Sciencemadness Discussion Board - Sodium Metal from NaCl Reduction using Lead

Sodium Metal from NaCl Reduction using Lead

ShadowWarrior4444 - 18-5-2008 at 00:51

Fairly explicit title. I intend to use the process outlined in British Patent 23, 689; Vournasos, A.C. 1908 which involves heating lead metal and NaCl to red heat, at which point the sodium distills over. Lead should be left as PbCl2. What would be the best method for recovering the lead metal: Electrolysis with lead anode and graphite cathode? (Perhaps I can use lead/tin solder for the anode, thereby allowing the recovery process to be a refining process as well.) The only problem might be the solubility of Lead Chloride. Are they any other practical methods for reducing the chloride to the metal?

Another process I may attempt is the reduction of Sodium Tetraborate or sodium bicarbonate with carbon. The drawbacks for this is that copious amounts of CO are produced. Perhaps I'll conduct it into a chamber with scrap nickel to make some "Death Juice."

Yet another considered is carbon reduction of NaCl using CaO. This would likely be the most inconvenient as it produces just enough CO to be annoying, and would require some method to recover CaO. I would likely rather use a Downs Cell, which I will probably end up doing anyway. (Now if I can get around to building a stirling cryocooler to condense that Cl2.) Or perhaps the gas liquefier posted previously on these forums [somewhere,] though I'm hesitant to put Cl2 though a compressor.

Well then, enough rambling! Discussion: Production of Sodium Metal using non electrochemical methods. (Preferably via the Lead reduction/recovery process, as the cycle is near fully regenerative.)

[Edited on 5-18-2008 by ShadowWarrior4444]

panziandi - 18-5-2008 at 07:20

You can electrolyse molten lead (II) bromide mp 367°C to recover lead (molten) and bromine (vapourises off obviously) - this was a demonstration of non-aqueous elcetrolysis in school chemistry class! Perhaps lead (II) chloride mp 501°C could be easily melted with a burner and electrolysed to recover your lead? Carbon anodes would be used but chlorine at that temperature would likely eat the carbon quickly so replacement would be needed.

There is a post on unconventional routes to sodium somewhere in this forum already but I hope the above will help you.

Good luck!

p.s. obvious but distilation of sodium must be done in inert atmosphere such as argon.

Aqua_Fortis_100% - 18-5-2008 at 08:02

ShadowWarrior4444, check out excellent papers scanned by BromicAcid in his page: http://members.aol.com/bromicacid/sodium/ or go to Tacho's page for an related experiment: http://www.tacho.kit.net/pag5.htm

Also like panziandi said, go to search and find out this thread:http://www.sciencemadness.org/talk/viewthread.php?tid=2105&page=1 or search other thread..IIRC is another very extensive thread on sodium making, as well an excellent work posted on prepublications (by len1)..

Also, IIRC there are numerous other routes like lead + NaCl , some of them react to give sodium at even lower temps and all are , at least, briefly described on Bromic page above.. have fun!

By the way, IMHO will be somewhat hard to recover your Pb after converting it to PbCl2.. But will you really need that? lead is usually very cheap and even free from some sources..
Besides you can convert the lead chloride in other lead compounds more reactive (e.g. lead carbonate) that you can have just lying around to make other Pb compounds (e.g. lead nitrate).

[Edited on 18-5-2008 by Aqua_Fortis_100%]

MagicJigPipe - 18-5-2008 at 08:36

I agree that it would definitely not be worth it to recycle the lead. I would also use it to make other lead compounds or just keep it as is! I think you have a pretty massive undertaking here. More so than you think. Especially the requirement of an inert atmosphere.

[Edited on 5-18-2008 by MagicJigPipe]

ShadowWarrior4444 - 18-5-2008 at 09:05

I would first like to state that I have indeed read both Bromic's site, and the unconventional sodium threads. I have also searched the entire forum using Google, and the find feature on the printable version of nearly every sodium thread.

As for the inert atmosphere requirement, it is quite a simple undertaking should I construct a full-scale reaction chamber; though, I have been considering that it may be possible to produce sodium metal in its molten form simply by adding NaCl to molten lead. The molten sodium may be decanted off the PbCl2, as I doubt it is soluble in molten sodium.

