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Author: Subject: Extracting Sodium from Sodium Chlorate
raistlin
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[*] posted on 9-7-2002 at 06:07
Extracting Sodium from Sodium Chlorate


Does anyone have an idea as to how I can separate sodium from sodium chlorate? My problem is, with what few materials I have to use, I dont know what would work to remove the chlorine....

Raistlin
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[*] posted on 9-7-2002 at 09:03


The only realistic route to preparing sodium metal is electrolysis of a molten sodium compound in an oxygen-free environment with inert electrodes. Sodium hydroxide is recommended for such purposes, due to its relatively low melting point.



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raistlin
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[*] posted on 9-7-2002 at 10:13
OK......


That just brings me to another question, where can I get sodium hydroxide?
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[*] posted on 9-7-2002 at 18:26


Sodium hydroxide is sold at hardware stores - it is sold for cleaning out problematic plumbing. It's often sold as "Red Devil Lye". It's extremely caustic, and will instantly blind you if any gets in your eyes.



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kingspaz
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[*] posted on 23-7-2002 at 06:35


if you can't get sodium hydroxide then would it be possible to use sodium nitrate (lower melting point, 308*C)
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raistlin
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[*] posted on 23-7-2002 at 10:19


Ok, Ill try that.

Thanks,
Raistlin




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Rhadon
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[*] posted on 23-7-2002 at 10:32


Is getting sodium nitrate easier than aquiring sodium hydroxide in the US??
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raistlin
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[*] posted on 23-7-2002 at 10:39


I really cannt answer that, because, as you probably know, I am just begginning to delve into chemistry. lol

Raistlin




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[*] posted on 23-7-2002 at 11:58


NaOH should be much easier to get because it can be used in the household. NaNO3 can't.
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[*] posted on 25-7-2002 at 03:54


Does anybody has information on current/voltage setting to optimise the seperation of the sodium metal ???

I tried it once with my 50 volt power source, just melted the NaOH on an electric fire, and aplied the voltage through 2 carbon electrodes (from batteries), the moment the second electrode contacted the fluid, it started to spark ... Does NaOH conduct well ?
I think there was a bit of water still in it.

I don't know about the current that flowed, but this trafo can give quite a bit (I can pull a spark of allmost a cm with these carbon electrodes)
Also, should I use AC or DC current ?




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[*] posted on 25-7-2002 at 08:06


I don't know much about this, but these are my best guesses for what should be used:

-low voltage, high current (use a transformer)
-non-alternating current




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kingspaz
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[*] posted on 26-7-2002 at 05:27


the reason AC is used is because then sodium will only form on one electrode. wait a second....what forms at the anode?
OH- ---> OH + e-......seems to be short of hydrogen....ahh fuck knows...atleast with NaNO3 the products are gaseous so don't contaminate the molten salt.
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[*] posted on 26-7-2002 at 05:29


sorry it won't let me edit....i meant to say 'the reason AC is NOT used' ;)
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[*] posted on 21-9-2002 at 05:04


Obviously, you have to use direct current to electrolize molten NaOH. I used a car battery charger at 12V and it drew 5A.

I scooped off the sodium and kept it in paraffin (you can use kero)

NaOH melts over a spirit lamp.

I would (delicately) further suggest that anyone who uses an AC source doesn't understand electrolysis enough and should stick to aqueous solutions before trying to discover electrolysis with caustic soda.
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[*] posted on 21-9-2002 at 05:21


But how much Na did you obtained in this way ? Hm, it sounds promising...I never tryed it since I have enough lab grade Na.
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[*] posted on 21-9-2002 at 09:43
Very interesting.


Daryl, I must say, it sounds like you've done a lot of interesting experimentation in your time. Do you ever post on Usenet? Because I remember somebody who seemed to describe making white phosphorus from red in the same way that you have described here...

Can you elaborate on how you produced your sodium? I have talked to people who produced tiny amounts via electrolysis of chlorides or hydroxides, but not enough to be of any use. How did you scoop the sodium up and transfer it to hydrocarbon storage? Were you using an inert gas blanket to protect the metal?

I've considered trying a mixture of some of the alkaline earth metal chlorides/hydroxides along with the alkali metal chlorides/hydroxides because of the greater density of the alkali earth metals. Theoretically, these metals should sink to the bottom of the reaction vessel as they're produced (protected by molten salt) instead of floating to the top and oxidizing.
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[*] posted on 21-9-2002 at 13:56


Admittedly, the yield of sodium is small, but if you've never seen it before, it's a good feeling. No the sodium doesn't sink to the bottom, it floats to the top. You just scoop it up with a nickel spatula.

Hot molten sodium can burn in the air but most of it won't. It's not as reactive as the other alkali metals.

I used a nickel crucible as the anode and a hand held coat hanger wire as the cathode.

You have to watch the temperature because if it's too hot, the sodium burns in the air and if it's too cool, the NaOH starts to solidify in big bubbles and escaping hydrogen starts to collect under the bubbles.

My main problem was to find a suitable power supply. I just remembered... I used a car battery and a battery charger provided the ammeter. Any other mains supply of 12V doesn't have enough current.

