Sciencemadness Discussion Board


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Xenos - 13-1-2003 at 19:17

Ok, well my power supply came from ebay (5V @ 50amps) :D so i tried some simple electrolysis. Using carbon electrodes, as close together as possible i first tested it in a salt water solution. It bubbled rapidly, and i plan to attempt to make and capture the Hydrogen and fill a balloon with it. I then proceeded to try KOH. I heated it on my hotplate and put the electrodes in. Soon it began bubbling rapidly, and forming a kinda grayish mix. However, i stoped that experiment. Yesterday i tried to heat some NaOH, but my hotplate wasnt getting hot enough, so i put a torch to the crucible and it began to bubble away. This happened for about 30 min, while the NaOH was turning a silvery black color. ( I suspect this was the C coming off, and they were a bit smaller than before). However, my crucible got a hole in it, so i stoped. That sat overnight, and i then wetted the mix, to find some whitish specks on the side got hot when water touched it. I think if any Na or K formed, the water from the + electrode converted it back to OH as soon as it formed. Today i made a crucible with fire cement and put an C rod in the center. Hopefully that will prevent anything from getting to the material.

madscientist - 13-1-2003 at 20:32

I prepared some sodium metal this last Saturday myself. I placed 20g of NaOH in a ceramic crucible, and proceeded to heat it on a hotplate until it melted. I inserted two electrodes, both graphite, and bubbling began at the anode. I left it unattended for around an hour. When I returned, the crucible contained a greyish, silvery liquid, permeated with greyish-black lumps. I allowed it to cool, and poured in charcoal lighter fluid to protect the sodium from oxidation. I found that the somewhat grey and faintly silvery solid was brittle when cold. It clearly contained a lot of graphite impurity. I chipped some of it out, and saved it in a beaker containing charcoal lighter fluid. I then reheated the contents of the crucible to about 100C; and poured as much of the liquid sodium as possible into that same beaker. Localized boiling occured for a few moments; then several dull silvery pieces of sodium metal settled at the bottom. I allowed the crucible to cool, and rinsed it with cold water. There was several sizzling sounds, and what was left of the heavily-corroded cathode disintegrated completely. The mass of the ceramic crucible fell from 73.5g to 72g over the course of the experiment. The anode was not corroded to any noticeable degree. Next time, I will use a copper cathode, and will pour off the sodium metal while it is still molten (immediately after electrolysis) into chilled xylene.

speculating from my armchair

Polverone - 13-1-2003 at 23:37

In an ideal world, DuPont would deliver ingots of alkali metal to our doorsteps for production costs + 15%, or we would all have sophisticated machine shops and expendable hunchbacks to operate the apparatus that we built. In the real world, I have never been able to confirm anybody's home production of significant quantities of such metals. I recall reading somebody's old Usenet posts that claimed the writer had build some fairly good apparatus and didn't have nearby neighbors, and so was able to produce kilogram-range quantities of potassium and sodium from the chlorides. Wouldn't I like to do that.

With alkali electrolysis the main problem (apart from atmospheric isolation) is (IIRC) be protecting the metal from the products produced at the anode, which dissolve in the molten alkali and can re-oxidize the metal. Unfortunately, this process is so old that references to its specifics are difficult to find.

I would like to try using an equimolar mixture of NaOH and KOH, which is a eutectic melting around 200 degrees C (lower MP should significantly increase ceramic crucible lifetime). The resulting Na/K alloy may have an unmanageably high reactivity and low melting point, though.

vulture - 14-1-2003 at 09:26

I think if you were to electrolise that when molten, firstly all the sodium would be produced because of it's higher electronegativity and lower reactivity.
This will ofcourse have its effect on melting point, because the mixture will shift to Na + KOH.

Na/K alloys are often liquids at roomtemperature and very reactive, that is more reactive than pure sodium or potassium metal.

trinitrotoluene - 14-1-2003 at 10:26

I think what can be done to protect the K , Na alloy is do to pour a layer of mineral oil on top of the mixture and then do electrolysis.The only problem is what kind of mineral oil has a boiling point of above around 200*C? I also consiture to do the same with Na metal.If the mineral oil can remail a liuqid and not boil. If theres a hydrocabon that remains a liquid at the melting point of KOH then I guess it can be used to protect the reactive metal.

Marvin - 21-1-2003 at 18:30

Sodium Ive done in small amounts without a blanket, and the air oxidation was amazing to watch, the metal constantly had a yellow aura around it. Small hydrogen explosions tended to propell sodium across the room and this is something of a hazard. I used a copper cathode and a steel anode/container, this seemed to work quite well.

Since I tried this, about 10 years ago, Ive picked up a fair amount more information from a number of sources.
The most common mistake is to overheat the hydroxide, in which the sodium is soluable. Too hot and the hydroxide/sodium produces a metaloid which conducts electricity and electrolysis stops. Numbers from a book on sodium were, 5 degrees above the melting point of sodium hydroxide you get an 80% yeild on a current basis, and 25 degrees above it you get nothing.

This partially explains why a lot of early references dont use gas burners to make sodium, they rely on the current itself to do the heating, and controlling the current gives much more precision in controling the temperature of the cell.

vulture has a very good point about NaK alloys. This are much more reactive than either of them alone and burst into flames on contact with air at room temp. The general concensus of the people Ive talked to, is that the difference in electrode potential is too small to produce only sodium from a molten mixture of both salts. This is potentially very hazardous.

The book recommends I think 12% of carbonate in the mixture for optimum yeilds, and suggests a 50:50 mixture of sodium hydroxide with sodium sulphide to reduce the melting point. This is aparently patented. Ive found information on hydroxide/nitrate/nitrite eutectics, but its impossible to guess how reduction/oxidation at electrodes would affect the bulk mixture. Something that also bothers me about sulphide mixtures.

a_bab - 26-1-2003 at 04:25

I saw in some chemistry manual the construction of cells for Na production from NaOH, and the Na is produced under some sort of bell-shaped thing. Actually the (-) electrode it's inside of that bell. In this way, because sodium it's lighter than NaOH it'll float over the molten NaOH, under the bell, and it'll be protected in the same time from the air. It's verry simple.

Na with Hg catode

menchaca - 13-3-2003 at 08:52

first of all sorry for my english

Ill try to explain as well as i can

i ve never try this but i think it can work

the idea consist on a quicksilver catode
where the Na forms an amalgam. After the quicksilver takes all th Na that it can (i dont know how long must it take) just take the amalgam and heat it until Hg boils when Hg desapear solid Na should be there you can use an alembic to recover the Hg i´ll try to send you some draws of the idea i had, you can send it to others and they will be able to tell you if it works or it doesnt and improve it

hot electrochemical sodium

It burns - 14-3-2003 at 14:15

Originally posted by trinitrotoluene
I think what can be done to protect the K , Na alloy is do to pour a layer of mineral oil on top of the mixture and then do electrolysis.The only problem is what kind of mineral oil has a boiling point of above around 200*C? I also consiture to do the same with Na metal.If the mineral oil can remail a liuqid and not boil. If theres a hydrocabon that remains a liquid at the melting point of KOH then I guess it can be used to protect the reactive metal.

Im new and possibly wrong but couldnt you use tar, its got a realatively high melting temperature and is a hydrocarbon (at least im pretty sure it is) might be a pain to get the everything back out of though

vulture - 15-3-2003 at 13:11

Mencacha, the quicksilver amalgam becomes solid when the Na content goes above 0,7% and the maximum obtainable is 5% IIRC. Add to that that you're going to need large amounts of Hg which is very expensive.

rikkitikkitavi - 16-3-2003 at 00:18

When you increase the percentage of Na in Hg due to electrolysis , the emk increases and about 0,6-0,7 % Na, it will be the same as for the
H2O=> H2 + OH- reaction , thus it will not absorb any more Na by electrolysis but only decompose water.

So from 1 kg Hg(which is very expensive and hard to get ,lucky me that I have it:) you will recieve about 5 grams of Na.

