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Author: Subject: unconventional sodium
not_important
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[*] posted on 13-9-2006 at 07:10


Pb-Na might work in place of Hg-Na if it were in the form of a fine powder, otherwise much of the sodium content is not accessable for reaction.

Indium alloys have low melting points, gallium is liquid at slightly above STP, Galinstan is used as a mercury replacement. Those metals all have high boiling points, unfortunately indium and gallium also have high costs. Like mercury and sodium, solid intermetallic compounds are formed if the sodium content is great enough.
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tupence_hapeny
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[*] posted on 26-3-2007 at 00:02


For the diehards wanting to get to sodium metal, I enclose the following patent, which describes the use of a divided cell (only using fibreglass matting) and the use of an ionic liquid (methanesulfonyl chloride) to realise sodium metal <200C:

http://www.freepatentsonline.com/6787019.html

I realise methanesulfonyl chloride may be expensive and hard to acquire, so I hereby provide access to the means of producing the same:

http://www.freepatentsonline.com/4997535.html

Uses methane, sulfur dioxide and chlorine gasses, in a glass tube, under UV light (specifically stating that low vapour mercury lamps work).

BTW, Heres another method by which to get to the borohydride:

http://www.iop.org/EJ/abstract/0953-2048/17/12/L01/

In this synthesis they used magnesium powder and powdered elemental boron, in an aluminium container in the microwave to make magnesium diboride, which can be turned into x.BH4 (where x=Na, Ca, K, (NH4)2, etc.) by introducing the diboride to a concentrated solution of x in water (yeilds suck according to the patent), releasing copious quantities of hydrogen and some x.BH4. I assume anything else in the reaction vessel would be reduced however.... (In which case, fuck the x.BH4):cool:

I read somewhere that boron can be made from boric acid via reduction (at some considerable temp's) with carbon (although be careful, boron carbide will form if you go too hot) and also by electrolysis (presumably of the boron-trihalides - also perhaps BF (DON'T DO THAT ONE ANYWHERE NEAR ME)).

[Edited on 26-3-2007 by tupence_hapeny]




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tupence_hapeny
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[*] posted on 27-3-2007 at 21:28


Yup, that was a fuckup.... LNG is what you need.... Town gas as it called in the patent

EDITED FROM HERE ON 4 April 2007 by tupence-hapeny

The other ionic liquid (the phosphoryl one) is an absolute bitch to make, given that the only synthesis that I could find requires (as a starting material) PCl5. Anyhow, here is the synthesis:

http://answers.google.com/answers/threadview?id=433962

From what I have read, white phosphorus is reacted with HCl (not for the fainthearted) and then reacted further with Chlorine to give the PCl5:o As the easiest low temp route to phosphorus is via red phosphorus (~290C it supposedly distills over - let me suggest an inert atmosphere) this would appear to be unworkable and unfeasible.

Unfortunately, the low temp electrolysis patent does not detail high yeilds (or high current attempts) with the methanesulfonyl chloride. However, here is a more detailed version of the original patent (unsure if it is a co-application or a later application) which seems similar but includes better details:

http://www.freepatentsonline.com/20030094379.html

Here is another method of making methanesulfonic acid:

http://www.freepatentsonline.com/6207025.html

[Edited on 5-4-2007 by tupence_hapeny]




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[*] posted on 26-6-2007 at 21:49


Some quick questions on sodium electrolysis, pertaining to solvents. I'm asking here because I haven't found answers through the usual channels (Google, wikipedia, etc)

If I were to try to electrolyze a sodium salt in the following, what (theoretically) would happen?
An alcohol (such as ethanol or isopropanol)
A ketone, such as acetone.

I'd assume that sodium would react energetically with any of these these. I just don't know what the products would be.
Electrolytically, though, without water, I'd imagine the usual production of H2 and NaOH at the cathode (from using water as a solvent) wouldn't happen. Something else would - is there any chance that sodium could be made this way? And if not metallic sodium, are these likely to produce anything else that may be useful or interesting?




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[*] posted on 27-6-2007 at 17:45


Alkali metals are extraordinarily basic.

Ketone: probably enol form, then attack at the hydroxyl. "Acetonate"?
Alcohols: the hydroxyl is attacked, giving off H2 and an alkoxide salt.

Alcohols are much less acidic than water (Ka in the 10^-20s range, vs. 10^-14 for water), so there are fewer protons (H+) around to attack and reduce (2 Na + 2 H+ = H2(g) + 2 Na+) and it goes slower.

Alkanes (like mineral oil) are very, very non-acidic (typical methyl or methylene group Ka < 10^-40 or so, isn't it?), so they aren't attacked and are safe to store alkalies under.

