caused me issues. I put off reproducing this experiment until I had a kipps and a sufficent dryer so that I could do this experiment the way I knew I should have been doing it.

I know its not very important but any idea on the amount of Hexane used? I always used Ether which contained a small amount of Hexane.

Im sure I have many more question but thats the basics for now, im going to go back to admiring your pictures now. Awsome work. This has always been one of my favorite reactions I have performed because the end result is just mesmerizing to look at to the point where I never really made much real use of it. I just played with it till it decomposed

plastics - 27-2-2011 at 04:08

Yes I agree this stuff is very eerie and mesmerizing especially as all the globules formed, got bigger and then coalesced to a contiguous layer.

For the ammonia generator I used 2 mol (80g) NaOH (sink unblocker) and 1 mol (NH4)2SO4 mixed in 500ml erlenmeyer fitted with single holed rubber stopper and glass tube to the U-tube containing CaO. The NaOH was small prills appox 1mm diameter and the ammonium sulphate similar sized crystals. On adding the water the whole thing went crazy and started effervescing immediately, ammonia bubbling through my scrubber without even dissolving in the water resulting in some loss. The ambient temperature was about 10 degrees C. Slowed down gradually and ultimately required some gentle heating. Tried to keep a balance between production and dissolution via the inverted funnel

Lithium prepared from small chunks I acquired - all very old and covered in grey hydroxide/nitride. Used single 1.5g piece prepared as per Vogel ie pounded with a hammer until thin sheet, cut into strips into anhydrous ether and then into small squares directly into the hexane in the Dreschel bottle

No stirring of the hexane/lithium mix. Bubbling of the ammonia was sufficient to make all the pieces 'dance'. Initial hexane volume was 100ml - unfortunately I didn't think to weigh the mix before and after - wasn't sure it was going to work!

garage chemist - 27-2-2011 at 04:40

I did something similar a while ago, when I got my cryostat. I condensed about 20ml liquid ammonia and dissolved a few pieces of lithium in it, it went deep violet/black.
Then I let the ammonia boil off and was left with what looked like liquid gold, at room temperature. It also formed a firmly adhering mirror on the glass that looked like a copper mirror.
I added a drop of this "liquid metal" to a bucket of water, and it VIOLENTLY EXPLODED with a deafening crack!
That was enough to put me off further experimentation with concentrated alkali metal/ammonia solutions. The thought of the whole vial detonating like that was horrifying, and I let it lay in the backyard until atmospheric moisture had taken care of it.

I would think that atmospheric oxygen had reacted with the solution, forming a hyperoxide of some sort which would detonate with the residual alkali metal when provoked.

In Brauer, I found a method for preparing potassium hyperoxide by simply bubbling oxygen into a solution of the metal in liquid ammonia. It also warned of explosions frequently occuring with this method unless a special apparatus is used that refluxes the ammonia and washes down the solid crusts that tend to form on the sides of the reaction vessel.

aonomus - 27-2-2011 at 07:20

Quote: Originally posted by garage chemist

I did something similar a while ago, when I got my cryostat. I condensed about 20ml liquid ammonia and dissolved a few pieces of lithium in it, it went deep violet/black.
Then I let the ammonia boil off and was left with what looked like liquid gold, at room temperature. It also formed a firmly adhering mirror on the glass that looked like a copper mirror.
I added a drop of this "liquid metal" to a bucket of water, and it VIOLENTLY EXPLODED with a deafening crack!
That was enough to put me off further experimentation with concentrated alkali metal/ammonia solutions. The thought of the whole vial detonating like that was horrifying, and I let it lay in the backyard until atmospheric moisture had taken care of it.

I would think that atmospheric oxygen had reacted with the solution, forming a hyperoxide of some sort which would detonate with the residual alkali metal when provoked.