Quote:

Besides you can convert the lead chloride in other lead compounds more reactive (e.g. lead carbonate) that you can have just lying around to make other Pb compounds (e.g. lead nitrate).

Would anyone happen to know of a way to convert PbCl2 to other lead compounds that does not use an alkali nitrate?

panziandi - 18-5-2008 at 09:29

Hey,

Sodium is soluble in lead. In fact an alloy of sodium and lead can be used as a drying agent for some solvents instead of using neat sodium! Its a bit like the amalgam.

I expect the red heat is vital in that lead is a poor reducer and the whole reaction would be driven by the fact the sodium is distilled out.

One thing I am looking into (not for sodium mind as I have an plentiful stock!) is the electrolysis of certain metal salts in organic solvents. I had to put it on hold for the time being but will get back to it soon and will post my results on my website and of course on this forum!

12AX7 - 18-5-2008 at 14:21

I find that hard to believe. Pb is quite unreactive, and its chloride quite volatile. Nascent Na would surely reduce it to Pb.

SW4444: CO burns readily in air.

Tim

woelen - 19-5-2008 at 02:07

I have the same feeling as 12AX7. I would even expect it the other way around. If you mix PbCl2 and Na-metal, then I expect formation of NaCl and Pb. If making Na from lead-metal and a sodium salt does work, then I am very interested in an explanation for this.

Formatik - 19-5-2008 at 02:42

Quote:

Originally posted by ShadowWarrior4444 Another process I may attempt is the reduction of Sodium Tetraborate or sodium bicarbonate with carbon. The drawbacks for this is that copious amounts of CO are produced. Perhaps I'll conduct it into a chamber with scrap nickel to make some "Death Juice."

Heating bicarbonate or carbonate of potassium or sodium with carbon is risky, because it can form an explosive by the combination of e.g. hot potassium (or analogous sodium) metal with carbon monoxide to give the black potassium rhodizonate (K2C6O6), this compound explodes spontaneously when dry and reacts violently with water. I've thought about roasting carbonates and carbon in pipes, but it sounds too risky.

[Edited on 19-5-2008 by Schockwave]

ShadowWarrior4444 - 19-5-2008 at 11:38

Quote:

Originally posted by woelen
I have the same feeling as 12AX7. I would even expect it the other way around. If you mix PbCl2 and Na-metal, then I expect formation of NaCl and Pb. If making Na from lead-metal and a sodium salt does work, then I am very interested in an explanation for this.

It may be that the reaction is driven by the evaporation of sodium as previously mentioned. I also seem to recall that PbCl2 holds on to its chlorine quite tightly.
Salient reference: http://members.aol.com/bromicacid/sodium/4.jpg

Quote:

Originally posted by Schockwave
Heating bicarbonate or carbonate of potassium or sodium with carbon is risky, because it can form an explosive by the combination of e.g. hot potassium (or analogous sodium) metal with carbon monoxide to give the black potassium rhodizonate (K2C6O6), this compound explodes spontaneously when dry and reacts violently with water. I've thought about roasting carbonates and carbon in pipes, but it sounds too risky.

I am immortal, infallible, and fearless. *smirk*

Also:
http://www.sciencelab.com/xMSDS-Rhodizonic_acid_dipotassium_...
http://msds.chem.ox.ac.uk/RH/rhodizonic_acid_dipotassium_sal...
In addition, Rhodizonic acid is used in Gun Shot Residue tests to detect lead. Rhodizonates should also decompose at around 300C, so a reaction carried out above that temperature would not allow them to form.

As for the viability of the lead reaction, it seems that Lead Chloride boils at 950C whereas Sodium boils at 881.4C; were the reaction to be conducted between those temperatures, sodium should distill smoothly.

[Edited on 5-19-2008 by ShadowWarrior4444]

Formatik - 19-5-2008 at 11:50

Quote:

Originally posted by ShadowWarrior4444 I am immortal, infallible, and fearless. *smirk*

Also:
http://www.sciencelab.com/xMSDS-Rhodizonic_acid_dipotassium_...
http://msds.chem.ox.ac.uk/RH/rhodizonic_acid_dipotassium_sal...
In addition, Rhodizonic acid is used in Gun Shot Residue tests to detect lead. Rhodizonates should also decompose at around 300C, so a reaction carried out above that temperature would not allow them to form.