The to evaporate this amounts of Hg I can imagine the difficulties setting up an adequate distillation arrangement.


Polverone - 16-3-2003 at 12:05

I find it interesting that you refer to mercury as being very expensive. I can order it for about $30/kg, which makes it moderately priced to my way of thinking. That's still an outrageous price compared to what it would cost in bulk, but not so bad compared to many other laboratory chemicals. It seems strange that chemical availability and pricing varies so much by region, when you would expect these commodities to have near-identical prices without geographic differentiation.

a_bab - 17-3-2003 at 01:28

Hmmm...30 $/Kg is very cheap. Is it a regulated chemical ? In my country is banned and you can go into prison if you are caught with it. The price on the black market it's said to be around 1000 $/kg.

Polverone - 17-3-2003 at 10:40

No, the personal possession of mercury is not regulated here, though I am sure it is illegal and punishable to dispose of it improperly. That does explain the price difference, though. Commodities can become very expensive when they are not legally available.


Theoretic - 18-6-2003 at 07:23

Everyone, do yourself a favour and make do without NaOH. How? Well, just take salt with a bit of soda to make a low-melting eutectic (600 Celcius). No corrosion - no water evolved - no trouble getting NaOH.

vulture - 18-6-2003 at 08:15

NaOH has a much lower melting point, secondly, electrode corrosion is very high at 600C and you're dealing with Cl2 vapors.

Besides that, you always give me the impression of being a smartass when posting. Maybe you should calm down a bit and read the threads first before you post something.

jimwig - 18-6-2003 at 15:49

industrially mercury has historically been used to produce chlorine from NaCl acting as the anode (?) for a Whorill cell.

I think it goes like this

Na Cl electrolitically decomposes to Cl as a gas and Na as an amalgum with the Hg acting as an electrode on the bottom of the cell.

Castner Tiegel

Organikum - 18-6-2003 at 20:33

NaOH sodium electrolysis is done with this:

The trick is the iron net between the electrodes (cathode - copper, anode - nickel) which are only 2cm apart. This is a very tight net (100/per cm*cm) and yes it divides also the voltage of about 4V.
If interest I can post more data on this.


Theoretic - 24-6-2003 at 00:30

I think a graphite electrode would be resistant to chlorine, as it's hard to get chlorine to attack carbon.

Organikum - 7-7-2003 at 17:52

I would say the mercury amalgane process is indiscutable for home use - the separation of the amalgane is plain to dangerous. (I don´t say this often, yes?)

The NaCl electrolysis has in my eyes no advantages over the NaOH electrolysis. The "Castner Tiegel" shows the principle for how to setup a electrolytic cell for the NaOH process. Not so diffficult. Materials are common and cheap.

Now I answer all the questions nobody has asked:
Pot: Iron
Cathode: Copper
Anode: Nickel
The distance between anode and cathode is only 2cm. The amperage at the cathode is about 1,6Amp/cmxcm at the anode about 1,1Amp/cmxcm. (Smaller units call for higher current density). The voltage should be about 4V to 5V then, the current density is the more important point.

The magic part:
Between anode and cathode is a iron fabric as diaphragm which has about 100 mesh/cmxcm. This works naturally as a voltage divider. The fabric is fixed in a way that 40% of the voltage are on anode - diaphragm and 60% are on cathode diaphragm.
The maximum temperature is 330°C, 20°C higher as the melting point of NaOH.

Suggestions: The use of steelwool between glassfibre fabric or glassfibre fabric with a steelnet embedded would do for a diaphragm. The better this is the higher will the purity of the Na. I believe the voltage divider function being essential and that the ratio may shift to 30/70 instead 40/60 but not 50/50. The distance of diaphragm - anode/cathode IS critical.

The sodium from NaOH electrolysis is also cleaner as the one from NaCl - simple sedimentation is all whats needed. For even higher quality filtration through an ironfabric or glassfibrefabric under inertgas or petroleum is used. (petroleum is not preferred - fire hazard). Remelting under paraffine is another possibility.

The diaphragm as voltage divider and the bell on top - thats it. This sounds doable in a safe enough way to me and some steelwool for a improvised diaphragm is sure for everybody available without loosing the improvised touch.... ;)


Organikum - 8-7-2003 at 19:29

The thread is on sodium, so a photoreactor is perhaps not the right picture.

I apologize

Downs NaCl Sodium Production Cell

this is on topic I guess and answers what theoretic suggested: Yes! Graphite electrodes are useable in this process! The principle is the same working with an iron diaphragm. I would suggest to exchange anode and cathode for to be able to harvest the Na central what seems much better for a smaller scale unit. But I would fill it with NaOH anyways - no chlorine and much lower temperatures - no question.

blindreeper - 9-7-2003 at 03:24

Does anyone who is good with pottery wanna try to make one of these out of ceramic? I imagine it would not be too difficult. If i had a kiln and some clay and a wheel thingy I'd try it.

rikkitikkitavi - 9-7-2003 at 14:30

pottery usually contains SiO2 wich would be attacked by the molten NaOH. Except from precious metals like Pt, I believe that Ni is fairly resistant to molten NaOH.

Ordinary stainless steel is corroded by lye above 60 C @ 0,1 mm/year, @ 100 C corrosion is about 3-4 mm/year so I expect molten NaOH @ 400 C would show significant corrosion on stainless. But if it was hi-Ni content perhaps?

I can se a problem with the cathode , where Na is formed. the anode reaction would probably also be under quite reductive conditions since H2 is evolved ?

OH- => H2O + H2 + e-


kryss - 15-7-2003 at 13:34

anyone ever try using a carbon fibre elctrode - should nbe shapable.


Organikum - 15-7-2003 at 22:41

cathode = copper
anode = nickel
pot = iron
diaphragm = iron net, 100mesh/cmxcm

for NaOH electrolysis.

Thats the material used in industry and as it is cheap and easy I strongly suggest to use these materials also - whats wrong with them? Every scrapyard provides a pot of thick iron for near free, the nickel willl be the hardest to get as I believe but far from impossible.

The "Castner Tiegel" is more on the point as the more beautiful second graphic!

Also thick iron will stand the molten NaOH for a long time there is a trick: heating from the center has the effect that NaOH solidifies on the bottom and the walls of the cell. This protects from corrosion and contamination.

So preheating by a propane burner to about 290°C and then over the melting point and holding the actual temperature of maximal 330°C (400°C is far out) by resistance/thermoelement inside the cell.

Not bad, isn´t it?

The hydrogen formed together with the Na provides the inert atmosphere over the Na. No problem - a feature.

If you want problems here they are:

- A pain in some bodyparts is the exact fixation of the diaphragm - I suggest strongly to use wider space between anode/cathode (more as only 2cm) and compensate by higher voltage. This provides a good part of the needed heating this way - two with one.

- OTC NaOH is dirt like shit often. Big bad surprise! Use labgrade or test before use a small sample - cleaning the cell is - yes what? A work for vulture? :D

blip - 16-7-2003 at 14:42

Are you sure hydrogen would provide a safe inert atmosphere over the sodium? I've read that lithium reacts with hydrogen without any coaxing and I'd think the same would happen with sodium.

Organikum - 16-7-2003 at 20:22

After five or more books on chemical engineering describe the the Na from NaOH process this way I see a small chance that it can be done this way.

Actually I see this the best way to do this if one wants to use electrochemistry.

Sorry - but I am a little pissed - what more than the original picture from a serious book and the correct translation does it take?
Look at the "Castner Tiegel"!
You will see in the middle above the cathode "Na". And above this "H2".
If the sodium would react with the hydrogen the whole fucking Na from NaOH would be complete impossible!

I am pissed.

rikkitikkitavi - 18-7-2003 at 14:21

organikum, what books did you find this information in. I have been looking for info on electrolysis of NaOH,since it is much lower tempereatures involved.

