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neutrino
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[*] posted on 11-2-2006 at 08:30


Lithium is the impossible one. The patent describes every alkali except that one.

Why is the vacuum distillation necessary? Couldn't the metal simply be allowed to coalesce into one big piece and taken out?

If you're sure that Al will work, I'll try this with shredded Al foil. First I need to find a source of samarium cobalt magnets for a high-temperature stir bar.
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[*] posted on 11-2-2006 at 10:45


Well it takes a lot of heat, so sublimating or distilling the alkali earths out under vacuum is more convienient than blasting the snot out of it at white heat or so.

Aluminum might work with the lower temperature method that's been the recent topic, but it still has that problem of forming oxides.

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[*] posted on 27-3-2006 at 20:34


The last time I worked with the NaOH + Mg --> Na + MgO + H thermite type reaction, the result was a porous and friable mix of magnesium oxide, sodium and sodium hydride that was for all intents and purposes impossible to separate. At the time I had done some thinking about the possibility of using a hydrogen scavenger to bind the H and leave the elemental sodium, however no reasonable ideas surfaced so I put the project on the back burner.

Last evening I came up with the following:
4NaOH + 4Mg + B --> 4MgO + 3Na + NaBH4

This was tested today with good results. I took pictures but didn't have time to resize tonight. Tomorrow I will post those.

Experimental

1.599 gm - Sodium hydroxide, tech
0.972 gm - Magnesium, (-325m spherical)
0.108 gm - Boron 97%, 5 micron.

The above was ground lightly in a mortar and pestle and then placed in the Parr combustion bomb and sealed. The ignition wires attached and fired. The casing became very hot, ~200c, which was allowed to cool and then opened. No pressure escaped nor was any expected. The contents removed with a lab spoon and dumped into a beaker with kerosene. Several large chunks including one of obviously metallic sodium which had agglomerated and solidified on the bottom of the bomb. The surface was sliced into with a blade revealing very shiny silver metal.

For those of you who stopped thinking about sodium 14 lines ago......
This was also tested today:
NaOH + 4Mg + H3BO3 --> NaBH4 + 4MgO

First the magnesium and hydroxide was ground in a heat dried mortar while in a clear poly bag filled with argon. Then the boric acid was added quickly the previous and all was placed in the Parr bomb that had been placed in the bag with the mortar. This must be done fast as the mix begins heating almost immediately. The wires were connected and the bomb fired.

The result was mostly light gray with white streaks. It did not behave the same as the earlier run for sodium. It seemed stable in air but effervesced nicely when placed in water. More testing certainly necessary but I think I am on to a good synth here.
I however have not decided on an easy way to separate the NaBH4 from the MgO.
Does anyone know if N-Methyl Pyrrolidone might be a suitable solvent for this purpose or have a suggestion for another?
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[*] posted on 27-3-2006 at 23:59


Couldn't the rapid heating simply be an acid-base reaction between the NaOH and the boric acid? Perhaps you could first react the sodium hydroxide and boric acid, heat to drive off the water formed, then re-powder and add the magnesium. I don't think that NaBH4 should effervesce in cold water. Whether you find a novel way to NaBH4 or simply manage to make the "sodium thermite" useful, either way it's quite interesting.



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[*] posted on 28-3-2006 at 01:35


In cold water, the decomposition of NaBH4 is almost too slow to be visible. The solution (of reagent grade NaBH4) is only a little cloudy due to extremely small hydrogen bubbles.

However, if any acid is added, especially HCl or something similarly strong, an extremely violent hydrogen evolution is observed.
I think the effervescence of your NaBH4 product in water was due to residual boric acid which decomposed part of the NaBH4.

You should try to extract your (powdered) NaBH4 reaction product with dilute cold NaOH solution and add some HCl to a sample of the liquid.
If there is any hydrogen evolution, you can be sure that NaBH4 is present.

Industrially, NaBH4 is made by reaction of trimethyl borate with NaH suspension in mineral oil. It gives an aqueous solution of 3 mol NaOH and 1 mol of NaBH4 as an intermediate product.

