Duster
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Sodium Synthesis
Alright well I did a search and didn't come up with anything really, one thread did deal with one way but it seemed a bit too complex... Anyway,
my question is this:
How do I make pure or as pure as possible sodium?
I remember a while ago I watched an hour long show on salt (yes salt... an hour... ask me anything I know it all about salt now), and they menchined
NaCl was used to make Na... I was contemplating this and consulted some other chemists and the general idea seemed to be, get something that reacts
with the Cl, filter out whatever you get (if a gas, then nothing) and your left with pure (or near pure) Na... This seems way to easy, but does anyone
know of such a chemical that would be used for this?
I also belive I remember something dealing with electricity... Though I don't remember that very well.
Any thoughts? Thank you again.
Duster
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Blind Angel
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There is already a big thread on sodium isolating.
Look:
Sodium !
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Duster
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I did look at that thread Angel and took notes, but most if not all of that deals with electric means... I may still try that but i curious if there
is a strictly chemical way to procure sodium.
Like what you said Ross lol thank you, ill keep that in mind... Though, if you could get Li, wouldnt you probably be able to get sodium? (both are
same faimly of metals)
Duster
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The Ed
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not nessecarily
if you want to waste an incredible ammount of money you could get lithium out of alkaline batteries. I think calcium is somewhat available. Potassium
would be an incredible waste. Anyone who has potassium doesn't need sodium.
[Edited on 29-9-2003 by The Ed]
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Duster
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Hmm, perhaps I'll have another look at that other thread lol, I am mostly intersted in saving money and staying simple...
Ross, of course, and I actually know a few guys around here with some spare Li... Though I doubt they would just chuck it over to me (worth a shot
though).
Duster
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JustMe
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Per the Textbook on Inorganic Chemistry:
Sodium may be obtained by heating sodium peroxide with carbon, or sodium hydroxide with magnesium. It doesn't give any more details regarding
these two methods.
The Castner process for producing sodium was to heat sodium hydroxide with carbon and iron at 1000 degrees...
6NaOH + 2C = 2Na + 3H2 + Na2CO3.
Sodium is a very active alkali metal, it is not easily liberated. Hence, it is not really cheap or simple. The "cheapest" way is via
electrolysis of molten salt or sodium hydroxide, the cathode should be iron. If salt is used, watch out for the chlorine!
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Blind Angel
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I think that bubling Cl<sub>2</sub> in water to yield HCl would be a good idea. Also, spa store have NaBr which has a lower melting point
tha NaCl and the Br<sub>2</sub> could be retrieved and used in bromination... NaI is also available...
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the_alchemist
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Electrolysis
The Castner process may be the most cost efficient method of production, but it certainly isn't the easiest.
The way to go is electrolysis. Hell, you can do it with an aluminum can anode an iron cathode, and a battery, but it'll probably take a while to
collect the sodium. You'll probably find something on electrolysis of NaCl if you do a search but NaOH would be the way to go if you ask me, no
chlorine gas. Anyhow just electrolyze the solution and collect the metal as it forms. It's a relatively simple procedure and you should be able
to find a process on google or whatever.
[Edited on 2-10-2003 by the_alchemist]
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Marvin
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A rather cunning reaction between magnesium and sodium hydroxide in organic solution has been cited in a patent, though I forget the thread, and due
to what I assume is a downed router, causing merry hell with my access to these forums so I suggest you search for it, its very doable.
Sodium hydroxide and magnesium are stated to react if heated, but the reaction is supposed to be very violent.
Electrolysis in an alumininium can wont work If you use NaCl the can will melt before the salt does, if you use NaOH, the caustic soda will melt
first and then promply eat the alumnium, along with any glass its in contact wth.
Reducing Na compounds with K, or vice versa doesnt produce one or the other, it produces one of 2 eutectic alloys. These tend to be liquid at room
temp, and ignite at room temp in air. Its quite possible similar problems might happen with other combinations of group 1 metals.
Reduction of NaCl with Fe to produce Na gas, like carbon reduction of soda requires silly temperatures.
If you can get magnesium, go with reduction in an organic solvent, catalysed with an ethoxide, its probably the easiest to do.
If not, electrolysis of NaOH, in a steel container, with some sort of air excluding gas present. Cant be CO2 for obvios reasons. This is less
trivial, but with some planning and equipment it can certainly be done. Dont go with anything less than 5Amps, batteries will be useless unless you
use a car battery, and thats not a good idea, follow other instructions in the electrolyic sodium thread, or you'll most likley get nothing.
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a_bab
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"Anyhow just electrolyze the solution and collect the metal as it forms"
Sorry the_alchemist, but the solution of NaOH won't give you any Na. Well, at least not in the pure form. It'll react back with the water to
form NaOH again.
The molten NaOH will disolve the Na as it'll be formed, rendering a silvery stuff.