Can you scan the pages and put on the ftp (when it is up and running) ?


blip - 18-7-2003 at 20:05

Sorry about that Organikum. :( When I see that a thread has new posts, I usually don't reread the entire thread.

Theoretic - 24-7-2003 at 09:33

So, you'd like a low melting point, huh...
Well, here's a candidate. NaHSO4, melting point 177C. No ceramic vessel corrosion -although I'm worried about metal, since the salt is acidic.
Besides, I think the cathode product of the electrolysis would be H2S2O8 - perdisulfuric acid. Hydrolysis would yield H2O2, which is more volatile than the H2SO4 also produced, so it could be distilled off.
Or, alternatively, H2S2O8 could be employed as a powerful (2.07V potential - acid) but not-so-fast oxidizing agent.:cool:

[Edited on 24-7-2003 by Theoretic]


MaxPlanck - 29-7-2003 at 04:55

I think this way you will simply get
lots of H2 and maybe S03 on the other
electrode. As you already said,
this salt is acidic, so it will attack
the Na and H2 will evolve.

Theoretic - 31-7-2003 at 06:33

Oh well, :(.
But what about tribasic sodium phosphate? It's m.p. is 73-77C, and it's not corrosive. The cathode product P2O5 would dissolve in the melt to form polyphosphates, and from time to time one would add CALCINED soda, CO2 is given off - the phosphate is regenerated
(it's a sort of catalyst for electrolysis of, on end, soda).
Just then I thought about NaNO3 and NaNO2 - no, sodium would be reoxidized as soon as formed (but maybe not that fast - worth a try - :o). The same problem with NaOH. :mad:

no such problem with NaOH

Polverone - 31-7-2003 at 12:32

I've heard that sodium may dissolve in NaOH, but I don't think it will reduce it. What sort of reaction were you picturing?

Na reduces NaOH

MaxPlanck - 1-8-2003 at 02:31

at about 25 deg.Cels over NaOH melting point Na2O is formed and H2 is produced.

Marvin - 2-8-2003 at 01:48

I dont think so, though information on the subject is somewhat clouded.

According to a book on the manufacture of sodium metal that Ive been quoting for several years now (paraphased),
Yeild of a castner cell drops from 80% at 5C above to 0 at 25C above the melting point because the solubility of sodium metal in the hydroxide becomes so high the solution becomes a metaloid. This conducts electricity like a metal and electrolysis stops.

I did some searching today however and,
According to US patent number 2,202,270 however sodium metal and hydroxide start to react when there is hydride or hydrogen around at the melting point of the hydroxide. It goes on to say this reaction is preperativly useful for the monoxide at about 450C provided the hydrogen be removed from the cell to alow it to go to completion.

I'm still leaning strongly towards the metaloid explanation though I'm less sure than I was a week ago. Its also the more recent of the sources so the writers of the book should have considered the monoxide reaction allready. Where does your information come from Max?

[Edited on 2-8-2003 by Marvin]


Hermes_Trismegistus - 28-11-2003 at 14:02


listen guys, I had a thought. I am certain that our esteemed progenitors have given a great deal of "energia cogentis"
to the vexing issue of sodium production, while the organic chemists creed is to avoid ultra-high temperatures at all costs, perhaps we are unneccesarily complicating the obvious here.

The DOWNS cell is an apparatus equaled in its sophistication only by its simplicity.

Using the NaCl version seems a little foolhardy and expensive until you factor in the ruggedness and ease of manufacure. And the common availability of materials procurement.

If a clever guy was to enroll in a local community college metalworking course at night school, I am sure he could knock a fairly sizable one off for only a couple hundred dollars that would serve him well for the rest of his life.

and the skills and techniques learned in the course would be invaluable in future endeavors

As for a heat source, Dave Gingery's "lil Bertha" sixty dollar electric crucible furnace that will easily top 2200 degC will certainly be adequate, and last practically forever at the lower heat settings required for common salt electrolysis.

It is also about size of a five gallon bucket and has all sorts of uses.

Perhaps Cl gas is fairly reactive, but not unnaturaly dangerous as long as you were careful to use good, robust gas-tight fittings.

And I KNOW that your collective genius can come up with oodles of uses for the Cl that is evolved.....HMMMM?

Has anyone here looked at a site selling Lindsay's Books? Certainly worth a look:D

Marvin - 28-11-2003 at 22:03

Hermes, were you born with three necks or did that happen later?

I can see the aparent advantages to using NaCl, but NaOH is pretty cheap, widely available and melts much lower. It might cost a little more than salt, but youd probably save that on power bills and the apaeratus would be much safer.

Cl might seem like an added bonus but having a cell that forces you to produce both at the same time is a practical disadvantage at home. Furthur although Cl2 is fairly docile at RT, its almost certainly a different animal at 900C. Add to this that Na has a third of an atmosphere vapour pressure at the melting point of sodium chloride and this is enough to worry me a lot.

I would think most metalworking classes would use brazes that melted below 900C.

I am still convinced that a slightly butchered all stainless steel pressure cooker with NaOH offers the best bet for home sodium production. Keeping the current through the cell adjusted to maintain the caustic soda just above its melting point through nickel electrodes and with a nickel or iron diaphram.

The system has a big advantage that insulation/heating is not critical so you can daisy chain cells together to obtain a working voltage better matched to computer PSUs which will probably be providing the power for most peoples attempts. This would be very difficult with a furnace setup.

[Edited on 29-11-2003 by Marvin]

Hermes_Trismegistus - 29-11-2003 at 00:08

I bow to your obvious munificence

However, it mght be worthwhile to note that NaCl is liquid at only 804 deg C and that with the addition of CaCl (the common industrial process) the temp drops to around 600 degC

as for using brass brazing rods to cobble together your apparatus...:o


I dunno about that Paco....:o

brazed seals are notoriously liable to suprise failures.....:o

Watch the Simpsons much?

Remember the BEER BARON episode?

Remember when Homer was rolling around on the lawn, engulfed in blue flames?

if I was going to rely on a brazed seal in such spectacular fashion I would make damn sur I drilled and tapped the joints first so as to ensure that any failure wasn't what they call a


and while I was at it I would change my name to "THE INCREDIBLE, AMAZING, DEATH-DEFYING HERMES" and sew myself a pair of black and yellow tights and paint a crash helmet to match!


truth be told, I am thinking that having chlorine in my system might be a bit safer than having H and O in it, at least if there was going to be an accidental combination that overall pressure would decrease quickly, not that you would be decomposing at ambient temp anyway (for reasons you have already stated)

edit---ambient Pressure not ambient T

of course your probably right though!

[Edited on 1-12-2003 by Hermes_Trismegistus]


Organikum - 29-11-2003 at 05:40

Any pot made of solid steel or iron works, the thicker the better. Put an insulation of concrete on the outside then some glasswool on the sides. Heat up slowly.
Use a iron net as diaphragm or higher voltages and wider distances between anode/cathode.
Use NaOH.


Hermes_Trismegistus - 29-11-2003 at 12:16

Well if you are going to do the deed, at least use a pourable refractory material.

Or, in a pinch, you could even use common brick mortar mixed 2 parts to 5 parts with Pearlite (common inert potting soil additive).

However, this mixture would calcine after a few uses(temp dependant), causing the insulation to lose cohesion and become brittle and flaky.

Whatever you choose to do, don't use a heat sink like concrete!

I mean c'mon man! Why not just add cooling fins!



I am also pretty durn sure that the perforated Fe screen is more of a necessity than a performance option!

[Edited on 29-11-2003 by Hermes_Trismegistus]

BromicAcid - 3-12-2003 at 09:28

I don't know if anyone else has seen this website but when I found it I knew it would be invaluble to this thread.

It's the "Molten Salt Database -Eutectic Finder" you can type in NaCl and get dozens of eutectic mixtures, some as low as 50 C. FeCl3 and NaCl has an eutectic at 151 C, that's reasonable, KCl-NaCl-ZnSO4 (1:1:1) eutectic at 290 °C. You would have to look at the reduction potentials to see what metal you would get but I see this as quite the way to look into other eutectic mixtures.

