Sciencemadness Discussion Board

unconventional sodium

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not_important - 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.

tupence_hapeny - 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]

tupence_hapeny - 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]

indigofuzzy - 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?

12AX7 - 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

indigofuzzy - 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...

12AX7 - 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.

Tim

jimmyboy - 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

12AX7 - 17-7-2007 at 14:50

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

Tim

Zinc - 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?

Xenoid - 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

Regards, Xenoid

Tacho - 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...

Xenoid - 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.

Regards, Xenoid

Zinc - 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.

Centimeter - 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.

BromicAcid - 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.

len1 - 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.

Eclectic - 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]

len1 - 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]

BromicAcid - 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.

len1 - 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]

BromicAcid - 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.

len1 - 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

len1 - 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]

not_important - 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.

Magpie - 22-12-2007 at 12:26

I'm sorry I can't comment on the actual synthesis itself.

But I do have opinions on patents in general: I don't trust them. When attempting to make KOCN I had collected 4 or 5 patents, all claiming an advantage over the current method. Sometimes they even conflicted with each other. None of them worked for me.

My success with very early (~1900) journal articles is not very good either (re: malonic acid).

[Edited on by Magpie]

[Edited on by Magpie]

len1 - 22-12-2007 at 14:30

Yes thats what I thought, the H2 formed is meant to protect the metal, just as in the electrolytic method of K production from KOH. But in this case the K is heavier than the bath, so as soon as the reaction starts subsiding you can turn the stirrer off and the K, if it actually existed, would sink to the bottom. There would be no need for Ar at the end, if it wasnt required in the beginning.

These little details differentiate real work from fantasy. I noticed this oddity, but didnt put too much weight on it before I tried the experiment. Now I see its true import.

I didnt know about the legal status of this patent. How does one check that? To be honest I feel emotionally gutted, having been taken for a ride like that. I have been prepairing for this experiment for over a year.

I really dont believe anything can save this patent. Grignards are hard to start because of moisture and a protective MgO layer over the Mg. I believe the ultrasound is used to form a breach in this layer and expose fresh Mg to the organic halide. This reaction IS meant to happen with water present, the Mg acts as a dehydrating agent and supposedly violently. Well the Mg did react with water - it became a dull colour, there was no problem with initiation. Its just that it didnt react with the KOH. i guess I should have known - try heating Mg with KOH powder to 200C - they dont react. So stirred in a hydrocarbon its likely to be even more so. The fact that the supposedly measured quantities in the patent were actually calculated, and the elementary error that the reaction goes to MgO resulting in an impossible 'measurement' of H2 released in the 100-130 C phase, the lack of any K in my experiments, conducted exactly as in the patent, and the sluggish as one would expect, rather than 'violent' reaction of the Mg with water solidifies my belief that its all rubbish. The guy never did the exp.

[Edited on 23-12-2007 by len1]

reagents.JPG - 80kB

not_important - 22-12-2007 at 17:18

espacenet patent searches give an intro page with tabs for the full doc and legal status. This one ends with
Quote:

PRS Date : 2000/04/25
PRS Code : FP
Code Expl.: - EXPIRED DUE TO FAILURE TO PAY MAINTENANCE FEE
EFFECTIVE DATE: 20000216

if it was practical I'd expect they'd keep it alive. Combine that with an absence of reviewed publication and you have to wonder.

http://v3.espacenet.com/legal?DB=EPODOC&IDX=US4725311&am...

len1 - 22-12-2007 at 17:44

Thanks for that not important. I must say the owners look convincing, a german enterprise and several german scientists - you generally dont expect that sort of thing from them. That the results regarding H2 evolution are 'cooked', is without doubt, but I guess out of respect for the type of individuals involved, I will give one more chance to the process. Maybe it did work, but for one reason or another the did not measure the H2 evolved, but decided to cook it up for solidity.

I shall repeat the experiment, but will keep it going longer than their 4 hours, will use a hotter plate and more violent aggitation - although I dont want to break the flask. I havent got a ready means of applying ultrasound inside a rb flux thats heated and stirred - but the addition of the alcohol, I believe fulfills this roll - thats after all what the patent claims.

I will report the results shortly - but realistically I dont expect anything but confirmation of what Ive already written - and I will be able to leave the topic with a clear conscience. Len

len1 - 23-12-2007 at 00:43

This time I ran the experiment for 6 hours. Changes from the previous runs were as follows

1) The heat output was upped so the reagents were at a vigorous boil at 200C, to achieve this the RB flask was placed in a sand bath maintained at 240C.

2) Freshly cut Mg turnings, from a third source were used and thrown straight into the hot solution - there was no time for the Gringard stopping MgO to form

3) The mixture was stirred very vigorously from a special fast overhead motor

4) As before I did not use my good ground-glass multi-necked flasks, KOH at high T etches glass, and I dont want them stained, especially on an experiment that appears dodgy. I used a cheap flask with a rubber stopper - there was no perceptible leak or interaction with the rubber at 200C.

Result - some white percipitate, which would have to be Mg(OH)2 formed and Mg darkened from interaction with water. No K formed as evidenced by throwing contents of flask into water. Zilch in the language of the patent inventors.

I am now very confiedent this patent is rubbish and am sorry I wasted time on it. Reasons are:

1) No K formed in any experiments despite following patent closely, three different sources of Mg, the last freshy cut

2) The patent clearly 'cooks' the amount of H2 evolved. In water Mg + 2H2O -> Mg(OH)2 + H2, not Mg + H2O -> MgO + H2. SO 278mmol H2O releases 139mmol H2 not the 278 that the patent quotes. The hydroxide doesnt dehydrate at 200C. The very accurate wrong 'measurement' is in exact accord with the incorrect theoretical formula, which shows that it was never measured. The same goes for the H2 released subsequently - construed to add up exactly to 1000mmol, the amount of pure KOH used. If they cooked one part of their patent, they are likely to have cooked the rest.

3) This type of reaction was not really expected. KOH and Mg dont react at 200C the KOH form only an emulsion in the hydrocarbon and as such its interaction with the Mg in this medium is not expected to be better than in the solid state.