I'd stick to either castner of to chloride electrolysis if I would need large amounts of sodium. Or I'd buy it from Ebay or some chem supply
house.
[Edited on 4-10-2003 by a_bab]
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BromicAcid
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**Reducing Na compounds with K, or vice versa doesnt produce one or the other, it produces one of 2 eutectic alloys.
These tend to be liquid at room temp, and ignite at room temp in air. Its quite possible similar problems might happen with other combinations of
group 1 metals. **
Im pretty sure that most of the potassium produced industrially is from reduction of potassium chloride with sodium metal. The reaction is driven
foreward because the lower boiling point of the potassium allows for it to be continuously removed in the vapor phase. Average purites resulting from
this method are on the scale of 99%+ pure.
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Marvin
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a_bab, I assumed he meant molten rather than solution in water, its been so long Ive seen anyone make that mistake.
Provided the melting point of the hydroxide is not exceeded by much, electrolysis of molten NaOH works fine (Ive done it). Its much easier and safer
than trying NaCl.
BromicAcid, The eutectics are melting point eutectics, I have no information as to if either exists in the vapour phase.
A mixture of Na and KOH heated together gives NaK2 eutectic if allowed to go to completion according to the encyclopedia of chemical reactions. Its
something I remeber becuase one would think potassium wouldnt be produced at all, and here it is in 66% yeild without being externally driven. The
reason is becuase the energy/entropy of forming the alloy is so much greater than the energy difference between the reduction potentials.
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Wolfram
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When you talk about sodium and other elements or small inorganic substances I would´nt call it "synthesis".
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BromicAcid
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Yeah, decent point, usually considered isolation. Some newsgroups if you say synthesis they'll start talking about bombarding neon with
highspeed hydrogen atoms.
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Alkali
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Hi Marvin:
How much sodium you were able to synthesise??, I mean how many grams??
I mean, is it effective to use NaOH instead of NaCl?
Thank you you.
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Marvin
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Yes it is effective 80% yeild based on current as I said, the only thing that stops it still being used commercially is that the NaOH would have to be
made from NaCl anyway.
Electrolysis of molten NaCl is not feasable as a home method.
My test method didnt produce a lot, but then it wasnt meant to. I used the underside of a steel drinks can, the very shallow 'pan', a small
amount of NaOH heated with an alcohol flame, a 5A (7A surge) 12v power supply, the pan was the anode and I think I used a copper wire cathode. I
tried using a carbon rod cathode from a battery but it ate it, faster when the current was on. No gas blanket and the metal globule constantly had a
glowing yellow aura round it. Periodically it would set fire to a hydrogen bubble, and send sodium metal flying across the room. Quite a spectacular
experiement, but not practical for making sodium in quantity. I think I removed about a dozen globules about the size of 1/4 of a pea during the 2
days I was doing it.
My one attempt to scale the process up hit a few problems, the nice strong steel tin I chose turned out to be rather weak after I cut the top off, and
the molten hydroxide wont yeild sodium until its been electrolytically dehydrated. 5A for a large tin setup is far too low to be useful. Attempting
to dehydrate the hydroxide by feeding it aluminium - well the theory was sound - doesnt seem to help at all. Some sort of gas blanket is essential,
N2 could be used, Ive heard of people using propane, butane could also work. They dont ignite at the temperature the cell is working at.
Speaking of temperature, it needs to be maintained at only a fraction over the melting point, this is best achieved by using a strong enough power
supply that the waste heat produced by the electrolysis is enough to keep the mass molten, by altering the current you can very easily maintain the
temperature. If you need external heating this is much more difficult and you tend to be constantly working with hydroxide thats slowly freezing like
my experiments in the shallow dish were. (The alternative is that the melt is too hot and you get no sodium at all). You only have about a 20C
success range but this is very easy to maintain electrically or while the solution is freezing (because that gives out energy to maintain its
temperature).
If you attempt a scaled up version, be aware that leaving the solid hydroxide in contact with air..... It will absorb water, liquify, absorb CO2 from
the air, climb out of the container by capillary action, go all over your work surface and when you come back after a few weeks will have a thick
layer of sodium carbonate snow over evertying. The hydroxide to carbonate I was expecting, the leaping out and running round took me by surprise.
This tends to be more offputting than the outright reaction-didnt-work problems.
If this description sounds amateur, its becuase it was over 10 years and a halfdecent education in chemistry ago. I still have most of the sodium I
made under oil. Someday I will revisit this and do it all properly.
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Theoretic
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The mistake I see is using Al to dehydrate the NaOH. They react:
6NaOH+2Al=>2NaAlO2+2Na2O+3H2.
Otherwise it made quite absorbing reading  .
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Marvin
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Yes, and the formed sodium oxide will do what to the water in the hydroxide melt. Young was I, yes, born the day before, no. 
In practice, as I said, it did not improve matters much.
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