I am a fish - 4-12-2003 at 00:57

Unfortunately, both zinc and iron have higher reduction potentials than sodium, and so these eutectics wouldn't be suitable.

Theoretic - 5-12-2003 at 07:22

"351 CO(NH2)2-NaCl 90-10 112 C"
This one should work. Although I don't understand, what does mol percent mean? In this case, 0.9 mols of urea and 0.1 mols of salt?

BromicAcid - 6-12-2003 at 14:28

Looking up more of the recent subject matter on electrolysis of molten salts to get sodium metal I'm finding that in almost every case the sodium metal is the cathode, eg. they add sodium to the melt before hand and run the current though the sodium metal where it collects a alumina ring. However, here is one patent that sounded kind of nice.

91781n Electrolytic method of recoving an alkali metal from a molten salt of the metal. Kummer, Joseph T.; Weber, Neill (Ford Motor Co.) Ger. 1,783,137 (Cl. C 22d), 22 Aug 1974, US Appl. 507, 624, 22 Oct 1965; 4 pp. Division of Ger. 1,596,077 (See Fr. 1,491,674, CA 68: 92484r).
Molten alakali metal is electrolically sepd. into 1 arm of a U-tube device, the other arm of which contains the molten salt, the 2 arms being sepd. by a solid electrolyte of cryst. structure of which contains essentially Al and O ions in its lattice. Thus, the 2 arms of a U-tube are sepd. by a sym. positioned plug of Na B-albumin [12005-48-0] e.g., Na2O.11Al2O3. In 1 arm an eutectic mixt. of Nano3 and NaNO2 is held at 245 C and molten Na [7440-23-5] is introduced into the other arm. Electrodes are inserted into the melts and connected to a source of d.c. whereupon Na+ passes through the solid electrolyte.

Just think it's neat that NaNO3 is used and such.

One more thing, about my earlier thought about electrolysis of Lithium Chloride in ethanol. It would be interesting that if lithium was produced that it would immediately react with the ethanol to produce the alkoxide. At the anode chlorine would be produced which would react with the ethanol to produce CH2ClCH2OH but it would be nice if there was some catlyst to produce CH3CH2Cl from chlorine and ethanol because then it could do a Willamson Ether Synthesis with the alkoxide to produce Diethyl Ether and lithium chloride which would react back with the electrodes and basically serve as a catlyst to make ether. Also, the HCl produced as a byproduct of the chlorination of ethanol would react with ethanol being that it's anhydrous and make CH3CH2Cl and H2O so you would end up with all kinds of crazy stuff.


Organikum - 8-12-2003 at 12:39

have you ever seen the picture on page ONE of this thread showing a Castner-Tiegel?


A Castner-Cell wont spill any Cl2 - thats one of the reasons I propagate its use.
Put the cell in a metal bucket if you have fears but be assured any molten NaOH will immediately get solid as it escapes the cell - the temperature isnt so far above the melting point at all in the NaOH process - another reason I propagate this way.

I have to repair my famous "MEGAFLAMER" now so excuse me please...... :D

Das Lernen ist difficult.The Geschmack meiner eigenen Schuhe ist schlecht. Mein Kiefer ist wund.

Hermes_Trismegistus - 8-12-2003 at 15:23

Do ya feel like posting some more about that Castner-Tiegel cell?......Organikum?

When I searched the net...the top two posts were you and the third was in German.

[Edited on 8-12-2003 by Hermes_Trismegistus]


Organikum - 10-12-2003 at 11:00


Orgy: your electrolytic cell may work marvelously, but does it glow with heat and belch burning hydrogen? Does it?! I thought not.

In Europe iron doesnt glow at about 330°C. Of course hydrogen is produced and should be burned directly at the outlet of the bell - this is favorable in special if heating is done with an open flame.

Hermes you nitpicking nervesaw: You dont need pure nickel for the electrode. You can use plain copper, or nickel plated on copper or nickel plated on whatthefuckever, or even plain iron, you can even leave away the diaphragm and use the iron pot as electrode. Also your fingers may work - but I guess only for a VERY short time - it for sure makes an EXITING experience.

OK. I dislike this but as everyone seems to want the recipe to enable retarded idiots to cause trouble:

- Take a soupcan and burn the paint/laque away with a propane torch. Clean it insides. Put some concrete around the sides - at least 3cm, reinforce the concrete with some steelfabric/wire or similar. Put it on a hotplate/gasstove - on the gasstove put an steelnet between flame and can to prevent holes being burned into the bottom. Fill halffull with dry NaOH fresh from the box. Heat it till the NaOH melts and regulate the temperature just to hold it in liquid stage - dont overheat! Put one SS nail in the can and one coppernail. Connect them to a car-batterycharger or better to a car-battery with charger attached. Over the coppernail you have to put a iron-bell with a little hole on top for the hydrogen to escape. Plug in charger. Run. Unplug charger and stop heating. Your bell is full SODIUM now.

And thats more than close enough IMHO. :mad:

- Please dont ask whats the + and whats the - pole.
- Please dont ask how to determine the distance between the nails.
- Please dont ask how to determine if the Castner-Tiegel is actually working.
- Please dont ask how to make the bell.
- Please dont ask how to attach the electric cable to the coppernail.

As if you cannot answer this for your own you definitivly shouldnt even think on making sodium metal.

Regarding pictures - see my "I am back" thread

[Edited on 10-12-2003 by Organikum]

Marvin - 11-12-2003 at 08:39


Quick question, if the bell is electrically connected to the copper nail, as the description implies... What stops the sodium forming on the outside of the bell closest to the other nail rather than on the copper nail?

I have no information on overpotentials for copper, iron, or sodium in molten sodium hydroxide or for that matter on the conductivity of molten hydroxide but Id be worried that even if this is in the right direction that it wouldnt be enough to prevent problems.

I dont like the idea of using a car battery either, they can push a lot of current at 12v and its only power you have to replace with the charger anyway even if this can be done after the experiment.


Organikum - 11-12-2003 at 11:14

- the bell isnt connected.
- the power provides internal heating
- of course there should be an isolated iron lid covering the cell to avoid H2O sucked from the air.
- the battery works mostly as a capazitator equalizing the usually pulsed output from the charger.
- if I would answer your problem with the current, I would answer one of my "Dont ask" points, sorry.

[Edited on 11-12-2003 by Moonmonster]

[Edited on 11-12-2003 by Organikum]

BromicAcid - 11-12-2003 at 12:55

Hey! How's about the electrolysis of molten sodium cyanide!

2NaCN ----> 2Na + (CN)2

Supposedly it works but I really don't fancy the idea of working with cyanogen and I don't know much about the behavior of molten cyanide salts. 563 C is'nt too attractive a melting point either. But would the cyanogen react back with the sodium metal? If not you could enclose the whole apparatus and run the exit gas though a solution of sodium hydroxide to convert back just run the reaction till all electolyzeable material is consumed and you might have a big jolly pot of sodium.

Back on the Castner process. I think that I will attempt to make said apparatus so I figured the best place to start would be some good electrodes so I picked up some nickel-silver rods, turns out that is just a term used and the rods are like 65% copper and 12% zinc with the remainder being nickel and molybdenum so just a warning to others that might be attracted by them, I don't think they will hold up to well. But they do sell titanium rods and I have heard good things about its corrosion resistance to molten hydroxides. :D I found a nice picture of a Castner cell in a book at my library and made a photocopy, it details all the parts in a seperate paragraph nicely although I liked the origional one on the first page better. I think that I might post a the picture of this one later though.

vulture - 11-12-2003 at 13:09

I would suspect the NaOH would dissolve the TiO2 layer and form sodiumtitanate, but I may be wrong...

BromicAcid - 11-12-2003 at 20:38

Like I said, a picture of another, Castner Tiegel. This one with additional explantions.