In short 'BUSTED'. I feel unclean for having participated in this rort. I look forward to doing something more interesting.

apparatus.JPG - 29kB

Fleaker - 27-12-2007 at 09:10

Pity. Len, I think you would have better luck distilling potassium from KCl using some of your Na to displace it. That's at least tried and true.

len1 - 27-12-2007 at 14:03

Thanks. The whole point in this for me was to witness what I thought was an amazing reaction. Until you look closely the patent seems really genuine, at worst I thought Ill get a low yield. I am not sure what the patent authors wanted to achieve by faking somthing that is so easy to check. Len

BromicAcid - 27-12-2007 at 14:34

When I made my run at this I didn't have any special preparation of my reagents. My kerosene was taken straight from the gas station, my isopropyl alcohol from the drug store, and my Mg out of a plastic bag that it came in from the pyro people when I bought it. My KOH was in large prills and taken straight from the container, it was originally sold for soap making. No clue why I got such glaringly different results.

I had a strong effervescence throughout the procedure and I ran it at notably lower temperatures. Granted there is the possibility of water contamination but I would not think that it would be to such great extent as to dissolve as much Mg as dissolved.

Because I have not posted them elsewhere (or maybe I have) here are more pictures from my attempt.

http://www.destructve.com/sciencemadness/IMAGE011.JPG
http://www.destructve.com/sciencemadness/IMAGE016.JPG
http://www.destructve.com/sciencemadness/IMAGE018.JPG
http://www.destructve.com/sciencemadness/IMAGE021.JPG

They are in chronological order, with the first picture taken within 30 minutes of beginning the heating and the second was taken after 3 or 4 hours IIRC. The third picture was taken after the mixture had been stoppered and allowed to sit approx. 15 hours and you can see the strange color change that I attributed to any K working around in there chewing things up and reacting. The final picture is form when I filtered the mixture and decided to dump some of the solid material into some water, you can still see bubbles from the somewhat vigorous effervescence along the rim of the water in the container.

Again, I don't know what happened differently, maybe your reagents were a bit too pure or maybe t-butanol just works like crap. My main reason for posting here though is just to show exactly what I observed. With the validity of this patent being called into question I wanted to show that the results that I posted about were not exaggerations and were what I actually saw, no matter their origin or how much of a fluke they may have been.

len1 - 27-12-2007 at 15:19

I dont think your results are different. A white precipitate is generated from the reaction of Mg with H2O, that would occur in pure hot water as well. The reaction generates H2 although at half the rate to what the patent authors 'measured'. Thats the effervesence. But the point of the patent is not to form magnesium hydroxide.

My request to you or anyone else who thinks this reaction works, is

please show the potassium.

Anything else is pure speculation. Len

[Edited on 27-12-2007 by len1]

crazyboy - 18-2-2008 at 20:13

OK so I just repeated BromicAcid's experiment detailed on page 3 except I only used half the hydroxide and magnesium and I substituted potassium hydroxide for sodium hydroxide. I also used a bean can with a 1/4" hole drilled in the top; I inserted a short piece of copper pipe into this hole and plugged it with steel wool. (I did this to limit oxygen to the reaction chamber without causing too much pressure to build up. then I heated the can with a blowtorch after about 30 seconds smoke began to come out of the hole I continued heating for 20 or so more seconds until all smoke had stopped. I removed the copper pipe and began to use a can opener to remove the lid. When the lid was only partially off I could tell there was no fire in the can when it was removed the contents lit with a bright yellow glow it wasn't as bright as magnesium and it went on for about 50 seconds. Then it died down and I let it cool. I came back 5 minutes later and poured in some cold water the mass relit with brilliant yellow flame then was extinguished. Did I produce pure sodium? Or was it just magnesium burning in the presence of a sodium salt with little oxygen that caused the slow yellow flame? I plan on repeating this with more Mg and NaOH and pouring xylene in the can after it is cooled instead of opening it to the air does this sound like a good idea?

crazyboy - 18-3-2008 at 18:18

Well I believe I just made impure elemental sodium metal. I heated 15g Mg and 33g NaOH in a sealed soup can in an argon atmosphere with a blowtorch for about a minute or two. Then I let it cool and poured in xylene (kind of screwed everything up) the chunk was broken up with a screwdriver and the chunks and soot where found to release hydrogen gas and combust when dropped into water.

The xylene screwed the whole thing up because it was oily and wouldn't get off the chunks so when they were tossed in they only liberated hydrogen and only a few actually burst into flames. I know this was the xylenes fault because they burnt with a sooty flame and when they where heated with a lighter they burned sootily but when the burnt piece was tossed back into the water it burned much better then the ones still coated with xylene.


Here is some short demos of my homemade sodium: http://s179.photobucket.com/albums/w318/crazyboy25/?action=v...

It might seem to burn with a oily flame not like that of sodium this is due to xylene not the sodium.

[Edited on 18-3-2008 by crazyboy]

kilowatt - 19-3-2008 at 19:50

Nice work. Next you could try distilling it under argon to see if you can get any pure sodium metal out. Sodium boils at 883°C. Make sure everything is thoroughly purged of air of course. Be sure to allow venting, since if you got any sodium hydride that will want to decompose to sodium and hydrogen. Maybe this process could even be coupled with just such a distillation, so it is a one step thing.

[Edited on 19-3-2008 by kilowatt]

crazyboy - 19-3-2008 at 21:59

Thanks! How would you suggest doing that? I mean an all glass setup would corrode with the molten NaOH and if it didn't I would think the heat would be too much for the glass. Perhaps a metal can with an all copper condenser? Perhaps even a feed of argon into the can at all times? I'm not trying to shoot down your ideas i'm just looking for practical ideas.

kilowatt - 20-3-2008 at 02:07

I was thinking an iron reflux still made of threaded pipe fittings, with stainless steel scrubbers for column packing. The magnesium oxide smoke would come right over with a simple retort. I would definitely keep argon flowing through it slowly. This will help the distillation along as well as insuring there is no air inside.