From this solution, the NaBH4 is isolated by extraction with isopropylamine (2- aminopropane) as solvent.

Isopropylamine could be made by reductive amination of acetone. I think that you'll be able to find something on the general procedure for reductive aminations on Rhodium.
After all, amphetamine is made by reductive amination of phenylacetone.
You'll be doing the same, just without the phenyl.




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[*] posted on 28-3-2006 at 02:37


so boron is the way to go? is boron easy to make? what i have read so far is melt down boric acid for the oxide then reduce with magnesium - so we are basically using the magnesium twice.. first to make the boron then the sodium metal

too bad boron is so friggin expensive :(

[Edited on 29-3-2006 by jimmyboy]
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[*] posted on 28-3-2006 at 18:29


Pictures as promised.

This is the magnesium and boron I used for the reaction producing mostly sodium.
http://img73.imageshack.us/img73/5230/mg1112kw.jpg
http://img529.imageshack.us/img529/2589/boronsb955vu.jpg
http://img73.imageshack.us/img73/6664/boronsb9524gd.jpg

The first shot is of a chunk of reaction mix pulled up from the bottom of the Parr cup. You can see the sodium has pooled and solidified there. I gently tapped the assembly not long after firing while still very hot.
http://img73.imageshack.us/img73/250/na10eu.jpg
I scraped the bottom with a knife exposing a fresh silvery layer.
http://img529.imageshack.us/img529/7126/na21fr.jpg
A shot from the side where you can see the magnesium oxide layers over the metallic layer.
http://img529.imageshack.us/img529/4883/na37el.jpg

Tossing selected pieces into some water.
http://img73.imageshack.us/img73/7702/naburn5yn.jpg
http://img73.imageshack.us/img73/8508/naburn20zk.jpg
http://img73.imageshack.us/img73/1827/naburn39po.jpg



This was the result from the NaBH4 tests. As you can see the material is more powdery right out of the cup. It is not brittle as the above mixes came out. The color is lighter as well.
http://img73.imageshack.us/img73/5337/nabh4mgo0tp.jpg
http://img529.imageshack.us/img529/1515/nabh4mgo20ax.jpg

As Garage Chemist suggested, I extracted the powder with a cold NaOH solution.
http://img73.imageshack.us/img73/5249/naohsolelution1ut.jpg
The beaker on the left contains 10% HCL solution, the vacuum flask on the right contains the suspect NaBH4 extract in NaOH solution.
http://img73.imageshack.us/img73/3679/10hclnabh4sol2gt.jpg
The result after they were added together.
http://img73.imageshack.us/img73/9707/mixextracthcl2wk.jpg
A good deal of gas immediately evolved smelling like hydrogen and HCL.
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[*] posted on 29-3-2006 at 11:38


Ordenblitz this is some really fine work - producing elemental sodium and NaBH4 all before supper. This is a quantum leap for MadScience! ;)

I'm intriqued by your Parr bomb calorimeter. I have never seen one before. Do you use it in support of your "real" work? What size is it? Can you show us a picture of its parts in more detail and describe their features? I looked on the Parr website but they don't give much for construction details.




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[*] posted on 29-3-2006 at 17:05


Thanks but I was only building on other peoples work.

Bromic's posts in the NaBH4 thread got me to thinking about my work on sodium that was again an expansion on his fine experiments. There may be other scavengers for the H that would work better or be easier to obtain than Boron and that should be the next direction for the evolution of this sodium process.

The big problem in making NaBH4 via this method is identifying available solvents or processes for its extraction. Isopropylamine has been suggested but I am afraid that the synthesis would be quite involved. Maybe someone with some NaBH4 could do some solvent testing.

I wanted to post more on this in the NaBH4 thread but the two methods are so interconnected in my process keeping it together seemed appropriate.

