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Author: Subject: Sodium!
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[*] posted on 3-9-2004 at 04:10


today I am going to try it with my variation of bromic's cell. I purchased a 10 dollar water heater for the heat source but I cracked the metal pipe the contains the heating wire and some dust like material (do you guys know what that stuff was?) Since my power source blows I an going to heat it additionally with a propane torch. I'll let you guys know how it turns out!
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[*] posted on 3-9-2004 at 07:40


Don't heat with a propane torch, the CO2 will react with your NaOH.



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[*] posted on 3-9-2004 at 07:55


Quote:

...some dust like material (do you guys know what that stuff was?)

Something ceramic probably, heating coils are usually embedded in ceramics. If you cracked open the heater to get to the coil, why didn't you just make your own coil in the first place? *confused*




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[*] posted on 21-10-2004 at 11:28


Success. I managed to make myself some drops of sodium using the wire slope setup. However I found out that the NaOH solidifies to fast using such small quantities and stainless steel eggcup.

By the way. Corrosion was massive and rendered the dark but clear NaOH a brown "mush". Stainless steel was used both as cathode and anode. (Cathode was made out of "steelwire" and it didn't seem to be zinkplated, I guess this is stainlesS?). Anyways, the anode (stainless eggcup) was heavily corroded and a yellow salt was formed on it's side above the NaOH surface.

I found out that the intense gas release ejected the formed droplets of Na in my slope which made extraction very hard. And further on I can tell you that sewing-machine oil is not inert to Na. :(

Need to get myself some charcoal lighter fluid. Thanks everyone for giving me the inspiration and providing information. I'll be getting back to this experiment in the near future.
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[*] posted on 3-4-2005 at 03:37


I've tried it again, this time with a nickle crucible and steel electrodes. NaOH melted very easily, seems Ni has a rather good heat conduction.

Electrolysis went smooth with lots of little globules and lots of corrosive vapor.

The only problem is that I now have a large black mass of NaOH/Na which is rock hard and I don't want to wreck my crucible.

I don't think melting under xylene will work as there's a massive layer of crystallized NaOH above it.




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[*] posted on 3-4-2005 at 11:49


That's a toughie.

Maybe add another salt to it when molten to decrease the solubility of the NaOH. Maybe melt under rock oil. Maybe try to heat and remove liquid Na. Nothing I can think of sounds pleasent.

BTW, I kept my 2 kg lump of NaOH I got from my Castner cell run and I use it for things like making aqueous NaOH solutions to neutralize chlorine and bromine and such. I tossed a piece into some water the other day and it turns out it was hiding a small piece of sodium which quickly made its way to the top and bursted into flames.




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[*] posted on 30-4-2005 at 13:57


would any of these battery chargers work?
http://www.argos.co.uk/webapp/wcs/stores/servlet/ArgosBrowse...
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[*] posted on 30-4-2005 at 14:43


Quote:

and some dust like material

Magnesium oxide.




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[*] posted on 2-5-2005 at 13:31


will either nichrome wire or copper wire react with NaOH?
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[*] posted on 2-5-2005 at 17:25


Nichrome might (as I recall, most chrome alloys are protected by a layer of chromium oxide), copper shouldn't. Copper also tends not to react with acids either, at least very quickly.

Tim
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[*] posted on 23-5-2005 at 12:53


THE RATE OF DISSOCIATION OF PERCHLORATE ION IN FUSED SODIUM HYDROXIDE

1,2 Ralph P. Seward, Harry W. Otto;
J. Phys. Chem.; 1961; 65(11); 2078-2081.

Quote:
Aluminum containers were used for the sodium hydroxide fusions in this work. When aluminum is immersed in fused sodium hydroxide, gas evolution from the surface of the metal occurs but this lasts
only a few seconds. It is proposed that the protective coating consists of a layer of an insoluble sodium aluminate since if the metal is removed, washed and then returned to the melt, the brief attack occurs again. While nickel is satisfactory in resistance to corrosion by fused sodium hydroxide at moderate temperatures, the unfortunate tendency of the liquid to creep up the walls of the container, to solidify when it reaches a cooler spot, is much more noticeable in nickel than in aluminum. Corrosion of the aluminum containers did occur but slowly enough so that they could be used for many hours with only a few milligrams loss in weight.
Yes... I may have guessed that aluminum may form a passive coating against molten NaOH, however I would have never truly expected it.
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[*] posted on 23-5-2005 at 15:13


Some time ago, I prepared several grams of sodium with the soup can method.
I melted approx. 200g NaOH in an old soup can (previously heated with bunsen burner to burn off the plastic coating on the inside) and connected the can to the positive pole of a laboratory power supply (5 Amps at about 12V- more current would be better).
The NaOH was melted (just melted- NOT heated further, this is very important!) and an iron wire, connected to the negative pole, was inserted about 1cm into the melt.
When the temperature is right, the Na collects as a single globule which is attached to the wire and floats on the NaOH. It grows bigger and bigger and then seperates from the wire and swims around. It is taken out with a previously heated spatula and immediately tossed into a small beaker filled with low viscosity paraffin oil (or xylene or BBQ lighter or whatever you prefer).