[Edited on 20-3-2008 by kilowatt]

crazyboy - 7-4-2008 at 14:46

Sorry guys I’m going to have to side with len1 on this I just retested the experiment with NaOH and aluminum results seemed positive gray solid burns yellow reacts with water almost everything you would expect; almost. I noticed the "sodium" was very brittle previously I had just attributed this to impurities but rethinking it this is what is probably happening:

The sodium hydroxide melts and forms a very intimate mixture with the aluminum. The gray stuff is just sodium hydroxide and aluminum fused together when they are added to water the same reaction occurs as if you were adding aluminum foil to NaOH soln. hydrogen and heat are produced. But because the mix is so intimate much hydrogen and heat is produced and sometimes that heat ignites the hydrogen.

hashashan - 16-11-2009 at 12:34

I probably don't understand something but had anyone tried electrolyzing a NaCl solution in glycerol?
what will happen?

bbartlog - 16-11-2009 at 14:05

Quote: Originally posted by hashashan  
I probably don't understand something but had anyone tried electrolyzing a NaCl solution in glycerol?
what will happen?


I wouldn't know the specifics, but given that sodium glycerate is a known compound I think you can be fairly sure that you aren't going to end up with elemental sodium. If you poke around in the scimadness library a while you may be able to find a reference in one of the older books on electrochemistry; they record a lot of results for simple experiments such as the one you are describing.

Ephoton - 17-2-2010 at 14:22

think I found a new one. it uses amalgumated copper electrodes and a salt solution.

it works for gallium as well :)

http://www.patentstorm.us/patents/4012298/description.html

madscientist - 6-11-2010 at 20:17

ordenblitz, if you're still around, I'd be interested to hear what the results would be with an attempt to prepare NaAlH4 in an manner analogous to your NaBH4 preparation:

6NaOH + 4Al ----> NaAlH4 + 2NaH + 3NaAlO2

Reading this thread I can't help but think the most useful application of this thermite reaction is not to be found in the alkali metals, but rather a handful of powerful reducers and bases that are otherwise unavailable to the amateur - aluminum hydrides in particular.

production of sodium metal

symboom - 21-11-2010 at 19:44

i have made sodium metal from sodium hydroxide with great success using a wide iron container with the cathode patuding from the bottom of the container then putting iron screen would have been more pure if i used argon gas

potassium metal however from potassium hydroxide i am going to have to use argon

i read that molten fused sodium cyanide electrolysis but cyanide compounds are toxic and a problem and correct me if i am wrong but the products formed are cyanogen a toxic and flammable gas and sodium metal

does anyone what is the electrolysis of molten sodium cyanate products

[Edited on 22-11-2010 by symboom]

Potassium made in Germany

chedo - 2-12-2010 at 13:14

Quote: Originally posted by len1  
I dont think your results are different. A white precipitate is generated from the reaction of Mg with H2O, that would occur in pure hot water as well. The reaction generates H2 although at half the rate to what the patent authors 'measured'. Thats the effervesence. But the point of the patent is not to form magnesium hydroxide.

My request to you or anyone else who thinks this reaction works, is

please show the potassium.

Anything else is pure speculation. Len

[Edited on 27-12-2007 by len1]


Look here:
http://www.versuchschemie.de/topic,14677,-Synthese+von+Kaliu...

Quote:
Abstract not available for EP0244894
Abstract of corresponding document: US4725311
A process of producing alkali metal by a reduction of alkali metal compounds with metallic reducing agents, which is simplified and avoids corrosion of the reaction vessel and part thereof in that the reduction is effected with particulate magnesium in an organic solvent which under the conditions of the process is inert to the alkali metal to be produced. The reaction is effected at temperatures from 100 DEG to 300 DEG C. In an embodiment of the process the reduction is effected in the presence of an alcohol used as a reaction accelerator.

runninfarmer - 5-12-2011 at 19:17

I've seen a lot of people make sodium metal from NaOH using Mg as a reducer. I've also read the two large threads about it on here.

But I can't see where anyone used NaCl alone with Mg to produce Na metal. Does it not work? Do you have to use NaOH to get the reduction to work?

Thanks for the help

zuluwhiskey - 27-1-2012 at 06:45

I've been scouring the internet on making sodium and here's what I found:
http://en.m.wikipedia.org/wiki/Sodium_amalgam
If you look at the "uses" section, you'll find:
Quote:
Sodium amalgam is a by-product of chlorine manufactured by mercury cell electrolysis. In this cell, brine (concentrated sodium chloride solution) is electrolysed between a liquid mercury cathode and a titanium or graphite anode. Chlorine is formed at the anode, while sodium formed at the cathode dissolves into the mercury, making sodium amalgam. Normally this sodium amalgam is drawn off and reacted with water in a "decomposer cell" to produce hydrogen gas, concentrated sodium hydroxide solution, and mercury to be recycled through the process.

Just distill of the mercury in an inert atmosphere?...
Sodium amalgam in action

Alastair - 1-3-2012 at 00:08

len1 and everybody else, i understand that skimming sodium off a solvent is perfect and you get it nearly clean.
however, has nobody thought of dipping a strong magnet into any nps and thus recovering sodium from MgO or Al2O3 or whatever acids and salts?

Or am i just clumsy? :D

watson.fawkes - 1-3-2012 at 04:06

Quote: Originally posted by Alastair  
dipping a strong magnet
What substance in the reaction mixture would respond to a magnetic field, and how would it respond?

Alastair - 2-3-2012 at 08:21

Quote: Originally posted by watson.fawkes  
What substance in the reaction mixture would respond to a magnetic field, and how would it respond?


Sodium is paramagnetic, so "is only attracted when in the presence of an externally applied magnetic field" (wiki)

Maybe it would just stick to the magnet leaving behind most of impurities. But that is just a guess. It might not work with simple magnets, and i for one wont test it right now, though i wish i could.

Just it seems to be an idea worth trying, if one doesn't want to distill.

watson.fawkes - 2-3-2012 at 17:16

Quote: Originally posted by Alastair  
Sodium is paramagnetic, so "is only attracted when in the presence of an externally applied magnetic field" (wiki)
The strength of that interaction is minuscule. It also decreases with temperature. Plus you're talking about a liquid, not solid, so there's yet another source of spin disordering and an even lower strength of interaction.

Alastair - 3-3-2012 at 02:45

Quote: Originally posted by watson.fawkes  
The strength of that interaction is minuscule. It also decreases with temperature. Plus you're talking about a liquid, not solid, so there's yet another source of spin disordering and an even lower strength of interaction.


I won't start a pointless discussion about sodium being a solid.

watson.fawkes - 3-3-2012 at 07:50

Quote: Originally posted by Alastair  
I won't start a pointless discussion about sodium being a solid.
That's good, because in all the elemental sodium syntheses I've seen here, the sodium comes out at elevated temperatures, as a liquid.