If you don't know what Ka is, it's the equilbrium between something complete and something that lost a proton. It mostly applies to water (i.e., H2O <--> H+ + OH-), but anything with hydrogen can lose hydrogen in an equilibrium, it's just a matter of how much prying is required to yank it off.

Another important issue is the solubility of the salt in the solvent. Salt typically has very little solubility in anything other than water and a few other solvents that happen to be reasonably ionic. Even if the sodium can be formed without reacting with the solvent, there may not be enough ions in solution to reduce at an acceptable rate. And that's why a salt melt is so much preferred: it's ionic and doesn't react with the metal.

Tim




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[*] posted on 15-7-2007 at 18:54


more of my darned curiosity....

what about electrolyzing a sodium salt dissolved in liquid anhydrous ammonia? I know that may be out of the reach of the average experimenter, but does anyone else think it would work?
I have a thought that the chlorine may form chloramine though instead of bubbling out of the solution, which may be problematic, though, if there's an anion to attach to sodium that wouldn't react with the ammonia when liberated (electrons removed, and it recombines, as the halogens do.), we might have yet another way to get sodium...




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[*] posted on 15-7-2007 at 20:50


Ah, so instead of sodium combining with the solvent (which does happen, just really slowly), the chlorine does? Could very well be. :P

Electrolysis of sodium amide is probably straightforward. But that doesn't help any. :D Hydroxide probably works, but that can be done somewhat easier in its own melt.

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[*] posted on 16-7-2007 at 15:38


all this talk of melting sodium chloride at home is pointless - you would need an inert gas and very high temps to get the sodium under those conditions --- you could just buy the sodium instead at that price -- the objective is to get small amounts at a lab scale
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[*] posted on 17-7-2007 at 14:50


Well, no, not if you lead it off as in a Downs cell.

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[*] posted on 14-10-2007 at 14:35


I have a small furnace that can heat to around 1200 C and since sodium boils at around 900 C could I distil it? As far as I know it was made before by mixing charcoal and sodium carbonate and heating the mixture to around 1000 C. I was thinking about using a steel pipe closed at one end and an elbow at the other end and then a another pipe that would lead the molten sodium in mineral oil. The pipe would have to be long enough so that the sodium cools enugh. Would it be possible?



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[*] posted on 14-10-2007 at 15:01


Quote:
Originally posted by Zinc
As far as I know it was made before by mixing charcoal and sodium carbonate and heating the mixture to around 1000 C.


Castner's original process involved heating NaOH with a mixture of carbon and iron at a temperature of 1000 oC.

6NaOH + 2C ---> 2Na + 3H2 + 2Na2CO3

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[*] posted on 15-10-2007 at 08:20


Zinc,

It may not be impossible, but sounds like a process that will require a good investment in time and equipment. I would invest that in an electrolictical cell like len1 has built.

http://www.sciencemadness.org/talk/viewthread.php?tid=2103&a...
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[*] posted on 15-10-2007 at 10:22


Quote:
Originally posted by Tacho
I would invest that in an electrolictical cell like len1 has built.


It doesn't need to be as sophisticated as len1's cell, if you only want small quantities of sodium.
All you need is some NaOH, a steel dish, some bits of wire, a gas burner and a battery or power supply.

http://www.sas.org/E-Bulletin/2001-10-05/chem/column.html

NOTE: Pay attention to the stability of the apparatus (I had a bad accident, trying this as a teenager). See my comments on len1's apparatus in the other thread. Do this outside, but not if rain is forecast! Be careful not to short circuit the electrodes whilst they are beneath the molten NaOH. Wear protective clothing and glasses.

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[*] posted on 15-10-2007 at 10:26


The distilation seems a lot simpler to me. If it could be done more than 100 years ago it can be done today.



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[*] posted on 31-10-2007 at 14:45


I just read this thread so please pardon the late comment!

Bromic Acid, I was very sorry to see your experimentation with US patent 4725311 not continue. I think that it is a rather novel method and it involves some excellent explosion possibilities. Fascinating!

You mentioned that you wanted a procedure to make t-butyl alcohol. We've been covering oxidation of alcohols in class so I thought I'd give it a swing just for fun.

Make some Jones Reagent, which is a solution of chromic acid and 8N sulfuric acid. Fill a titration burette with isopropyl alcohol and slowly add the isopropyl alcohol until the green chromium salt sludge stops precipitating out.

Since the initial alcohol was a second degree alcohol, it will form acetone. This is probably a good time to mention that you could also just start with acetone, but for education sake!

