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Author: Subject: Preparation of Lithium Ammonia complexes in EtO2.(Lithium Bronze)
Sedit
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[*] posted on 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.


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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:P) 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]





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[*] posted on 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:D





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[*] posted on 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?





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[*] posted on 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.
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[*] posted on 7-1-2011 at 10:40


Quote: Originally posted by aonomus  
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
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[*] posted on 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.





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[*] posted on 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:




IMG_5052a.JPG - 26kB IMG_5053a.JPG - 22kB
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[*] posted on 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.
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[*] posted on 26-2-2011 at 08:05


Impressive and beautiful. Liquid bronze!
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[*] posted on 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:D





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[*] posted on 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!

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[*] posted on 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.




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[*] posted on 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?
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[*] posted on 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)
This file has been downloaded 1326 times

-_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]
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[*] posted on 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.





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[*] posted on 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.


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[*] posted on 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.

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/ORG
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[*] posted on 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.


Attachment: McDevitt.The.Many.Uses.of.a.Seven.Inch.Sidearm.Test.Tube.pdf (1.7MB)
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Attachment: Ilich.etal.Solvated.Electrons.in.Organic.Chemistry.Laboratory.pdf (680kB)
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Attachment: Ilich.Solvated.Electrons.in.Organic.Chemistry.Laboratory.Supporting.Information.pdf (493kB)
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