The sodium hydroxide, contained in an iron pot set in brickwork, is melted by means of a ring of gas jets placed underneath; and kept about 20 C above the m.p. (318C) of sodium hydroxide. The cathode, H, of nickel or iron rises through the bottom of the iron pot, A, Fig. 8, and is maintained in position by a cake, K, of solid sodium hydroxide in the lower part of the pot. THe anodes, F, several in number, are suspended around the cathodes from above. A cylindrical vessel, ND, floats in the fused alkali above the cathode, and the sodium and hydrogen liberated at the cathode collect under this cylinder. The hydrogen escapes through the cover and the atm. of hydrogen in the cylinder protects the sodium from oxidation. A cage of nickel-wire gauze, M, seperates the anode, F, from the cathode, H. From time to time the sodium D, is skimmed off by means of a perforated ladle, which retains the liquid metal, but allows the molten hydroxide to flow back. The oxygen liberated at the anode escape via the vent P.

Simply a beautiful process. One day it will be mine. As for vulture's comment on the use of titnium I didn't even think of that, I was more worried about hydrogen embrittlement, titanium is strikingly succeptable to that.

Also, I found "Sodium chloride mixed with powdered lead heated red hot in a closed retort gives metallic sodium:
2NaCl + Pb -----> PbCl2 + 2Na
I guess I would list that just for completeness seeing as how our current research into aluminum is paying off so well.
Also, while I'm on a roll, "Upon heating a molten mixture of sodium carbonate and sodium cyanide at a temperature of 1000C, metallic sodium, carbon monoxide, and nitrogen are formed."
2NaCN + Na2CO3 ----> 4Na + 3CO + N2
[R. Franchot, Ind. Eng. Chem., 16, 235 (1924)]
I also found references to sodium carbonate being reduced by both magnesium and aluminum turnings to yeild sodium metal, which I believe has been hinted at if not outright stated already.

Marvin - 12-12-2003 at 10:12


I'm not convinced the aluminium method is worth following up, for reasons Ive explained.


I concurr with the worry about titanium in molten hydroxide.


Then there is a problem, becuase if the bell isnt electrically connected to the copper nail at the start of the reaction, it soon will be. A cell where both electrodes enter from the top seems to require a molten hydroxide stable insulator just like the other designs, and I havnt been able to find one I can use. The one good suggestion that was made to me was sintered magnesia containing 1% or so waterglass, which only reacts with the hydroxide very slowly, but I lack a suitable furnace to make it in.

Any idea what insulators the Castner Tiegel cell used?

It would also be nice to know how big the holes in the ladle need to be for surface tention to retain the sodium, but allow the hydroxide to pass through.

If the battery is just for smoothing the current, then it is not useful in this application. If the cell has 12v over it then too much power is being wasted for the current that is flowing and it needs redesigning.

Hermes_Trismegistus - 12-12-2003 at 10:38

Look on the first diagram on the first page, The Holzring is the anode insulator, its made of wood.


Organikum - 12-12-2003 at 11:29


Then there is a problem, becuase if the bell isnt electrically connected to the copper nail at the start of the reaction, it soon will be. A cell where both electrodes enter from the top seems to require a molten hydroxide stable insulator just like the other designs, and I havnt been able to find one I can use. The one good suggestion that was made to me was sintered magnesia containing 1% or so waterglass, which only reacts with the hydroxide very slowly, but I lack a suitable furnace to make it in.

No really not. Hole in bell with filed down nail through it fixed by ovencement as no molten Na should reach the holes for hydrogen escape would be blocked and boom it makes.


It would also be nice to know how big the holes in the ladle need to be for surface tention to retain the sodium, but allow the hydroxide to pass through.

Eh? surface tension? I guess this is an misunderstanding. Hydroxide passing through what? Why?
Refill: Take plier. Open lid. Throw in NaOH. Close lid. Do fast.

for the rest:
no. no. no. no.
current efficiency in a quick and dirty tomatosoupcansetup?
no. no. no.
No. sorry.

Marvin - 14-12-2003 at 18:10

Hermes, there seems to be something inbetween the solidified melt and the holtzring. Its just labled isolation in what I can see, but it must be molten sodium hydroxide stable and an insulator.

Wood is attacked by molten sodium hydroxide and starts to decompose over about 270C anyway. I dont think we can use this.

From the description it seems the molten sodium will short circuit the nail and the bell long before the bell has filled with sodium. I dont understand how this is prevented.

One of the features of the castner process is that the sodium is removed by a perforated ladle that retains sodium while draining hydroxide. It would be nice to know what hole sizes are needed for this.

Adding a battery is only useful over a narrow range of cell voltages. If the cell is not designed to work over this range, the battery will not help. If it is designed to work over this range then its been designed specifically to work at a voltage that wastes power. Its a lose-lose situation.

Bromic acid,
Very nice going. If you are willing to swap magnesium for sodium potassium etc, and I would be if I could get magnesium cheaply, then you might prefer the organic solvent method in the patent. Decent yeilds and its easier to get the metal as a single ingot relativly free of magnesium oxide/hydroxide.

Aluminium and sodium hydroxide is supposed to go mainly to Na3AlO3, a mixture that was formerly used to defrost oil wells and which produces no sodium directly, formation of sodium is a side reaction with the coproduction of sodium metaaluminate. Magnesium seems the only feasable metalic reducing agent which limits most peoples ability to use this.

Kaboom, now that would be hydrobromic acid wouldnt it ;)

[Edited on 15-12-2003 by Marvin]

molten sodium

Organikum - 15-12-2003 at 05:16

seems to be a bad conductor for electricity, Marvin.

Marvin - 15-12-2003 at 08:57

Molten sodium is one of the best liquid conductors.

Solid copper around e-8, liquid sodium around e-7 liquid mercury about e-6 (perohmpermeter). Its undoutably much better than the molten salt its displacing/floating on.

Theoretic, based on what?

Organikum - 15-12-2003 at 12:39


but it works nevertheless, also I never thought on this ??


hodges - 3-1-2004 at 14:59

I know this thread is about making sodium, but calcium is a similar metal. I found out that CaCl2 actually has a slightly lower melting point than Ca does. Calcium also does not burn nearly as readily as sodium does (in fact, I had a hard time getting a small hunk of Ca to burn even with a blowtorch). Yet it still reacts fairly well with water. And CaCl2 is cheap and plentiful (dehumidifier powder or ice melter).

I would think Ca would be significantly easier to make than Na. Might be a good first try for someone who has never made Na.

I tried making sodium from lye when I was 14. I heated up some in a large spoon, using the spoon for the anode and a wire for the cathode. As I recall, I got vigerous bubbling (much more so than when electrolysing water, for example). The hydrogen kept exploding and blowing out the alcohol lamp I was using to heat it. I had a small gray lump near the cathode that I tried putting some into water with a tweezers. Nothing appeared to happen. But when I touched the lump again with the wet tweezers there was an orange spark. Have never tried making it again since then.


Tacho - 5-1-2004 at 05:32

Yes, calcium would be easier to make than sodium.

I remember reading a description of the setup to make calcium in the lab. It involved the use of a screw-like iron electrode which would be lifted from the molten CaCl2 as calcium was being deposited, leaving only a tip of calcium in.

I'm traveling now and can't read my books to find out why can't you just leave the deposited calcium in the molten salt. This would even prevent oxidation.

When I return home, I will find out and post. Maybe we should start a "Calcium!"

Although you can use calcium to dry solvents and make some interesting reactions, sodium BURNS, or even EXPLODES with water! UNDER water! You can't beat that! A "Calcium!" thread sure would not be as popular as this one.

The_Davster - 6-1-2004 at 16:23

A bit off topic but how much shouldc a graphite electrode cost?, I saw one today for 12$ CDN. It is about 6mm wide by 15 cm long

Saerynide - 7-1-2004 at 03:00

Wow, thats really expensive. I got a pack of 6 of carbon rods about 1 cm x 15 cm for less than $2.