Make some chloromethane by reacting methane and chlorine gas at elevated temperatures. Hell you might even throw in a light source so you can say you employed photochemistry too. Chill in a freezer to condense the alkyl halide. Dissolve in acetone and poor over magnesium metal. A grignard reaction will hopefully take place generating a methyl radical that will attack the acetone and form (CH3)3CO- which will pick up a hydrogen to produce t-butyl alcohol.

After several months of labor, I’m sure you could pump out tens of milliliters! Is this an accurate reaction?

There’s probably a more efficient method that more learned members can extrapolate from the fundamental procedure that I have posed here.

CH3COHCH3 + CrO4H2 --> CH3CCH3=O
CH3Br + Mg --> CH3* + MgBr2
CH3CCH3=O + CH3* --> (CH3)3CO-
(CH3)3CO + H+ --> (CH3)3COH

Anyway, please continue with the US patent 4725311 experimentation so that I may experiment vicariously through you.
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[*] posted on 31-10-2007 at 17:52


I'd love to do more with the patent work. I had several modifications that I wanted to test, the 2-propanol that I was using seemed to be working okay, but sadly I won't be getting to any of them any time soon. I moved from my previous home to an apartment building and I honestly don't think my neighbors want me heating mixtures of magnesium metal with alkali hydroxides under flammable solvents to high temperatures.

Oh well, one of these days hopefully I will get around to having a place of my own but for now someone else will have to resort to this line of experimentation, thanks for your interest though.




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[*] posted on 31-10-2007 at 17:59


Im planning on checking this method shortly, to see how it compares with my earlier work on electrochemical Na cell as a method for getting Na. I have one problem left to solve before I can try the method: I need a non-carcinogenic solvent with a density greater than 0.98 and inert to sodium. Unfortunately dioxane and CCl4 fit the bill but are carcinogenic.
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[*] posted on 31-10-2007 at 19:31


You DON'T want to mix chlorinated solvents with alkali metals. :o

Sodium-potassium alloy in an ampoule inside a test tube of chloroform is a classic "flash-bang" (and shrapnel) device.

Why do you want your sodium to float on the surface of your solvent?

If that's really what you want, propylene carbonate might fit the bill.

[Edited on 10-31-2007 by Eclectic]
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[*] posted on 31-10-2007 at 20:12


Sorry you are right CCl4 reacts with sodium if initiated. I had thought there would not be a reaction because the C-Cl bonds are highly covalent. That there would be a reaction with chloroform is clear. In any case they are both carcinogenic.

The way this solvent is used to separate sodium is described in the patent.

Your suggestion of propylene carbonate seems a very good one. Its a non-carconogenic ester and has a high sp. Thanks

PS I have now found a literature reference where Na catalyst is used in the presence of propylene carbonate, which means the later must be unreactive to Na. So that solves that problem. It seems to be a bit trickier to make than dioxane (passing urea and propylene glycol over an iron-zinc catalyst) nut is still a better solution than dealing with dioxane i think.

[Edited on 1-11-2007 by len1]

[Edited on 1-11-2007 by len1]
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[*] posted on 1-11-2007 at 13:13


Why do you need a solvent that sodium will float on? Just melt and filter it, you'll leave your MgO far behind.



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[*] posted on 1-11-2007 at 17:57


Quote:
Originally posted by BromicAcid
Why do you need a solvent that sodium will float on? Just melt and filter it, you'll leave your MgO far behind.


Quote:
Originally posted by BromicAcid
No potassium was recovered. My reasons for failing, the isopropyl alcohol has such a low boiling point and I was running the reaction at the boiling point of kerosene, the patent recomends t-butanol which I did not have, this was part of my problem. Also filtration was not the best option, in theory the kerosene should have been distilled off and dioxane added which is more dense then the potassium and causes it to rise to the surface, then it is either skimmed off or filtered off, the solid all fell out at once so I think that stopped efficent filtering from happening.



Weren't you the one who posted the excerpt above? So why are you asking me this question?

In reality the surface tension to density ratio of liquid Na is such that substantial pressure differential is required for it to pass through a filter with even large pores. This pressure diffrenetial increases with decreasing pore size so that to eliminate microscopic MgO particles a substantial pressure difference at 100C plus and inert atmosphere is required. Its not for nothing that the inventor in the patent chose a density separation stage - its much simpler.

[Edited on 2-11-2007 by len1]
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[*] posted on 2-11-2007 at 13:24


Yeah, forgot about that post. Although now I am more familiar with pressure filtration apparatuses that would be better suited to this were I could use nitrogen to drive the filtration and not really worry about air contact.