Alastair - 4-3-2012 at 11:52

Quote: Originally posted by watson.fawkes  
That's good, because in all the elemental sodium syntheses I've seen here, the sodium comes out at elevated temperatures, as a liquid.


That's good, now you can start a chemistry/physics revolution and change all the literature on sodium, and everything else that liquifies just below 100C. Good luck ;)

watson.fawkes - 4-3-2012 at 15:04

Quote: Originally posted by Alastair  
That's good, now you can start a chemistry/physics revolution and change all the literature on sodium, and everything else that liquifies just below 100C.
Are you actually asserting anything? I can't tell. Otherwise I'll have to assume you're trolling.

Alastair - 5-3-2012 at 12:25

Quote: Originally posted by watson.fawkes  
Are you actually asserting anything? I can't tell. Otherwise I'll have to assume you're trolling.


You can assume whatever you want. I feel that this "dialogue" is getting quite boring for others to read. If you think it is a big problem to cool sodium down to room temperature, you can aswell go on theorizing about it. My point was just to propose an idea for someone who would actually try it, wouldnt hurt i think? Anyway, thanks for your oppinion.

watson.fawkes - 5-3-2012 at 12:45

Quote: Originally posted by Alastair  
If you think it is a big problem to cool sodium down to room temperature, you can aswell go on theorizing about it. My point was just to propose an idea for someone who would actually try it, wouldnt hurt i think?
I see. Cool it down and lock in all the inclusions into a solid matrix where you're going to get zero separation. Well, that won't work either.

So "someone who would actually try it", SWWATI, is apparently not you. Well, now I know. At least now I realize I'm not trying to help you, but rather the anonymous future reader who might not realize that they'd be wasting their time.

SWWATI, while not as obnoxious as SWIM, has at least a certain responsibility to make sure your proposal isn't totally hare-brained. If you're going to do armchair chemical engineering, then, at least have the grace to put a sanity check on your idea, like, oh, calculating some numbers.

metalresearcher - 22-12-2012 at 13:08

Last week I got sodium by heating Na2CO3 + C in a steel tube to 1200C but capturing under oil did not work yet.

<iframe sandbox width="640" height="360" src="http://www.youtube.com/embed/9l_DojAugyg" frameborder="0" allowfullscreen></iframe>

And the same for K2CO3 + C:

<iframe sandbox width="640" height="360" src="http://www.youtube.com/embed/nUQNjO32vyQ" frameborder="0" allowfullscreen></iframe>

m1tanker78 - 24-12-2012 at 08:35

Nice demo MR! I've had some success in the past from condensing sodium vapor on a cooled metal surface. The vapor was generated in an overheated Downs cell (the reduced sodium boiled off instantly) and collected inside of a tube. Collecting the sodium in oil is definitely worth a try but you may find quite a bit gets trapped in the retort outlet.

How are you drying the carbonate salts before charging the retort?

Tank

aliced25 - 28-5-2013 at 03:51

Right, I posted in the wrong forum as I was a little stunned by this video (http://www.youtube.com/watch?v=seSg_GWj1b0).

Now, there are some serious problems with what he is doing, but the basic procedure is exciting. If Mg can reduce NaOH, then if we heat the material in a furnace/kiln - a propane furnace will hit the right temperatures, then a finely divided, briqueted form of NaOH (fused) and Mg could be heated under a mild vacuum and distilled (the boiling point of Na is 883C at 1Atm, so it would be considerably less under vacuum, at 20mtorr it comes over between 300-350C, (http://www.osti.gov/bridge/servlets/purl/4478378/4478378.pdf) we won't be going that far, but 500-600C is likely with a normal single stage sliding vane pump at 10pa (or 75mtorr).

It is a variant of the Pidgeon Process, which works with Magnesium (Silicothermic), Calcium (aluminothermic), Lithium (Aluminothermic/Silicothermic or Magnesiothermic) and presumably quite a lot of other chemicals.

Attached is a paper on the Design of a Distillation apparatus for Lithium, the v-shaped apparatus looks like something we'd use - although, if we had a furnace with a gas take-off tube for the gaseous metals, then that could be cooled enough to allow the liquid metal to run down into the condensing area.

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aliced25 - 31-5-2013 at 02:12

This shit, Self-propagating synthesis, is insane. Microwaves with a susceptor (generally SiC powder packed between two crucibles) have been used to ignite what are essentially thermite mixtures, which burn in a vacuum/gasless environment, producing everything from TiC and other refractories to NiAl Alloy (useful). I suspect they could be harnessed for the preparation of alkali metals, as the volatile components of various Pidgeon-type processes - just put a hole in the top of an old MW, insert a large crucible, then pack it with SiC, then another crucible, with a cover that can hold a vacuum (no MW are reaching this, so bolt-down lids would be a good idea).

This is some sick shit, thermite in a MW - perhaps we can make Si without needing Sulfur (YAY). :D

blogfast25 - 3-6-2013 at 08:29

Quote: Originally posted by aliced25  
This is some sick shit, thermite in a MW - perhaps we can make Si without needing Sulfur (YAY). :D


Sure. That thermite is slightly exothermic, so if you heat it enough you'll end up with molten alumina and molten Si.

The Standard Enthalpy of Reaction for SiO2 + 4/3 Al === > Si + 2/3 Al2O3 is about - 212 kJ/mol, that's quite a bit for free. Provide the missing heat (to get to about 2,500 C) with MW and Bob should be your uncle. Just don't expect chip grade! :D

aliced25 - 3-6-2013 at 19:21

Don't want chip grade, what I want is a reductant for various other metals, including lithium, possibly strontium, and especially calcium. I'd also like to make the aluminium silicon alloy the Russian's used to reduce KCl with (http://www.dtic.mil/dtic/tr/fulltext/u2/a359639.pdf). What is your idea, first go with a propane kiln and something like this (http://www.scribd.com/doc/123531537/Magnesium-silicothermic-...)? A simple rotary-vane pump should get enough vacuum, if not we need additional heat. Working with thermite mixtures we run into the issue that even under vacuum they start a self-sustaining combustion reaction (http://en.wikipedia.org/wiki/Self-propagating_high-temperatu...) which will rapidly reach and breach the ability of most metals to contain the reaction with some products.