I dont live in the US, so you wouldnt know were I got them from anyways :P

But if you insist, I got them at a local shop that sells plants stuff.

[Edited on 8-1-2004 by Saerynide]

Get it while it's hot!

BromicAcid - 12-1-2004 at 18:10

I don't know how long I'll leave this up, maybe someone can convert to .pdf and host it but for now here it is, go to the page and gander at the scans I made from "Sodium: Its manufacure, properties, and uses" It is the best reference Ive come across yet regarding both chemical and an overview of the electrochemical methods to this wonderous element. I have the whole chapter on the production of sodium metal scanned in. Like I said, get it fast, download it, whatever!

[Plus when I take it down I have to delete this post so that my last post is not a lie.....]

[Edited on 1/13/2004 by BromicAcid]

Polverone - 12-1-2004 at 21:15

Ahhh, I remember reading that book some months ago. BromicAcid, do you have higher-resolution scans on your computer, that OCR would work on? Whether you do or not, I'll gladly host a PDF made from the scans (someone else will need to assemble the PDF as I don't have access to Acrobat at the moment).

Theoretic - 13-1-2004 at 05:38

On the topic of separation of sodium from carbon that fell off the electrode... ...sodium forms sodium carbide with carbon, separation must be quick. An easily meltable salt could be used to do that, it would sink to the bottom, and carbon would sink to the bottom of it, while sodium would float on top. Also sodium carbide would dissolve in the molten salt.

sodium production chapter converted

Polverone - 22-1-2004 at 13:50

BromicAcid: the images you had on your website have been turned into a PDF and are now in the <A HREF="">library</A>. If there are other materials that you want to scan and put online in the future, I suggest scanning in black and white at 300 or 600 DPI (or scanning in grayscale and then converting to black and white) since it's cleaner, compresses down to a smaller size, and looks better on screen/in print. But thank you for making that chapter available the way you did.

KABOOOM(pyrojustforfun) - 5-2-2004 at 21:22

Polverone:<blockquote>quote:<hr>BromicAcid: the images you had on your website have been turned into a PDF and are now in the library<hr></blockquote>I downloaded it. it gives this error when I open it:
<i>There was an error opening this document. The file is damaged and could not be repaired.</i>

(Thank you in advance)

you are quite right

Polverone - 5-2-2004 at 22:18

The file I had locally was twice the size of the one in the library. I guess the upload got interrupted before. I have verified that the file now in the library can be downloaded and viewed.

Castner Tiegel remark

Organikum - 8-2-2004 at 13:32

I was irritated all time because of Marvins remark that the molten sodium would give a shortening of the electric current between the "bell" and the central electrode and make the design unworkable.

But I had done it and it worked.
How could this be?

Solution: The iron bell works is cooled by the air. At the edges inserted in the molten NaOH the NaOH solidifies and forms such an isolator on the bell. This works because the temperatur of the NaOH is only slightly above the melting temperature of NaOH and the superior heattransfer properties of iron compared to NaOH.

Try it!
Just put something from iron into just molten NaOH and you will see the solidified NaOH forming a stable protection layer on the iron. And this layer wont dissolve also not after times.

Nothing shortens here.

Hope this answers your question Marvin?

Electrolysis of Sodium Hydroxide

BromicAcid - 4-4-2004 at 12:18

  1. Picture shows my nickel crucible with some sodium hydroxide pellets in it. Anode is just the crucible and the cathode that decends into it is nickel as well. Below it is the torch that I was using to heat it.
  2. Too awhile for everything to melt. The hydroxide pellets on the top solidified together and formed a crust, I didn't realize it was all liquid underneath till I poked at it with a nickel rod, it all caved in and melted together. The solution was white till it all dissolved then turned crystal clear. Electrolysis was begun using a 9V battery.
  3. Very soon after this there were bubbles at both the edge of the crucible and at the cathode. No sodium though, but the solution darkened to a yellow color.
  4. Thinking that maybe the solution was too hot and that any sodium being produced was reacting back with the hydroxide at this temperature I turned off the heat. A crust quickly formed over the top that bubled up due to gas formation. Some sodium may have been formed but I could not tell. After awhile the bubbles stopped so I turned the heat back on.
  5. Not long thereafter I decided that maybe it wasn't a good idea to use the crucible as the anode so I grabbed a piece of nickel rod and put the clip on that and manually held it in place. Bubbles started to rise immediately but the hydroxide solidified around it and it took awile to take on enough heat to where it conducted again. You can see a black crud around it, I believe the sodium was going into solution and where ever the hydroxide was solidifying it was darkening more and more.
  6. The hydroxide solution turned almost black and no longer produced bubbles. The metaloid that has been mentioned. I discontinued electrolysis and heating. I let it cool a bit and added water, very violent reaction but it was a block of hot hydroxide after all.

During the whole procedure I saw no sodium gobules. But as the solution darkened I was sure there was some around there somewhere. The nickel crucible came out with no noticeable errosion but it had a black residue on it. What's everyone think, was the NaOH too hot, too much voltage, not enough, current, electrolysis really isn't my strong point. Also this cell really isn't anything like the one that I've been working on, this was mostly just a proof of principle thing makes it even worse that it didn't work for me.

[Edited on 4/4/2004 by BromicAcid]

hodges - 4-4-2004 at 12:41

I think if by "9V battery" you mean one of those tiny things a few centimeters square and about 1 centimeter thick, it did not have enough capacity to do much electrolysis. A power supply is best. Lacking that, use a 12-V auto or motorcycle battery. Lacking that, use a 6V alkaline lantern battery.

chemoleo - 4-4-2004 at 13:21

I have also tried NaOH electrolysis a number of times (under conditions as you describe), with the same results as you. A blackish residue was left, that violently reacted with water.
I used a powersupply, with as much amperage as it could take (the amperage being restricted by the resistance), @ 12 V.
Never any success that way. I believe it can't be done this way, this has to be done in a properly designed vessel, with proper temp. control, etc.

Marvin - 5-4-2004 at 00:26

I had several sucesses with small amounts of hydroxide and one failed attempt when I tried to scale it up (before I had a good idea how.

Hodges is right, you arnt getting enough current from the battery. I used a 12v PSU capable of delivering about 7A. The voltage was much higher than it needed to be, the PSU sounded almost like it was directly shorted when running. 7A though is the very minimum Id consider using as a proof of concept. A car battery or motorcycle battery might deliver far too much cuttent for too short a time to be useful. Formation of sodium was very slow at 7A though its halo was impressive, best part of the experiment IMHO.

Ideally you should be pumping almost almost enough power into the cell that it stays molten without external heating - this makes it much easier to keep it around its melting point. I would also expect an induction period before sodium starts to form as the mixture dehydrates.

Also watch out for the small explosions as the sodium sets the hydrogen bubbles off. This tends to send sodium metal flying round the room and limited the amount I could let form as a globule before I had to remove it, or lose it. Eye protection a must, though I think you know this. Thick gloves also as the fumes tend to attack skin.

[Edited on 5-4-2004 by Marvin]

Tacho - 5-4-2004 at 03:35

Originally posted by hodges
I think if by "9V battery" you mean one of those tiny things a few centimeters square and about 1 centimeter thick, it did not have enough capacity to do much electrolysis. A power supply is best. Lacking that, use a 12-V auto or motorcycle battery. Lacking that, use a 6V alkaline lantern battery.

I agree. You need much more current than this batteries can provide. At least a couple of amperes (edit) considering the size of your setup.

When I tried to do an open air electrolysis of NaOH, I could not get any metal due to the fact that sodium burns at this temperature, but I could see the little orange sparks in the cathode. Can you see those?

You seem well equiped, can't you built the "bell" proposed some time ago in this thread?

[Edited on 5-4-2004 by Tacho]

BromicAcid - 5-4-2004 at 07:20

Actually I bought that nickel crucible for use as my bell in the larger setup that I've been working on. It will hold about 1 Kg of hydroxide and I am going to purchase a good power source. I basically just wanted to see what molten hydroxide looked like and how it flowed etc before I did my massive cell. So I guess the consesus is that I should have massive amps. Not just a car battery charger or something of that sort, right?