The temperature of the molten NaOH is VERY critical, only 10°C too hot and the sodium will simply dissolve in the NaOH to form a black conducting metalloid from which no sodium can be isolated. The heating should be very gentle and some solid NaOH should be present at the sides of the can. The NaOH should be as close at its solidifying point as possible.

This process is repeated several times and the paraffin is then heated until the sodium liquefies. The Na drops are melted together by gently stirring them with a bent piece of strong wire (this is not as easy as it sounds, but it can be accomplished). Bits of NaOH are removed by poking the liquid Na and removing the solid pieces.
By this method, one can collect about 5g of quite pure sodium a day, depending on the current and temperature.
There's no need for huge amounts of NaOH or special apparatus. Only a soup can, a good power source, some strong wire and patience are required.

I use my sodium mainly for drying organic solvents, only very small amounts are needed. 5g of sodium last a long time.

[Edited on 23-5-2005 by garage chemist]
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[*] posted on 23-5-2005 at 15:22


A pic of my entire setup (I used a computer PSU the 12V line was used- with 5V, not enough current was flowing).

[Edited on 30-1-2007 by chemoleo]

Img_0037.jpg - 54kB
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[*] posted on 23-5-2005 at 15:24


The molten NaOH

[Edited on 30-1-2007 by chemoleo]

Img_0038.jpg - 85kB
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[*] posted on 23-5-2005 at 15:27


The sodium pieces under paraffin, prior to melting.

I have no pics of the actual process, sorry.
I have some videos of the process (and melting together the Na, plus a video where the Na reacts with water, it caught fire), but they are way too big to upload.

[Edited on 30-1-2007 by chemoleo]

Img_0043.jpg - 46kB
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[*] posted on 2-6-2005 at 12:01
my own Downs cell


I am constructing my own Downs cell, The general construction is an "oven" consisting of the 1/4 inch thick steel rectangular trough 8" by 4" by 6" tall, (outside dimensions), and at the temp of 873K this radiates 4700 watts (black body temp). This trough is wrapped in ¼” asbestos.
The heating element is a bicycle chain in a coil insulated electrically from the steel by thin strips of asbestos. The dimensions of all this were chosen carefully so that all this could be surrounded by one layer of firebrick. This firebrick is surrounded in fiberglass and sheet metal, (to keep the atmosphere from entering this chamber) and is then surrounded by the 4” by 4” by 8” cement brick. The overall dimensions of this are about 2.5 feet cubed. I currently have no method of attaching photos.

My question is this, what ceramic will withstand a 575-585 C solution of nacl and cacl2, the ceramic is necessary to separate the steel trough (the cathode) from the graphite anode sitting in the center.
Second, will the calcium precipitate out from the liquid sodium and settle back into the bath where it dissolves in the molten salts or will it react with nacl and become Cacl2, thus requiring only the addition of Nacl?
My power source for the electrolysis is a 900-watt transformer that I completely made myself from two sets of E cores that I removed from two home backup DC to AC supplies. It can produce 4.5 vac at 200 amps or 9 vac at 100 amps, this is then rectified to dc. The steel bicycle chain draws 60 amps at 120 volts (I power this though a triac)
The Idea is to have the sodium run out a pipe into a one-gallon can full of oil. The shielding gas will be CO2 blown though another pipe into the trough, thus the CL2 will be blown out the top, I have the nastiest idea to collect the CL2 and compress it into a can at about 200 psi, at which the O2 N2 and CO2 can be released and the liquid CL2 will remain in the can. Will the sodium that will float on the Nacl Cacl mix be a mix of Ca and Na that will solidify in the runoff pipe? Or will the Ca sufficiently react with the Nacl and thus not be present in the floating sodium?
Most of the physical construction is complete, I have yet to get the steel welded together and the heat turned on full, I have proved that 2000 watts is all that is necessary to keep everything at 600C based on calculations at reduced power.
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[*] posted on 2-6-2005 at 15:26


Wow, sounds neato :D

Calcium is insoluble, heavier and has a higher melting point than sodium, and more than the electrolyte hopefully (hence using enough CaCl2 for the low melting point eutectic). It also has a somewhat more negative reduction potential than sodium, so it probably does react to produce sodium metal. Hence the calcium chloride ought to be inert. :)

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[*] posted on 2-6-2005 at 18:31


I believe that sodium produced in this manner has a calcium precentage of <1% the calcium does indeed crystalize out and fall back into the melt and this does react further with the NaCl there to produce more Na instead of calcium.

As for the ceramic material, alumina and magnesia based ceramics may work, the chlorine being the main culprit of damage here to ceramics so things are somewhere easier then using a castner cell where the melt will eat a number of ceramics and glasses. I had plans for a similar cell on my site but alas I am lazy and haven't started construction. My intention was to just use a high-magnesia pot from my gardening supply and invert it, drilling the hole in the bottom (which would be attached to the top of the lid) and having the carbon anode come through that attached to a copper wire (which would be running through a copper pipe and hanging there).