As I mentioned earlier in this thread I acquired the Parr setup on Labx for next to nothing.
http://img98.echo.cx/img98/4245/parr0ee.jpg
It is pretty old, circa 1950. I have been to Parr's site as well and really cant find anything that is similar there. I think these things fell out of favor when machines like TGA and DSC came along. Parr still makes some specialized bombs for oxygen combustion work though but they are very expensive. I'm sure there are a few more orphan units like mine, lurking in university labs looking for a good mad scientist to take it home and love.
Making your own out of SS pipe and a modified spark plug would be just as good or better even since it would have far more internal space for actually doing some production. One would only need to drill and tap out an end cap for a small diameter, long reach spark plug... think small engine etc. I would bend up the ground tang then get some fine diameter nichrome wire and wrap around the center electrode and the ground tang leaving a loop for contact to the powder. Fill, cap and spark it up!
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[*] posted on 29-3-2006 at 17:49


So thats why a bomb calorimeter is used and not an open reaction vessel....


I added 0.4g NaOH prills, 0.8g Mg filings and 0.5 g boric acid to a 100mL beaker and mixed well with a stirring rod. I then added two scoops of this mix to a testtube which was then heated over an alcohol burner. White fumes were given off then a orange fireball shot out of the tube, bounced once, then came to a stop leaving a glowing pile of something. The glowing pile(after it cooled) was added to cold water, the aqeuous part decanted and acidified. No bubbles formed. The remainder of the residue in the tube was heated for a few more seconds on the burner, then removed and allowed to semi cool. Cold water was added to the tube and some hissing was observed, it might have been some sodium in there reacting with water or it might have just been really hot still. The aqueous layer of the tube was decanted and acidified. Strong bubbling occured.:)

Seems simple enough, but now I gotta air out my basement....:P

But in any case, this is a cool method for making borohydride or sodium, I just gotta get some pipe parts I guess. What are the chances of by doing this reaction in a sealed container that it could rupture and kill me? Or is the pressure produced minimal?


[Edited on 30-3-2006 by rogue chemist]




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[*] posted on 29-3-2006 at 18:26


I would add a tee to your pipe and screw in one on those emergency blow off valves for hot water heaters just to be safe.
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[*] posted on 29-3-2006 at 18:29


It’s a very pretty yellow flame huh Rogue!
If you measure reactants properly this should be a net gasless reaction.
The O is grabbed by the Mg and I suspect pretty fast. If the NaBH4 is forming as we have assumed, the H goes to the borohydride. Maybe next time you try this in an open test tube you might try warming the contents then igniting it, say by dipping in with a steel wire heated red on a propane torch. This should keep the reactants in the tube as it would burn from the top down not the other way round. It is possible that the H would get scavenged before getting away but of this I have no idea. You may be able to determine weather a sealed chamber is necessary or not.
I have had trouble getting the contents to ignite in my bomb from the resistance wire when using magnesium of a mesh size larger than 325. Using the coarser mix you have described, one would have to preheat the chamber before firing to get the thing to start.

I have done the reaction 5 times with varying amounts of mix to better than 3/4 full by volume in the bomb and never did I hear any escaping gas when opening the bomb. One could, for only a few dollars more fit a reducer in one end of the pipe and install a small ball valve to safely vent the chamber before attempting to open it. From what I have experienced so far this is a relatively tame reaction and probably not necessary but for a few dollars of peace of mind, I would.

[Edited on 30-3-2006 by ordenblitz]
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[*] posted on 29-3-2006 at 18:47


Sounds like you have had some experiance with the flame as well eh;).

I am thinking the white fumes were just MgO being lofted by hot air in the tube or even potentially water vapour from NaOH absorbing water from the air.

When I have more time(ie not a few weeks before exam time like I am now) I would like to try this again with a tin can with a diameter a couple inches at least and a lid on top held down with a brick. The reactants could be arranged in an inverted conical pile such that in the centre of the bottom there is little reactants and along the outsides the piles grow steeper. Heating would be started in the centre of the bottom of the can. Hopefully such a design could prevent the ignition from lofting everything everywhere.

Also this is an outdoor or garage experiment next time:P

By the way, the sodium borohydride MSDSs list its decomposition as 400, what exactly does it decompose into at this temp?




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[*] posted on 29-3-2006 at 20:14


Igniting a mix from the bottom seems like a bad idea...