The MSDS aren't very useful, you have to look here, here, or here. Phew. Explosions have also result from trying to get potassium this way, you can read about this here.

[Edited on 19-5-2008 by Schockwave]

panziandi - 19-5-2008 at 13:14

potassium carbonyl forms if the potassium vapour condenses in the presence of carbon monoxide. potssium carbonyl is given the structure of the hexa-potassium salt of hexahydroxybenzene. di-potassium rhizonate is a different compound.

ShadowWarrior4444 - 19-5-2008 at 13:55

Quote:

Originally posted by Schockwave
The MSDS aren't very useful, you have to look here, here, or here. Phew. Explosions have also result from trying to get potassium this way, you can read about this here, they call it the potassium salt of hexaoxybenzene.

[Edited on 19-5-2008 by Schockwave]

I should note, that were I trying to obtain potassium by reduction of its carbonate, I would be using elemental silicon, not carbon. That said, these reports of explosions are quite sketchy with some being attributed to the formation of potassium peroxide and its interaction with elemental carbon. I don't anticipate any such unsavory occurrences while producing sodium from the carbonate method.

Ancillary reference: PROCESS FOR THE PREPARATION OF HEXAHYDROXYBENZENE

[Edited on 5-19-2008 by ShadowWarrior4444]

Formatik - 19-5-2008 at 14:51

This will clear it up:

Carbonyl potassium is KCO, it is different than hexahydroxybenzene potassium K6C6O6. Both of which are actually different than potassium rhodizonate (K2C6O6). Then there is also potassium acetylenediolate (K.OC.:CO.K).

Hexahydroxybenzene potassium is made as the Lehrbuch der organischen Chemie by Victor Meyer describes by leading dry CO into molten potassium, it is a grey mass and becomes highly explosive upon standing in air (they say it's harmless freshly prepared), this comprises the most significant part of the black mass which had resulted in large explosions by the reduction to potassium, they say its formation can be completley avoided nowadays, but don't give the details afterwards. Though at higher temperatures (above 180ºC) the reaction between CO and potassium mainly the hexahydroxybenzene compound forms, below that near the m.p. of potassium around 62.3ºC acetylenediolate and an "organometallic compound" predominate. According to Meyer, potassium rhodizonate forms from the hexahydroxybenzene by washing it continuously with dilute alcohol where it results by oxidation. Carbonyl potassium is described by Gmelin and is made by leading CO at -50ºC into a solution of K in NH3, until it turns from blue to white-pink. It detonates at 100ºC, upon exposure to air or a drop of water it forms K2CO3, potassium oxide, and carbon. Sorry about the confusion, I didn't know so many different compounds form from CO and potassium.

The_Davster - 19-5-2008 at 15:12

Very interesting info Schockwave.

As for why this works, all reduction-oxidation potentials are in equilibrium, and while the electrical potentials do indicate that nothing should happen, as the equilibrium constant that can be calculated from the electrical potentials hugely favours lead and salt, keep in mind that since it is in equilibrium, a small ammount of sodium will be present, which can be continously distilled off shifting the equilibrium to obtain more Na.

12AX7 - 19-5-2008 at 19:55

I still don't buy it. The boiling point of PbCl2 is very close to sodium's. If the reaction proceeds at all, they would be co-present in the metallic (Pb, Na), salt (NaCl, PbCl2) and gaseous (PbCl2, Na) phases!

I would find Pb + NaOH more believable, as PbO has a somewhat higher melting and boiling point, but still nothing impressive, not like Al or Mg reduction.

Besides, if this reaction proceeds, why not with somewhat more reactive and much cheaper Fe? FeCl2 would be the product, which has a sufficiently high boiling point. FeO, Fe3O4 and Fe2O3 all have much higher melting and boiling points, yet there is no reaction between NaCl or NaOH and Fe.

Tim

ShadowWarrior4444 - 19-5-2008 at 20:34

Quote:

Originally posted by 12AX7
I still don't buy it. The boiling point of PbCl2 is very close to sodium's. If the reaction proceeds at all, they would be co-present in the metallic (Pb, Na), salt (NaCl, PbCl2) and gaseous (PbCl2, Na) phases!