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[*] posted on 13-12-2007 at 17:02


I have been looking at methods for synthesising propylene carbonate, and havent been spoilt for choise. Does anyone know of a method, have a reference which includes methods for isolating and purifying the product? thanks Len
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[*] posted on 22-12-2007 at 06:02


I have been fascinated by US patent 4725311 for a while now. KOH and CsOH can be reduced by Mg turnings to the metal at no more than 200C in a common organic solvent. It seems sweet, a less reactive metal yielding much more reactive ones at low temperature. Certainly other such reactions are known, for example with Zr, though this requires a vacuum and much higher T. The patent operates under much simpler conditions.

This reaction certainly seems well within the realms of reality, if one looks at the enthalpy of formation

KOH + Mg -> MgO + K + 1/2H2

425kJ/Mol - 603kJ/mol = -178kJ/mol

because of the high enthalpy of formation of MgO

Its one of those reactions which if I hadent been told has been proven, I would never have believed. The inert solvent does not promote intimate contact between Mg and KOH, two reactants which have a significant kinetic hinderance to reacting, and dont react if their mixture is heated to 200C, despite the negative enthalpy. The writeup in the patent seemed to be very reaslistic though. I went ahead and purchased Shellsol D 70 and tert-Butyl alcohol, two reactants which the patent requires which I dont normally have.

The reaction was carried out exactly as in the patent. Fresh Mg turnings stirred with ground KOH powder in Shellsol D 70 for 4 hours at its bp (200C), under an argon atmosphere. Tert-butyl alcohol was added in shellsol at 200 bp, slowly.

Result - no reaction, with two different sources of Mg. No violent evolution of H2 at 100-130 due to 'reaction with H2O', no K formed, just the Mg turnings.

Why? Everything appears to have been as in the patent. The possibilities are

1) The Mg was not in a reactive enough form

2) The reaction vessel was contaminated

3) The patent is bogus. Written on the basis of calculations on paper rather than work in the lab to help gain the author an authored patent.

Another reaction with Mg that is very temperemental is the Grignard reaction. But that is due to its sensitivity to the dryness of the reactants. Here the Mg is purported to dry them. Indeed the reaction is 'violent' if not dry enough. One could use pyrophoric Mg, but then the patent clearly states 'turnings' or chips, nothing about special preparation.

Of course the thermodynamics is favourable, but the following very similar reaction is even more so:

HOH + Mg -> MgO + H2

+295kJ/mol - 603kJ/mol

yet it doesnt go in cold water, and is very slow (forming the hydroxide) in hot water. And in the inert solvent the contact between the reagents is much worse than in boiling water. Add to this that this reaction is much less favoured with KOH, whose heat of formation is much higher (430kJ/mol) than that of water.

Having been led to consider 3) seriously Im now convinced this patent is a fake. Look at the H2 thats claimed to be evolved. 278mmol. Hardly measured so accurately -but OK let this pass. It agrees exactly with the excess Mg added (also 278mmol). That is funny because the Mg should go at half the mmol to the H2 from the formula.


My explanation is that the patent author took the above reaction seriously, that is he thought

H2O + Mg -> MgO + H2

whereas in water we have

2H2O + Mg -> Mg(OH)2 + H2

dehydration of the hydroxide to the oxide takes place at 580C dry, and would hardly occur in the hydrocarbon at 200C.

Clearly the excess Mg and KOH weight was calculated backwards on the basis of the H2 evolution, also calculated. The reaction was never performed.

Now note that the H2 evolved subsequent to the reaction with H2O at 130 degrees is 460mmol. Thats because 460*2+80=1000. SO THIS AMOUNT WAS NEVER MEASURED IT WAS CALCULATED. Indeed it would have been hard to measure if the vessel was being flushed with argon.

It is also strange (though not perhaps fataly) that flushing with Ar is mentioned only after the reaction is finished and has cooled below 70C. Exactly when its not needed, because the K is at the bottom of the hydrocarbon bath, indeed how it would usually be stored.

My conclusion is that the patent is bogus, and its author a lier. It is indeed a pity that in professional publications one can not use that term. Stark language might stop such people. Untruths can cost people time and money. I have little time for that sort of thing.


[Edited on 22-12-2007 by len1]
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[*] posted on 22-12-2007 at 10:00


Looking at the legal status also suggests it may bogus.

The patent itself suggests actual labwork, it is true. There stirrer may have been agressive, be interesting to try a small run under ultersound as that often can kickstart Grignards. The H2 supposedly being formed might have served to protect the product, OTOH H2 and the alkali metals react at the upper end of the suggested temperature range.
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