Actually the power supply is the last part that I need. So I'm glad I tried a proof of concept, I was just going to go with my car battery charger 12 V 1 Amp.

Also I got no sparks at the cathode, just a few bubbles that solidified in the air and re-melted shortly thereafter. I also vouch for Organikum's point further up thread that inserting metal into the molten hydroxide instantly coats it with a solid coating that prevents electrolysis, I had to stir my electrodes around until all the hydroxide had melted off and go deep. I'm sure that if they had remined stationary the coating would have remained and insulated.

Saerynide - 5-4-2004 at 09:02

May I ask why the nickel anode doesnt get oxidized?

BromicAcid - 5-4-2004 at 10:55


May I ask why the nickel anode doesnt get oxidized?

My nickel anode did get a slight grey tinge to it probably from it being oxidized slightly. But nickel is probably the one of the most resillant metals to hydroxides, it's just that the oxygen being produced there is doing a number on it. I would guess that the oxide will go into the melt as it is slowly formed and thereby allow the cell to continue electrolysis.

In the Castner tiegel on the first page I believe the anode is copper also I have heard of iron being used as an anode. However I have not heard of graphite being used. Probably disintigrates quickly, contamination possibly from flaking off? I don't know, but basically to answer your question the nickel anode does oxidize, albiet slowly.

[Edited on 4/5/2004 by BromicAcid]

chemoleo - 5-4-2004 at 12:36

yes, the graphite electrodes essentially are eaten up by the hydroxide. Just like when you electrolyse a NaCl solution.
I experienced both first hand, and moved on to different electrodes thereafter.

PS on the note of little explosions happenenig upon the electrolysis of molten NaOH - the same happened to me. Beware of that, molten NaOH on the skin is not pleasant

[Edited on 5-4-2004 by chemoleo]

darkflame89 - 9-4-2004 at 00:25

Dunno about this, i am new to electrolysis anyway. If have 2 bowls of aqueous NaOH, of which the anode side has a copper electrode, and the cathode side also copper electrode, and the entire setup is connected to battery supply, will the get Na metal at the cathode in the end. Since copper dissolves at the anode to form Cu(OH)2 and H2 gas is given off at the cathode, won't i get Na in the end??!!

If there are any flaws in the system, please enlighten me, and i will go back to the thermite mixture setup.

Marvin - 9-4-2004 at 05:33

Read the whole thread Darkflame, if there is any water sodium metal will not be produced. This is why molten NaOH must be used.

ballzofsteel - 9-4-2004 at 11:58

This is an interestinng patent describing
electrolysing sodium nitrate with sodium carbonate at lower temps.
A good read anyhow GB440139 .:)

patu - 9-4-2004 at 22:17

I obtained sodium from NaOH by using a battery charger. I used a soup can as the anode. For the cathode, I used an iron wire with a small loop at the end. I melted the sodium hydroxide on a hot plate and plugged in the charger. I dipped the loop on the surface of the melt. Liquid globules of sodium started forming and popping with an orange flame. After the liquid sodium filled up the loop, I took the loop out of the melt and tapped it on the side of a glass filled with mineral oil. Using this method, I produced a few dozen balls of sodium the same size and shape of B.B.s

ballzofsteel - 10-4-2004 at 03:42

Heh,I swear Ive heard this before.

Tacho - 10-4-2004 at 05:48

patu, I liked that.

You mean that the liquid sodium got trapped under the molten sodium surface, in a loop, like solder does to copper wire?

What size was your loop? Could you elaborate?

Edit: Have you tried multiple loops?

[Edited on 10-4-2004 by Tacho]

patu - 10-4-2004 at 10:20

The sodium formed around and inside the loop. When the loop was removed from the sodium hydroxide, all the liquid sodium on the outside of the loop collected in the middle. the wire was still very hot so the sodium stayed molten long enough to be shaken off in a cup of mineral oil. The loop was about a centimeter wide, and yes it's just like a bead of solder. I really should try multiple loops, but lately i've been trying to isolate potassium.

The_Davster - 10-4-2004 at 16:19

Patu:What voltage and current were you using to get the sodium?

[Edited on 11-4-2004 by rogue chemist]

patu - 10-4-2004 at 20:51

I used a 12 volt battery charger to make the sodium. When the sodium was being produced, the amperage needle jumped all the way to 15 amps!

Finally, Sodium from Sodium Hydroxide for me!

BromicAcid - 18-4-2004 at 16:35

  1. This picture was taken when I turned off the voltage for a minute. What I had done to this point was take a quantity of sodium hydroxide and place it on a watch glass. Electrodes were inserted into the mass and it was spritzed with water. When I saw a few bubbles spring up from the electrodes I just sat back and watched. As you can see eventually the sodium hydroxide melted in the middle and made a shell around itself to contain itself in the middle of the watch glass. For the power supply I used a car battery charger. I ran it at 12 V and when it melted the current was only 1 A or less. When electrolysis was commenced again bubbles continued to come to the surface, electrolysis of residual water probably, NaOH holds a firm grasp on some of it.
  2. Lower left-hand side, that little glistening spot, sodium! At first it looked like green coming up from the cathode and I thought that the nickel was corroding and salts were flaking off but after it moved away from the frothing it had the color of mercury like the time I made sodium by electrolysis of NaNO3 using soda glass as the membrane between the two. The gobule did not burst into flames though until provoked by jabbing at it with a separate nickel rod.
  3. Several gobules now, they grow in number, huzzah! The current had risen to 5 A by this time but was holding steady. If the electrodes were submersed in the melt more the amps would have risen but it even worked with just the tips in. Electrolysis was discontinued at this point as I had shown myself that the battery charger was sufficient to produce sodium.
  4. The Aftermath - See that burnt spot, that is where the watch glass that I did this in rested. Didn't expect that, also the watch glass survived surprisingly intact, it only had a depression in the bottom directly beneath were the electrodes contacted the melt. All in all I think this shows that the solid hydroxide is an excellent material from which to furnish the protective layer in a sodium production vessel.

I'm just glad that it worked. Finally some hope that my castner cell might work. I've got a picture of my progress

Saerynide - 18-4-2004 at 22:33

Wow :o Sweet :D

So the slightly moist NaOH will melt itself during the process?

Now I really want to try this sometime :)

Tacho - 19-4-2004 at 11:04

Congratulations Bromic!

Your work seems clean and easy! I tried melting some NaOH on a hotplate last weekend to try Patu's loop idea and all I got was NaOH all over my benchtop, my hotplate and my arms.

I realized that the orange sparks might be trapped hidrogen popping. Na ions all over makes the spark orange.

I'll try this as soon as possible.

BromicAcid - 19-4-2004 at 13:09

The slightly moist NaOH allowed a current to be established, I just piled some prills around both of the electrodes and between them and used a spray bottle for my hair just to put a sheen on them, not saturate them. Once the connection was established it generated enough heat to melt the NaOH but did not require any additional outside heating.

Actually I might run my castner cell with KOH in place of NaOH. My most recent reading has shown that for NaOH the yeild is usually 27% but with KOH it is 55%. However I really don't want to store potassium in that it forms those superoxides that can explode when you cut into them from being stored under oil. Well it's just still under consideration, let's put it that way.

[Edited on 4/21/2004 by BromicAcid]

BromicAcid - 20-4-2004 at 18:00

Here is a design I was thinking for collecting the sodium during the use of a Castner cell. The main part of it would be an inverted petrii dish so that the open part would be directed down toward the melt. Now this part would not have to be made out of glass but that would make it easier. Because it would only be dipping into the top of the hydroxide I would hope that it would not be attacked too readily. From this would lead a glass tube straight out from the side. This would follow along just an inch or less above the hydroxide. This would go up and over the rim of the cell and go into a jar. The jar would have that tube going in and another tube going out, applied to the tube going out would be a hand vacuum pump.