But done properly, this is possibly the best route for the home scientist to get access to some serious quantities of a variety of metals, no electrolysis, no exact heating - if the SHS kicks in heating isn't the issue anymore anyhow- just high quality distilled metals.

[Edited on 4-6-2013 by aliced25]

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[Edited on 4-6-2013 by aliced25]

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watson.fawkes - 4-6-2013 at 04:48

Quote: Originally posted by aliced25  
Don't want chip grade, what I want is a reductant for various other metals, including lithium, possibly strontium, and especially calcium.
If all you want is a source of silicon for reduction, there's no need to use elemental silicon. If crude silicon works for you, then ferrosilicon should as well.

There are four low-melting eutectic points, that is local minima on the phase diagram. The one with the greatest percentage of silicon is the eutectic of FeSi and Si, at around 60% Si. But the whole range from ~ 50% Si to 60% Si has a melting point at 1207 - 1220 &deg;C, a couple hundred degrees lower than that of Si itself at 1414 &deg;C. This is rather easier to get to, as the T4 ratio is 1.8 (that's ratio of radiative heat loss).

And it's a lot easier to make an arc furnace than screwing around with microwaves. Welding rods, castable refractory, and an ammeter (to know when to push the rods in) is all you really need. At the materials are dirt cheap: sand, carbon, and steel scrap.

blogfast25 - 4-6-2013 at 05:35

Quote: Originally posted by watson.fawkes  
If all you want is a source of silicon for reduction, there's no need to use elemental silicon. If crude silicon works for you, then ferrosilicon should as well.

There are four low-melting eutectic points, that is local minima on the phase diagram. The one with the greatest percentage of silicon is the eutectic of FeSi and Si, at around 60% Si. But the whole range from ~ 50% Si to 60% Si has a melting point at 1207 - 1220 &deg;C, a couple hundred degrees lower than that of Si itself at 1414 &deg;C. This is rather easier to get to, as the T4 ratio is 1.8 (that's ratio of radiative heat loss).

And it's a lot easier to make an arc furnace than screwing around with microwaves. Welding rods, castable refractory, and an ammeter (to know when to push the rods in) is all you really need. At the materials are dirt cheap: sand, carbon, and steel scrap.


Ferrosilicon was mentioned in the other thread.

Ferrosilicon can be made easily by mixed thermite (Fe2O3 + SiO2 in the desired ratio, + Al and CaF2). The Fe2O3 acts as the heat booster to provide the 'missing heat' for the straight SiO2 thermite. I've done this with good results a few years ago. Because various FeSi mixtures are more dense than pure Si, you also get much better gravitational separation which tends to be a problem with sulphur assisted SiO2 thermites. Not to mention the H2S!

FeSi globules from aluminothermy: still quite uncorroded after so many years:



Quote: Originally posted by aliced25  
What is your idea, first go with a propane kiln and something like this (http://www.scribd.com/doc/123531537/Magnesium-silicothermic-...)? A simple rotary-vane pump should get enough vacuum, if not we need additional heat. Working with thermite mixtures we run into the issue that even under vacuum they start a self-sustaining combustion reaction (http://en.wikipedia.org/wiki/Self-propagating_high-temperatu...) which will rapidly reach and breach the ability of most metals to contain the reaction with some products.


If it was me, definitely propane kiln.

Self-sustaining reactions obviously wouldn’t work here in most cases. It’s PRECISELY not self-sustaining reactions that should be the target of this ‘reactive distilling’.



[Edited on 4-6-2013 by blogfast25]

aliced25 - 4-6-2013 at 17:25

Precisely, these reactions aren't self-sustaining, they need vacuum in order to proceed. @Watson.Fawkes, we are trying to avoid the need for an arc, for one thing it is going to seriously complicate building a vacuum container and for other, external heat is applied to cause the reaction to kick in.

For mine, if we can make pure silicon, we can make pure silicon, it is about the same price on pyro sites as Ferrosilicon anyway, the only thing that has put people off it in the past is the slag forming H2S. If external heating can be used to kick the fucker off, then a simple Al/SiO2 reaction with Si and Al2O3 as the slag would be fucking brilliant. The less crud in the pot the better when trying something new out is how I look at it (although Pidgeon, et al, did use Ferrosilicon for the reduction of both Lithium and Magnesium, because it was cheaper than Al - not our problem, for us Al is significantly cheaper, as is Mg for that matter).

An externally heat-ignited thermite mixture, would presumably have very similar effects on crucibles as the normal thermite reactions, so I would suggest that any vacuum vessel be designed to have a single-use crucible, that is cheap and expendable and can be placed in there easily (ie. the thing should be able to be pulled apart, in fact modular would be great). The difficulty is, how to we get the external heat through an expendable crucible to the reagents, sufficient to raise them to initiation temperatures? How do we get a metal that will allow us to push that much heat through it without melting/deforming under vacuum? I know the Russian's used a Tungsten wire to initiate the SHS mixtures, which could be conceivably pushed through the vacuum housing and potentially through a small hole in the base of the expendable crucible. But that gets away from the propane idea into electrodes.

A lot of questions to be answered yet, not to worry, we have the collected literature of NASA, the US Military and the USSR's best Scientists to help us answer it.:D A straight crucible, made of something like alumina (or porcelain - such as a porcelain vase/cup) that costs fuck all, would be the first part. SS to contain it, which would need to be vacuum sealed to the main chamber (containing air/water cooled condenser and collection cup). I'm thinking flanges here, bolted flanges joining the two components with a gasket.

I'm thinking if we keep the original small, then catastrophic failure of any component would be no more dangerous than a normal thermite reaction. A vertical design for the area going into the furnace, simplifying the system and probably an angled condenser coming out of the area above it to both the collection cup and vacuum take-off (like a normal pot-still). I think given the range of materials this will be tested with, from alkali metals to phosphorus, SS is going to be the best option, I'm just worried about it's ability to cop the heat at the furnace end.


[Edited on 5-6-2013 by aliced25]

watson.fawkes - 4-6-2013 at 21:56

Quote: Originally posted by aliced25  
we are trying to avoid the need for an arc, for one thing it is going to seriously complicate building a vacuum container and for other, external heat is applied to cause the reaction to kick in.
I didn't say at all that you'd use the same reactor. Use an arc reactor for lots of ferrosilicon, which you take out and use somewhere else.