I tried the same trick with KOH, but it didn't work. either I needed more amps or potassium dissolves in the moltern KOH. I never did find out. The potassium would have been too difficult to handle.

In my time, I was able to buy sodium and potassium over the counter. A few years ago, I again purchased some sodium over the counter. Therefore the electrolysis method was a curiosity rather than a means to obtain sodium.

Be careful of the fumes of NaOH, it's not very pleasent.
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[*] posted on 21-9-2002 at 14:02


The reason your potassium metal attempt failed could also be because potassium metal is significantly more reactive than sodium metal. :)



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[*] posted on 22-9-2002 at 03:35
mercury?


Anybody ever tried electrolysis with mercury cathodes in solution? The only problem is the solvent here, which can't be water for obvious reasons...



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[*] posted on 22-9-2002 at 04:35
Glad you proposed that.


Vulture, Im glad you brought that up. I had planned to before I saw your post.

Firstly, do not electrolyse NaNO3, think of what is going to be formed: Nitric acid from the water present, and nitrogen oxides as byproducts. Not fun to work with.

Use NaOH. Or use NaHCO3/Na2CO3 in water, it doesn't really matter.

Kingspaz: The equation is 2OH- - 2e ---> H2O + 0.5O2

And on the other electrode, the equation is Na+ + e + H2O(solvent) ---> NaOH + 0.5H2

And with no water present Na+ + e ---> Na(l)

The water produced at the anode boils away instantly as it reaches the surface due to the high heat.

Electrolysis with a mercury metal cathode is tricky because the amalgamated sodium will solidify around 3-6% sodium, so the Hg must be mechanically pumped to prevent this. This is the oldest method for commercial production of sodium metal, used before popularization of the Downs Process.

A tank with a center partition extending nearly, but not completely, to the floor of the vessel, was employed. A mercury pool was introduced into the vessel, so that the partitioning wall was dipped into the Hg but not touching the floor. This created two seperate chambers seperated by the partitioning wall, with a mercury pool serving as the hallway between them.

One chamber was filled with aqueous NaOH or NaHCO3, and the other with nonpolar solvent.

Electrodes are introduced, electrolysis begun, and a mechanical rocking of the tank was started to make the mercury pool slosh from one side to the other, to prevent solidification as it absorbed sodium metal.

The mercury pool serves as the cathode, and electrodes dipped into the water solution served as the anodes.

The mercury was gradually replaced in a continuous process, and the material was distilled to remove the mercury, and the sodium was recovered as the metal in an inert atmosphere. Lot's of work and using lots of mercury, which is not only dangerous, but VERY expensive to purchase on this scale. Hg is sold for around $100 USD per pound, and a pound is only a tiny bottle. I'd estimate that this process usually employed several tanks weighing over 100 pounds each.

Mercury takes forever to distill by the way.

Potassium metal can be prepared the same manner. It does oxidize faster, but you shouldn't be doing this in an open atmosphere anyway, boys. You know that.

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[*] posted on 22-9-2002 at 08:13


Lol, don't try electrolysing molten NaNO3! You know what extremely flammable molten metals and molten oxidisers do when they combine, don't you?
As the cathode, use a piece of wire inside and upside-down porcelain crucible underneath the electrolyte. This will collect your metal and keep it away from the air (ensure that as much air as possible has been removed from underneath the crucible to minimise losses).
12V is sufficient. Any current can be used, but it's obviously quicker if you use a higher current. A car battery or two will work fine.

I know that rubidium is soluble in molten rubidium hydroxide (or was it chloride?), so maybe potassium does dissolve in its molten hydroxide?
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[*] posted on 22-9-2002 at 08:30


What I meant by "potassium is too reactive" is that the potassium metal likely is being oxidized on formation. Just look at the difference in violence of the reaction between sodium metal when thrown in water, and potassium metal when thrown in water. Another very likely possibility:

2K + 2KOH ----> 2K2O + H2




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[*] posted on 22-9-2002 at 10:13
That Reaction is incorrect


That last reaction does not happen.
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[*] posted on 22-9-2002 at 11:39
Mercury pricing


Quote:
The mercury was gradually replaced in a continuous process, and the material was distilled to remove the mercury, and the sodium was recovered as the metal in an inert atmosphere. Lot's of work and using lots of mercury, which is not only dangerous, but VERY expensive to purchase on this scale. Hg is sold for around $100 USD per pound, and a pound is only a tiny bottle. I'd estimate that this process usually employed several tanks weighing over 100 pounds each.


Mercury metal might be $100 a pound for small quantities of reagent grade material, but I'm sure that's not what was used in these sodium production tanks! I can buy technical grade mercury for about $16 a pound on a small scale. As a commodity, mercury is about $2 per pound. I don't know what historical pricing was like, though.
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[*] posted on 22-9-2002 at 13:27


since hydrogen has a very high overvoltage at a mercury cathode, sodium metall is formed instead of hydrogen when NaCl is electrolysed in aqeous solution. the sodium formes a amalgame with the mercury. No news here...

however, the difference in overvoltage and the electrode potential for Na+ => Na gets smaller and smaller when the sodium content is increasing in they mercury, and at about 0.3 % m/m they are equal , meaning that the sodium starts to react with the water.
thus making sodium this way means processing lots of mercury

/rickard
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