Anyway, to attach pictures just click the button 'Browse...' next to attachement under the box where you write your posts, this will attach one file of a decent size, multiple files can be attached in multiple posts or you can go to the topic 'Forum Matters' where there is a sticky for using the forum hosting services to upload multiple pictures, from there inserting them in the thread involves remembering the URL for them and clicking the button above showing the little picture of a mountain.

Looking foreward to hearing your results.




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[*] posted on 3-6-2005 at 09:57


Wow, tumadre, that's a huge container for sodium production!

Are you sure you need that much insulation? (I suppose you've calculated everthing out, and I haven't, but it does seem like a lot of insulation. Have you looked into using kaowool insulation?)

Hmm, ceramics that resist chlorine. I don't know... what would the reaction be? BromicAcid, why are MgO or Al2O3 containing ceramics best here? (just curious as usual) I did a bit of googling on chlorine and ceramics, but never found anything. Molten NaCl/CaCl2 shouldn't be a problem at 550 deg. C. or whatever it was, but it will leak through porous ceramics noticably. :(

I still want to do this reaction too, as soon as I get a good power supply. I suppose I could rig up something using the MOT I got. As for heating it, sticking the apparatus inside of a furnace would work, although it wouldn't be too precise. But I don't think NaCl electrolysis is as sensitive to changes in temperature as NaOH is.

Molten NaOH in an ALUMINUM container!!! :o:o Now that's confusing, especially because of the thermite reactions that occur with lye and aluminum. (Tacho was doing experiments with these)




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[*] posted on 3-6-2005 at 11:59
"Ceramics uneffected by chlorine"


According to patent US 1,926,072, (quoted elsewhere on this board for the production of PCl3), an alundum (aka alumina) boat was capable of withstanding an atmosphere of chlorine at >900C, even in the presence of finely powdered carbon.

[Edited on 3-6-2005 by Natures Natrium]




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[*] posted on 3-6-2005 at 12:08


The problem with ceramics is they are slightly soluble in NaCl. Al2O3 you may note is soluble to several percent in fluoride melts. Chlorides aren't as good, but it's still there. This would react with the sodium, calcium and electricity to form aluminum and silicon.

Metal is demanded. Stainless might work, Cr2O3 probably isn't very soluble in chloride. A copper vessel could be held in a reducing atmosphere, though it still readily reacts with chlorine gas.

Oh, the phase diagram for Ca-Na shows, at most, 1.5%at solubility. It is dramatically lower at lower temperatures, 99.95% or so should be had on cooling to 100-150°C and filtering the calcium.

Tim
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wink.gif posted on 3-6-2005 at 12:55


I know it’s big but the cost of Nacl is $3.12 for 25 pounds, and I have many uses for sodium, 100 amps will make about 1/5 pound per hour. at 7 cents per kilowatt hour electerical cost, this makes about $1.50 per pound.
The ceramic is the divider; it is necessary to prevent the sodium from touching the carbon anode. I bought 25 pounds of standard white clay from my high school art class ($7), I am going to fire a sample of it and let it sit in a container of liquid Nacl and Cacl2 to see what happens.
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[*] posted on 3-6-2005 at 13:14


How are you going to keep the melt at the proper composition of CaCl<sub>2</sub>/NaCl ? For my plan I was just going to use a mixture rich in NaCl and melting at a higher temperature and keep the electrolysis proceeding until the temperature to keep the mixture molten was exceeded (so the NaCl concentration would decline to the eucetic of the mixture then continue to decline during which time the melting point would rise) so are you going to have some sort of feeding system for adding NaCl during the electrolysis?



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[*] posted on 3-6-2005 at 13:34


If the Cacl2 remains the same, them only by adding too much Nacl will there be a problem. I have not yet constructed the steel and pipes so I have not designed the Nacl insertion apparatus
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[*] posted on 3-6-2005 at 18:05


Ya know, if you keep it just above the eutectic temperature, you have three conditions:
sodium-rich: NaCl crystallizes out
neutral: fully molten solution
sodium-lean: CaCl2 crystallizes out

Since sodium content is constantly falling, the only way the first condition can occur is if you add too much. If you add too much, you get a supersaturated solution of NaCl (in CaCl2) and no more will dissolve. If sodium falls too much, CaCl2 will crystallize on the coolest and/or most sodium-deficient surface, which means either the top or bottom corner of the cathode.

The best way to go about this, I think, would be to make (anhydrous!) salt briquettes. You could try making cakes of salt by evaporating water and drying them (heating slowly to a pretty good temperature, I'd say at least 200°C), or perhaps liquid-phase sintering a mixture of mostly NaCl with small amounts of KCl and CaCl2 which form a tertiary eutectic which fuses it together at say 400-500°C. The fusing would guarantee anhydrous conditions, always useful when adding to a very hot solution.

Then to use, just drop them in as needed, let it suck on the ice cubes (so to speak) and if you add too many, so what, the excess of sodium is localized right there, they'll simply dissolve as needed!

Edit: or yeah... you can just melt and cast salt ingots (mind that they are very fragile from such sudden cooling), if you can reach the additional temperature (red/orange heat).

Tim

[Edited on 4-6-2005 by 12AX7]
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