Don't forget that sodium is gaseous at the reaction temperature. The bomb keeps it pressurized in liquid form. That would explain the homely blob of sodium found in the first test.

So lemme see here,
NaOH + Mg = Na + MgO + 1/2 H2 = NaH + MgO. Lye reacts with magnesium to produce thermodynamically favored magnesium oxide (or likewise with aluminum), yielding sodium vapor and hydrogen gas. However, the reaction products can react further, producing hydride, useless for the goal of sodium metal.

So, you try something like,
NaOH + Mg + B = Na + MgO + BH3 (assuming boron is the stronger 'hydriphile' so to speak). Which works pretty well, but requires boron, so you might use more common materials:
NaOH + Mg + B(OH)3 = Na + MgO + BH3 [unbalanced], but there's too much oxygen and hydrogen to go from boric acid directly: NaOH + 4Mg + B(OH)3 = 4MgO + NaBH4. Same results as with plain lye, except boron has been reduced too. Of course if you want sodium hydride or borohydride, you can use either method, (somehow) extract them from the magnesium oxide, and smile.
One good way to reduce the hydrogen is to dehydrate things:
NaOH + 4Mg + B2O3 = 4MgO + 2B + Na + H, my this has some strong potential, let me rewrite that,
6NaOH + 9Mg + B2O3 = 9MgO + 6Na + 2BH3
Wonderful, it doesn't even need much boric oxide it would seem!

You could also start with sodium (or potassium? :D ) borate, probably a fused product (as with the boric anhydride) to ensure it is anhydrous.

Lemme see, borax is sodium "tetra"borate, which really means Na2B4O7 IIRC, so you'd have to add a good bit of lye even to that already.

Heyyy, when you melt things, gaseous anions tend to go away.. (at least, when glasses are analyzed, you only see oxides listed). You might be able to melt borax (or boric acid) with sodium carbonate, get it good and hot until it stops bubbling, then pour the molten glass into a bucket of mineral oil (obviously, something other than water) to quench it while remaining anhydrous, then wash with solvent, grind and you should have very little hydrogen whatsoever! Why, then you'd get a bunch of plain boron because there would be no hydrogen... oh, so you could reduce the boron present... ah, but the assumption is that sodium carbonate or hydroxide won't turn to sodium oxide on its own, so some must remain...?

Tim

[Edited on 3-30-2006 by 12AX7]




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[*] posted on 30-3-2006 at 12:07


boron oxide would be alot easier/cheaper if it worked - and not near as expensive - just add extra magnesium

hmm maybe this can be applied to phosphorus as well?

[Edited on 31-3-2006 by jimmyboy]
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[*] posted on 1-4-2006 at 20:37


2g NaOH, 4.8g of Mg filings and 3 g boric acid were mixed in a soup can. The can was tilted adn then returned to normal such that the reactants were slanted inside the can. The lit was placed on the can and secured with a single strip of ducttape. I could not find a brick so I used a slice of railway track on top of the can. A big plastic flower pot was placed over everything and the can heated with a propane torch on the side with the least reactants. After a few seconds of heating there was a hissing sound and the entire apparatus glowed orange...smoke excaped the can and floated up to the garage ceiling. I extinguished the flaming ducttape with some water and opened the garage door....that smoke can't be healthy.... I left and came back 5 min later when the garage was mostly vented and the can had cooled down. The big flower pot was white inside from the white smoke. 400mL of ice/wate was added to the can slowly...there was a small orange explosion so some sodium must have been formed. A sample of the water extract was acidified and strongly bubbled. It is being filtered currently of all the other crud in there. Is there any way that an aqueous solution of borohydride coul d be used to make isoproplamine via reductive amination of acetone, the places I saw were unclear on this?

EDIT: Weird...the aqueous filtrate does not bubble when acid is added, but the residue in the filter does...I thought borohydride was soluble in water? And it worked for you ordenblitz...weird...

EDIT2: Is it possible that boron could form? I got a good deal of black insoluble flakes...