The purpose of all of this. The dish would be over the cathode beneath the melt. Vacuum could be applied to draw a small amount of the hydroxide into the dish a tiny amount and as more sodium is produced vacuum is increased to pull up more sodium that would be formed. Being close to the hydroxide @300+C the tube would keep the Na nice and molten and at the end it would go in to the jar. In this way the collection of sodium would exceedingly easier.

Other Stuff:

A few patent numbers have been thrown around in this thread but here are a few more US patent numbers:

452030 is Castner's original US patent for his cell and ideas behind it, he emphasizes strict temperature control. Also he mentions the perforated spoon, it seems like it is just full of tiny holes and the surface tension of the Na keeps it in the spoon. I might make a couple spoons for use with my cell. But realistically I was expecting to use a ladle.

517001 is a patent for the production of both nitric acid and sodium or potassium metal from nitrates by electrolysis. Also patent 590826 issued to the same person later shows interesting schematics for a porous diaphragm for separating cells. Really though, I never thought that one could perform electrolysis on NaNO3 if the sodium ended up in contact with it, figured that Na and NaNO3 might lead to an explosion but the patent states that some of the Na might reduce the nitrate to the nitrite but it is not a major occurrence. Also the patent calls for everything to be made of aluminum, sounds like a job for axehandle ;)

What's the difference?

chemoleo - 21-4-2004 at 16:36

Bromic, your method doesn't seem all that different, except maybe the water addition.
I dont understand why you'd get metalic Na globules - everyone who's tried it (inc. myself), one gets a dark grey mass, with the occasional yellow explosion from igniting H2/O2/Na.
It's not like your temperature is low (judging by the mark on the wooden bench :P), or that there is an oxygen deficiency in your local atmosphere (you wouldnt be writing this otherwise :P)
But seriously, I can't figure out why you'd get metalic sodium, as a nice little globule.
How about upscaling this?

BromicAcid - 21-4-2004 at 16:57

Anode Material: Nickel
Cathode Material: Nickel
Heating Method: Current Heating
Distance between electrodes: 1 - 1.5 cm
Estimated depth of immersion: 2 - 3 mm
Hydroxide present in both solid and liquid state
Humidity: 83% ish
Air Quality: (Sensitive individuals should stay inside)
Hydroxide Source: Crappy OTC Red Devil
Power: 12V 5A at time of production
Observations: Initially tiny tiny bubbles produced at the top of the pellets where contact was made with rods. Amps were at 0 on the meter. Hydroxide around cathode took on a grayish tint, I continuously piled the hydroxide back between the electrodes. Hydroxide at the edges did their disquesence thing and liquefied on their own. Eventually the middle are liquefied and I thought that it had just taken on too much water from the air. Evolution of gas started to increase rapidly. At the outside edges solid hydroxide, middle molten. Gas evolution continued for about 8 - 12 minutes and then eventually started to yeild the gobules. Amps were at 5 at this time. The sun was about 3/4 of the way to the horizon and there were no spectacular planetary alignments in the sky. My power source was a battery charger as said before.

So what was different?

One other interesting thing:
Composition (mol percent):
77.7 NaBr 22.3 NaOH Mp 260C

My first Sodium :D

The_Davster - 26-4-2004 at 20:48

Thanks to BromicAcid. I tried your method but my power supply was 2A max so it did not melt the sodium hydroxide, either that or I was not patient enough:P. I used the concave bottom of a pepsi can as the crucible and melted the sodium hydroxide over an alcohol burner. Nickel electrodes were placed in the melt and 12V @ 2A from a car battery charger powered the electrolysis. No sodium was evident for about a minuit but there was constant orange sparks. At around the minuit mark there was a crack which sent molten sodium hydroxide flying, and now a tiny globule of sodium was visible. Not grey, but like mecury, nice and shiny. Most of the sodium hydroxide had solidified by this time except for the 1 square centimeter between the 2 electrodes. After a while the sodium that was formed must have shorted the electrodes because there was another crack, more molten material flying, and then the electrolysis stopped. I did 2 trials and both times I got loud "cracks" when the sodium started and completed being formed. Strange.
A few questions now:
1. If I scale this up will I need to use more amps to maintain a good current density?
2. What is the best way to extract/purify the sodium formed in a small scale experiment like this?
3. What is the best liquid to keep the sodium under? Is there any problems with using xylene?

Out of curiosity BromicAcid, How much sodium were you able to produce each run? I was only able to get a fraction of a pea sized ammount each time.

[Edited on 27-4-2004 by rogue chemist]

Saerynide - 27-4-2004 at 00:38

I got a question. How do you collect the sodium? Pipette it?

darkflame89 - 27-4-2004 at 01:46

I was jus about to ask that too. Do you pull the electrode out with the blob on it and immediately plunge it under oil or xylene?

And, must nickel rods be used? I can't get nickel..:(

BromicAcid - 27-4-2004 at 07:16

A few questions by Rouge Chemist:

1. If I scale this up will I need to use more amps to maintain a good current density?
2. What is the best way to extract/purify the sodium formed in a small scale experiment like this?
3. What is the best liquid to keep the sodium under? Is there any problems with using xylene?

Out of curiosity BromicAcid, How much sodium were you able to produce each run? I was only able to get a fraction of a pea sized amount each time

1. If you increase the depth your electrodes go into the melt you will of course have to increase the current to increase the current density step for step. According to Organikum in his posts on the castner tiegel :
The distance between anode and cathode is only 2cm. The amperage at the cathode is about 1,6Amp/cmxcm at the anode about 1,1Amp/cmxcm. (Smaller units call for higher current density). The voltage should be about 4V to 5V then, the current density is the more important point.

2. As for extraction you can skim off the sodium with a spoon with small holes in it or without and drop in xylene or toluene or kerosene or mineral oil or something of that nature. Then after you get your fill you put it in, jeeze, I forget the exact name so you're going to have to settle for a phonic equivalent, linindale or something of that nature, regardless it is an inert substance, I think a paraffin of some sort with a low density around .56 and apply heat. The sodium will melt and float to the surface leaving hydroxide behind from there scoop it back up and store under your final storage solution.

3. I have no clue what is the best but xylene will work although it will still oxidize to a small extent.

For electrolysis of molten NaOH I was only able to get small amounts with each run. I have got better results with molten NaNO3 with soda glass as a membrane but it becomes brittle and can basically explode, I will do some work with NASICON materials this summer though, supposedly much better.

A question from Saerynide:
I got a question. How do you collect the sodium? Pipette it?

Procedures that I have saw call for a spoon full of tiny holes, the higher surface tension of the Na(l) keeps it in the spoon while the hydroxide spills out. Although that really is only suited to a larger scale. On a smaller scale you could take a metal rod and sufficiently cool it then touch it to the top of a gobule freezing it and lift it up and scrape it off into your protective liquid. Or you could pipette it, take your pipette and put it into your protective solution being held at a higher temperature then molten sodium to keep it liquid inside and suck it up and put it right back into the protective liquid.

And from Darkflame89:
And, must nickel rods be used? I can't get nickel..

According to Organikum's first post on what kind of rods should be used the conclusion is Cathode: Copper Anode: Nickel also I have heard of iron being used as the cathode and copper being the anode. I'm sure to some extent you have your choice of Copper, Nickel, Iron, Platinum (and other nobel metals). Although if you're only going for it for a short period of time then there should be no problem experimenting with different anode/cathode materials.

Hermes_Trismegistus - 27-4-2004 at 09:00

Originally posted by darkflame89
And, must nickel rods be used? I can't get nickel..:(

Yes you can, just U2U whatever member you are currently most friendly with, Coiling up a nickel wire and sending it through the mail isn't THAT pricey for an electrode that'll last you practically forever.

If you need I'll send it, but you are probably better off getting it sent from the states, (everything seems to be cheaper there)

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