In Brauer, I found a method for preparing potassium hyperoxide by simply bubbling oxygen into a solution of the metal in liquid ammonia. It also warned of explosions frequently occuring with this method unless a special apparatus is used that refluxes the ammonia and washes down the solid crusts that tend to form on the sides of the reaction vessel.

I'm not terribly surprised by the potassium oxide/peroxide formation, leaving the metal instorage does form unstable peroxides, hence storage under argon.

If it was that unstable, what was used to quench this material, was methanol too reactive?

roamingnome - 5-3-2011 at 19:35

Ive been wondering, are these bronzes or royal purple magnetic in anyway.
What would a neodymium type magnet do ?

Also, anyone considering NH3(g) for experiment should read this bulletin
Attachment: 4385716.pdf (1.6MB)
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-_1 .4 - UNITED STATES ATOMIC ENERGY COMMISSION
-RMO- 203 6
AMMONIUM' SULFATE DECOMPOSITION

plus the ZnO can be recycled.
I mixed finely powered reactants and put them in a ceramic jar. I have not heated the mixture up at this time, but even at room temperature, opening the jar produces a "smelling salt" type jolt.
cheers...

[Edited on 6-3-2011 by roamingnome]

[Edited on 6-3-2011 by roamingnome]

Sedit - 6-3-2011 at 13:57

Don't take my word for it since its been a while since I reviewed the paper but I believe the document that Watson provided early in this thread stated that the complex showed Dimagnetic properties. They do discuss the magnetic properties as well as the conductivity of the complex so its would be worth the read.

I have been looking for a better means of Ammonia production because NaOH and Ammonium salt is just to finiky producing a rapid jolt of Ammonia then a decline not to mention its wet gas. However the temperatures run for that experiment seem rather high and alot of cooling would be needed on the gas before it was able to be used for this experiment.

roamingnome - 6-3-2011 at 17:00

thanks..

yes dimagnetic probably because the free electron is not orbiting around a F orbital causing a moment. Still sounds like a fun do-little physics project. Magnetohydrodynamics (MHD) or some thing

yes 500 degrees is pretty hot and some water is also created.

Diammonium Phosphate then....?
Hazardous Decomposition Products: Gradually loses ammonia when exposed to air at room temperature. Decomposes to ammonia and monoammonium phosphate at around 70°C (158°F). At 155°C (311°F), DAP emits phosphorus oxides, nitrogen oxides and ammonia.

there may be some water of crystallization to deal with.

Organikum - 1-4-2012 at 10:31

Reading on the dissolving metal stuff in Carruthers "Modern Methods of Organic Synthesis" lately I found that iron (collodial) and other metals are supposed to catalyse the formation of the amide.

Maybe thats of interest.

regards
/ORG

aliced25 - 29-6-2013 at 17:47

Quote: Originally posted by Sedit

I have been looking for a better means of Ammonia production because NaOH and Ammonium salt is just to finiky producing a rapid jolt of Ammonia then a decline not to mention its wet gas. However the temperatures run for that experiment seem rather high and alot of cooling would be needed on the gas before it was able to be used for this experiment.

Sedit, the better method of producing small amounts of ammonia is to make a concentrated ammonia solution (c.30%) and then heat it, pass it through a large side-arm test-tube with a small side-arm tube in the top of it full of CO2/Acetone as a cold-finger, with the side-arm of the larger tube taking the dried gas into Li/Hexane. If you use a dispersion (ie. melt the lithium on a stir plate with a stirrer) in the unsaturated hydrocarbon (as hexane can be a bitch to find), it should work better. I've attached an article on the many uses of various side-arm test-tubes from J.Chem.Ed. As well as a couple on the drying and condensation of NH3 from concentrated ammonia solutions.

PS I'm considering NH3/Unsaturated Hydrocarbons as an extraction solvent for an anhydrous AB on plant (and other) material. I strongly suspect it will be a highly reactive basic solvent, able to strip various amines from their salts, making them soluble in the NP.

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