[Edited on 2-4-2006 by rogue chemist]

borohydride.JPG - 23kB




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[*] posted on 2-4-2006 at 06:45


In response to the question about thermodynamics, the lattice energy of aluminum/magnesium oxide drives the reaction forward.

4Al + 6NaOH --> 2Al<sub>2</sub>O<sub>3</sub> + 6Na + 3H<sub>2</sub>

ΔH = -797.9 kJ




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[*] posted on 2-4-2006 at 09:12


After the reaction, I dumped the contents of the parr cup into a beaker. To this I added a 5% NaOH solution. I also saw a few yellow flashes as well, obviously from some sodium in the mix but there wasn't all that much. I stirred a bit then removed the solution through a 1.3 micron filter. When I added the HCL solution, as you can see from the pictures frothed quite nicely.

It makes sense that the sealed chamber is necessary since after all, the very reactive products of the in initial reaction are not going to politely hang around to form the borohydride without some encouragement from confinement.

I do not know much of anything about how NaBH4 is used in mad chemistry. The problem it seems is how to separate magnesium oxide and the borohydride. Is it possible that the NaBH4/MgO complex could be added to any reductive reaction and the MgO not substantially interfere?
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[*] posted on 4-4-2006 at 05:36
I was wondering something


What if you took lithium and dropped it into a solution of NaCl and some solvent other than water. Something that the Lithium and Natrium were unreactive to? I'm not sure what type of solvent would work though. Could something like diethyl ether be unreactive to the metals, or would this just react horribly. Maybe sodium could be precipitated this way somehow.

[Edited on 4/4/2006 by lacrima97]
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[*] posted on 4-4-2006 at 20:48


NaCl is insoluble in ether. You need something better, like uh, isn't it vaguely soluble in pyridine or somethin? Problem is reactivity tends to follow polarity, probably your best solvent (water) is unfortunately the worst!

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[*] posted on 4-4-2006 at 21:03


True, there are not many solvents that would be able to preform this trick, anhydrous liquid ammonia would be a possibility but only 3 grams of NaCl is soluble in 100 g of the stuff so you'd need a lot of ammonia though lithium chloride might be significantly more soluble allowing more the NaCl to be added slowly and in ever increasing amounts. Provided you had the proper reaction vessel you could also make it the same way potassium is made from sodium, via distillation. Distilling sodium from a mixture of sodium chloride and lithium metal would be one fun route. But why waste lithium metal, it burns prettier then sodium and has a better electropositive value. Still, I guess it is one possible viable way to sodium.



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[*] posted on 6-4-2006 at 04:35


Sodium metal is soluable in liquid ammonia, as are most of the alkali metals. The solution eventually decomposes to hydrogen and the hydride but for a while, its blue.

Also the 'lithium is stronger' is not really true, the unusual ordering only applies to the ions in water.
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[*] posted on 6-4-2006 at 07:16


It decomposes to the amide and hydrogen via:

2Na + 2NH<sub>3</sub> ---> 2NaNH<sub>2</sub> + H<sub>2</sub>

However if your ammonia is very pure this decomposition is retarded and one can attain the alkali metal by evaporation of the ammonia, first getting a golden solution and finally the metal itself. Still, you need good ammonia and it takes lithium metal, I know the potentials are only good for aqueous solutions but it's still something to go by.




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[*] posted on 12-9-2006 at 21:21


thats an amazing borohydride synth!

could lead be used instead of Hg in brine electrolysis?
i know that Na-Pb alloy is used to dry ether, could it be used as sodium amalgam is as well?
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[*] posted on 13-9-2006 at 06:35


I doubt it, it's notoriously difficult to get liquid water together with liquid lead (or bismuth).

It would happen in an autoclave under intense pressure. I doubt the little sodium content would not react under that kind of heat.

Easier, but still under pressure, would be the tin-bismuth eutectic. If your goal is distilling sodium from it, this would be easy as tin and bismuth have high boiling points.

You might take an alloy such as Wood's metal, which melts on par with hot water, but there are volatile elements like cadmium and zinc which would distill off with the sodium, if your goal is seperating sodium. If you just want amalgam, the other metals may cause trouble.

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