Edit: typo. clarity.

[Edited on 2013-6-5 by watson.fawkes]

aliced25 - 5-6-2013 at 04:59

Oh ok, my mistake

watson.fawkes - 5-6-2013 at 07:16

May 1933 article from Modern Mechanics
Experimental Arc Furnace Melts Anything

aliced25 - 5-6-2013 at 20:51

Yep, and there is this one (cited in the Phosphorus reduction papers). The difficulty I foresee is the circuit (I'll need help on that) and how to modify it for vertical operation (otherwise we'd be restricted to what could be put in a boat inside the furnace).

BTW Could you grab that entire paper?

PS Here is another laboratory furnace by Kroll, et al, (pdf) for the production of Zirconium.

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[Edited on 6-6-2013 by aliced25]

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aliced25 - 9-6-2013 at 04:56

Stauffer has probably got the simplest furnace design I have ever seen for this (on page 3/11) - a simple burner, in a kiln/furnace with bricks, and a condenser attached to a vacuum pump.

He uses Spodumene with ferrosilicon and uses calcination, which is going to be a headfuck and a half, given Li makes up only 3% of Spodumene even minor screwups will lose half your yield.

At the same temperatures though, using Lithium carbonate ==> Lithium Oxide (by thermal decarboxylation - mix with with Calcium Carbonate and they both decarboxyalate at a lower temperature I read somewhere), then silicon/ferrosilicon or aluminium to get lithium. Looks awfully simple, a decent vacuum would be needed, as would a propane torch/meker burner.

Like I say, start from the more expensive lithium carbonate which has a way higher percentage of lithium and save the fucking around (http://digitool.library.mcgill.ca/webclient/StreamGate?folde...). Same setup apart from that. As CaO is going to be needed, and there is apparently a eutectic (according to one paper), it might need to be looked into further. In the presence of Al powder, it undergoes decomposition to the oxide without being wetted (too low a temperature) at 700C or so.



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aliced25 - 11-6-2013 at 01:42

Probably the best way to work this is to try out the preparation of Calcium metal from a much cheaper ore, similar reaction, similar product, just a cheaper ore.;)

Attachment: Di.Wang.Tao.Feng.A.Novel.Vacuum.Aluminothermic.Reduction.Lithium.Process.pdf (305kB)
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aliced25 - 13-6-2013 at 00:49

These seem to take me over my file size, so I'll attach just two then the other one

[file]23831[/file]

[file]23827[/file]

halogen - 8-8-2013 at 11:48

http://link.springer.com/article/10.1007%2FBF00608791

I don't believe it! So what do you get on the other side, products from oxidation of propylene carbonate? Perchlorate?

[Edited on 8-8-2013 by halogen]

bfesser - 8-8-2013 at 12:19

<strong>halogen</strong>, <a href="viewthread.php?tid=15455&page=28#pid236897">that paper</a> is already in <strong><a href="forumdisplay.php?fid=21">References</a></strong>. This has also previously been discussed for Na, K, and Li in several threads, if I recall correctly.

[Edited on 25.9.13 by bfesser]

halogen - 8-8-2013 at 15:09

Ah, thanks. So they cheated using sacrificial alkali anodes.

Though, interestingly as the AlX3 adduct seemed to produce an alkali deposit, maybe some kind of aluminum anode or zinc would suit the amateur. Nonetheless I'm curious what might happen were there an inert anode. For what it's worth electrolysis of organics in HF is a known flurination, it might be interesting to examine the results of electrolysis of fluoride salts in non-aq. solution.

testimento - 25-9-2013 at 13:12

I am planning to reutilize my micro-castner soon and post some pictures because I'm gonna retry the chlor-alkali - combination process with molten NaCl:CaCl 70:30 at 650C. The reactor design is following (and later you see picture of the reactor I have constructed earlier at springtime) - ATTACHMENT. The reactor essentially consists of anode from underside and cathode around the reactor from upside, and I will contain it within clay an layers of rockwool to make it a lot more energy efficient. I'ts made out of 316 steel and I'll be using graphite anode and iron cathode with 5V at about 50-100A depending on the conductor's cross diameter and EC values. This is rather low power, but since it could procude sodium and clean chlorine gas for chlorination processes, I'll give it a go-code. The chlor-alkali will be waiting for polymer membrane or 2 kilograms of mercury until procceeding.

The reactor volume is about 4 liters so it should ideally hold about 6-7kg of chloride salts, which of about 5kg is sodium based and will be electrolyzed, and of this, in theory, about 2000 grams of sodium and 3000 grams of chlorine gas can be obtained. Heating should be carried out using 6kW propane torch.

downs-cello.png - 34kB

testimento - 28-10-2013 at 07:04

I made test run with my electric heater. This is molten sodium chloride from my reactor, 500 grams, the color is from temperature.

salts.jpg - 130kB

WGTR - 28-10-2013 at 11:53

testimento, maybe I'm misunderstanding the diagram.
Are you collecting sodium metal at the top of the cell, with the
molten metal in contact with the cathode? If so, this may give
you some trouble.

The old literature generally shows the cathode on the bottom of
the cell, with the sodium metal collected as it floats to the top.

I used a fireclay ceramic as a divider once, thinking that it
would allow me to collect sodium on one side, and chlorine on
the other; with the electrodes coming in from the top. After a
short period of time the ceramic became conductive (@#&^!).
Apparently the salt electrolyzed through the ceramic,
contaminating the melt in the process.

It takes a high-quality ceramic to work under those types of
conditions.

testimento - 28-10-2013 at 12:44

I was thinking this type:

http://images.flatworldknowledge.com/averillfwk/averillfwk-f...

The anode shall be put through the bottom with plaster fitting that insulates it from the bottom, and the cathode through the top flange that is lifted off when the system is opened.

How the sodium shall interfere with the cathode? What it does? Should I consider this structure straight away?

https://s3.amazonaws.com/readers/2010/02/17/downscellschemat...



[Edited on 28-10-2013 by testimento]

cell-downzs.png - 21kB

WGTR - 28-10-2013 at 13:29

In normal cell operation, the diaphragm (and collected sodium)
potential is partway between the anode and cathode potentials.
When the cell is operated at the correct current density (i.e.,
not too high of a cell potential), the metal diaphragm remains
electrically inert. If the cathode comes in through the top of
the melt, then the sodium can collect around it, shorting the
cathode to the diaphragm. At this point the diaphragm
becomes the cathode, and sodium is produced in the same
place as the chlorine.

If a ceramic insulator is used under the same conditions, with
the intention of isolating the sodium metal from the diaphragm,
care has to be taken in the selection of its material. The lousy
fireclay separator that I used was apparently not dense/inert
enough, and the sodium electrolyzed right through it, essentially
making it part of the cathode. Sodium beta-alumina is
discussed as a diaphragm material in at least one thread here
(maybe this one), and may work under these conditions.

The image that you linked to actually shows the cathode and
sodium metal electrically isolated from each other. That might
not be obvious right away.

If you want to run the cathode in through the top, you can see if
it's possible to electrolyze the sodium metal into a solid alloy at
the cell temperatures. That way the sodium metal is
immobilized. Na-Pb or Na-Sn systems might be a place to
start. If you make the container itself the cathode, you may be
able to electrolyze the sodium into a porous mass of iron
powder. If that worked current density would be critical, in
order to keep the sodium from just forming on the surface
closest to the anode. It'll probably be easier just to do it the
conventional way, though (with the cathode entering through
the bottom, or the side).

testimento - 28-10-2013 at 14:01

I have understood that no diaphragm is needed for this operation, as long as the electrodes are accordingly separated. I've redesigned the system to prevent the cathode from touching the sodium formed, and the metal is expected to accumulate on the top, and flow out of the reactor when the container is full and more sodium chloride is to be added.

Btw, as a quick off-topic, are there effective ways to make chlorine gas into hydrochloric acid?

WGTR - 28-10-2013 at 14:31

If you're not worried too much about efficiency, the diaphragm
is not needed. What it allows mostly:

1. The cathode and anode can be very close together without
the products mixing. The geometry of the cells is designed
mostly to allow maximum current and efficiency, and minimal
use of supplemental heating or cooling. Generally this means
that the electrodes have to be pretty close together, and the
current density has to be consistent across the electrodes. For
your purposes you can run a lower current density and use
supplemental heating, so the electrode spacing is not as
important.

2. It helps circulation of the electrolyte. The diaphragm is
porous enough to allow the molten salt to circulate through it,
without allowing the chlorine or sodium to mix. Again, if your
current density is low, and your heat comes from supplemental
sources, this is not so important. If you're using solid metal
containers around your cathode to catch the sodium as it floats
off, you have to be careful not to run the current up so high that
the resulting turbulence carries your sodium right back out and
into the chlorine compartment (or vice versa).

You can try this with regular water electrolysis. Even with both
electrodes in inverted test tubes, at high currents the turbulence
will carry the gas bubbles right out the bottom.

I think that other people can answer the Chlorine to HCl
conversion question better than I can.

[Edited on 28-10-2013 by WGTR]

testimento - 28-10-2013 at 15:38

I have my old power source rated at 400A at 6VDC, but with rewinding it can go as high as 600 amps. I was preferring a graphite electrode with CSA 500mm2 and a similar measurements for steel cathode. It's supposed to be heated up with electric heater, and I'm quite sure that the heat from the electrolysis is not sufficient to keep it over 600C, but the temperature shouldn't be a problem.

How many millimeters do you think the distance between the electrodes should be? In my concept the distance could be limited to ~100mm minimum, but changing the design to square instead of spherical, it could be lowered to maybe 10mm minimum. The schematics show that commercial cells use only steel mesh to prevent the chlorine and sodium mixing. I'd rather drive the cell with more power(to a sensible extent, of course) than make the structure more complicated, because in amateur setups, the simplicity usually inhibits costs and provides reliable operation. Membranes are, to my experience, highly prone to errors.

m1tanker78 - 28-10-2013 at 18:25

Quote: Originally posted by testimento  
I have my old power source rated at 400A at 6VDC, but with rewinding it can go as high as 600 amps[...]


You won't be able to push 400A (let alone 600A) through molten NaCl @ 6V with any reasonably-sized electrodes. If your supplemental heat can keep the salt molten by itself then you can use just about any DC power supply. A Downs cell (particularly straight NaCl - no fancy salt mixtures) is damned finicky. Keep in mind that your electrodes and any other 'features' will sink heat from the melt.

Tank


testimento - 28-10-2013 at 19:21

I have 20 units of 10x60mm graphite electrodes, 200mm length, and I checked their conductivity is high enough for the lower part even when over 600C. Of course, the resistance at very high temp conditions in molten salt and other stuff is not anything close from theory, but it's enough for me if I can push 50-100 amps off that record through the pot.

A mixture of nacl:cacl 42:58 will melt at 600C. I tested this yesterday, and it's the common mixture used in downs cells. Another mixture consisting of magnesium, barium and sodium chloride will melt as low as 450C temp, but I haven't got those at hand now.

WGTR - 29-10-2013 at 08:57

testimento, I can't say what your cell dimensions should
be without it being speculation, but 10mm spacing between
anode and cathode sounds too close for your type of setup.

For your NaCl/CaCl2 melt, I'd suggest solidifying some of the
molten salt on a steel rod, and trying to dissolve it in water. All
of the salt should dissolve. If some is left over, then perhaps
there is CaCO3/Ca(OH)2 contamination, and the melt needs to
be purified. The pH of the solution shouldn't be basic.

Zyklon-A - 28-1-2014 at 14:40

After reading most of this topic, it's clear that Mg or Al can take oxygen from (most) group 1 metals, does this reaction seem likely to happen: 2KNO3 + 4Al--> K + 2Al2O3 + N2?
Contrary to what some people have told me, it's easy for me to ignite a nitrate-Al mixture without sulfur.

Edit: Without waiting for a response, I went out and tried it.
I mixed KNO3 and Al in stoichiometric proportions, (as shown above). In a steel can with a brick on top, I proceeded to light it up. Soon the can was glowing red hot, and the wood which it was on, caught on fire.
After it cooled, I saw some little balls of metal on the bottom of the can (actually almost all of the bottom of the can had melted away), I dropped a drop of water on them, but nothing happened, I assume that they are steel nuggets. After dropping some more water on some more of the slag, I noticed quite a bit of fizzing! It must have been potassium, the can was way to cool to boil water, and there couldn't have been any other 'active' metals/materials in there AFAIK.
So the answer is yes, you can make potassium from KNO3 and Al, but it's very ineffective, and the product is quite impure.
I know that this topic is about sodium, and I'm sure if you switched the KNO3 with NaNO3, the result would be the same, except you would get impure sodium, instead of impure potassium.

[Edited on 28-1-2014 by Zyklonb]

quantumcorespacealchemyst - 24-10-2014 at 03:22

this seems the best place to inform that a 27ish mm quartz tube, domed on one end, with a magnetic stirrer and a small 12v UV lamp set up run on two 6v lantern batteries in series, enclosed with aluminum foil reflectance, all run on a mixture of atomized copper powder and ground iodine crystals in distilled water afforded a pure white powder, seemingly copper i iodide, CuI, which precipitated out of solution rather efficiently. settling in layers if reccolection is correct.

prior approach using electrolysis was unsuccessful with same 2x6v series lantern batteries.

[Edited on 24-10-2014 by quantumcorespacealchemyst]

Metacelsus - 24-10-2014 at 04:03

Right forum, but wrong thread.

m1tanker78 - 2-11-2014 at 07:11

Quote: Originally posted by testimento  
I made test run with my electric heater. This is molten sodium chloride from my reactor, 500 grams, the color is from temperature.


Judging by your pic, it's ripe for electrolyzing -- maybe just a touch too hot. You're looking for a black hot, so to speak, with a touch of red glow as seen in a dimly lit environment. If your PS puts out enough juice, you can directly brute force it by electrolysis to preheat everything then back off to that sweet spot. Introducing additional hardware for the apparatus may tax your electric furnace beyond its capability or fail to properly fuse the salt.

Tank

experimenter_ - 29-6-2016 at 09:12

According to this source: http://chemiday.com/en/reaction/3-1-0-8667
sodium oxide decomposes to sodium and sodium peroxide at high heat. Also, sodium oxide can be produced by thermal decomposition of sodium carbonate at 1000 C.

Do you think that heating sodium carbonate e.g. in a tube furnace might produce some metalic Na?
Special atmosphere might be required (vacuum). Also Na evaporates at 883 C ...

Another source refering this reaction:
http://www.allreactions.com/index.php/group-1a/natrium/sodiu...


clearly_not_atara - 4-4-2017 at 14:27

Most non-coordinating anions are hard to synthesize, but bisoxalatoborate salts can be made by simply heating oxalates with boric acid and oxalic acid, or by heating carbonates with both acids. This patent focuses on the purification but describes the prep:

https://www.google.com/patents/US20100145076

I've been thinking of what to do about this -- lithium bisoxalatoborate is soluble in most organic solvents and has received attention as a battery electrolyte, but sodium bisoxalatoborate is relatively unstudied. However, I believe the sodium bisoxalatoborate is soluble in solvents such as toluene or pyridine as well. This paper considered NaBOB for a battery electrolyte before settling on the related NaDFOB (diflurooxalatoborate), and notes that the solvent compatibilities of the oxalatoborates are better than Na perchlorate:

http://pubs.rsc.org/en/content/articlelanding/2015/cc/c5cc02...

Anyway, if you dissolve NaBOB in some sort of solvent, you can probably electrolyze this to get sodium. I'm not sure what the byproduct will be, though. One possibility:

(C2O4)B- >> e- + B(s) + 4CO2(g)

If boron is produced at the same time as Na this would be quite a useful coincidence but I suspect the actual product may be boric acid or sodium metaborate.

[Edited on 4-4-2017 by clearly_not_atara]

mayko - 6-4-2017 at 17:42

NurdRage had this to say this morning:

<blockquote class="twitter-tweet" data-lang="en"><p lang="en" dir="ltr">HOLY @#%!$ I ACCIDENTALLY MADE BULK SODIUM METAL WITHOUT ELECTROLYSIS. This is a MAJOR breakthrough for amateur chemistry.</p>&mdash; NurdRage (@NurdRage) <a href="https://twitter.com/NurdRage/status/849812574854160384">April 6, 2017</a></blockquote>
<script async src="//platform.twitter.com/widgets.js" charset="utf-8"></script>

the tone at least got walked back:

<blockquote class="twitter-tweet" data-lang="en"><p lang="en" dir="ltr">hmmm... looking closely now this might not be worthy of a peer-reviewed paper... but it&#39;s still a breakthrough for amateur chemistry.</p>&mdash; NurdRage (@NurdRage) <a href="https://twitter.com/NurdRage/status/849817272483028996">April 6, 2017</a></blockquote>
<script async src="//platform.twitter.com/widgets.js" charset="utf-8"></script>

<blockquote class="twitter-tweet" data-lang="en"><p lang="en" dir="ltr">okay, turns out this sodium method is not novel enough for a peer-reviewed paper. But it&#39;s useful for the amateur so i&#39;ll make a video on it</p>&mdash; NurdRage (@NurdRage) <a href="https://twitter.com/NurdRage/status/850126265298432000">April 6, 2017</a></blockquote>
<script async src="//platform.twitter.com/widgets.js" charset="utf-8"></script>

Still, I wonder what this could be? Might be worth keeping an eye on!

PirateDocBrown - 8-4-2017 at 08:34

If I had to guess, using Mg to reduce NaOH, and then some means of purifying that.

JJay - 8-4-2017 at 08:48

Accidentally making bulk sodium... I am burning with curiosity about what he was trying to do....

j_sum1 - 15-4-2017 at 21:04

Quote: Originally posted by PirateDocBrown  
If I had to guess, using Mg to reduce NaOH, and then some means of purifying that.
Nurdrage's latest video confirms you are correct. But it looks like we will need to wait a bit to learn what the purification process is.

If this is as successful as it seems it could become another hime chemist standard similar to chloroform, hydrazine sulfate, distilling nitric acid, and chlorine from TCCA.

symboom - 17-4-2017 at 21:54

Anyone try molten sodium magnesium oxide aggragate separation using candle wax

MgO density 3.58 g/cm³
Sodium density 0.968 g/cm³

ecos - 25-4-2017 at 01:06

is there any better way for collecting sodium in electrolysis cell than a syringe or a spoon ?
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