I would find Pb + NaOH more believable, as PbO has a somewhat higher melting and boiling point, but still nothing impressive, not like Al or Mg reduction.

Besides, if this reaction proceeds, why not with somewhat more reactive and much cheaper Fe? FeCl2 would be the product, which has a sufficiently high boiling point. FeO, Fe3O4 and Fe2O3 all have much higher melting and boiling points, yet there is no reaction between NaCl or NaOH and Fe.

Tim

There is a process that involves heating Iron and Sodium Fluoride under vacuum to obtain sodium; apparently it does not require either to melt. It cites the high volatility of iron halides as the reason it cannot be used with NaCl. It is also mentioned that iron displaces sodium from many other sodium compounds when heated under vacuum.

Hackspill, L. and Grandadam, R., Compt. Rend., 180, 68-70 (1925).

Soc. D'Electro Chimie, D'Electro-Metallurgie et des Acieres Electriques D'Ugine, French Patent 603,825 (1924).

Co-presence in the metallic phase to be expected, sodium is soluble in molten lead, and this solubility is exploited in many sodium production systems to remove the sodium; it can be easily distilled out of the lead. I also suspect that this is one of the principles behind the NaCl reduction--some sodium is formed, is absorbed into the lead, and is distilled out. It indicates that the lead is heated until just below red heat, at which point sodium distills over.

I suppose the easiest way to settle this is attempting it on a small scale in a crucible. One doesn’t need to collect the sodium, simply watch to the sodium oxide fumes.

[Edited on 5-19-2008 by ShadowWarrior4444]

panziandi - 19-5-2008 at 23:14

Perhaps a more practical version of this Pb/NaCl would be to electrolyse a molten sodium salt with a molten lead cathode? Sodium would dissolve in the molten lead forming an alloy, the sodium can then be distilled out of the alloy once you have accumulated enough? A eutectic of sodium salts being used to allow operation at a "reasonable temperature"??? The lead cathode would be under the molten electrolyte and so the sodium wouldn't be attacked by the atmospher or the by products of the reaction. Later the sodium could be distilled (perhaps a short path all-in-one setup under argon, the condensed liquid run directly under paraffin liquid for solidification and storage?

12AX7 - 19-5-2008 at 23:33

I suppose the advantage to that over, say, a Downs cell is, the sodium stays at the bottom, with lead? Might be a little easier.

Tim

S.C. Wack - 20-5-2008 at 01:25

No one can bother to actually read the patent before commenting unless there is a link, maybe even if there is.
http://v3.espacenet.com/origdoc?DB=EPODOC&IDX=GB19082368...

panziandi - 20-5-2008 at 05:54

Quote:

Originally posted by 12AX7
I suppose the advantage to that over, say, a Downs cell is, the sodium stays at the bottom, with lead? Might be a little easier.

Tim

Yes quite the sodium dissolves in the lead and stay at the bottom of the melt. You wouldn't use NaOH as it reacts with lead, but a eutectic of halides perhaps would work, IIRC a mix of sodium and calcium chlorides has a lower mp.

Quote:

Originally posted by S.C. Wack
No one can bother to actually read the patent before commenting unless there is a link, maybe even if there is.
http://v3.espacenet.com/origdoc?DB=EPODOC&IDX=GB19082368...

I didn't deny the validity of the patent. It may occur or it may not (like many patented experiments), patents never reveal the details or are peer reviewed. However I merely suggested a modification.

SW4444:

I would try this reaction firstly in a cruicible like they state and see if the fully cooled melt does indeed show alkaline reaction with water. If it does use a distillation method 2nd time around and try to isolate sodium. But do consider the electrolysis using a molted lead cathode, since I think that is a promising variation of the patent.

Picric-A - 25-7-2008 at 14:38

what temps do you need to heat the mixes ( of say, NaOH/Pb ) to?
surely it would be difficult to find a vessel able to take this heat and then find an attachment to connect a condenser?

ShadowWarrior4444 - 25-7-2008 at 15:08

Quote:

Originally posted by Picric-A
what temps do you need to heat the mixes ( of say, NaOH/Pb ) to?
surely it would be difficult to find a vessel able to take this heat and then find an attachment to connect a condenser?

The boiling point of sodium is 883C, this is *easily* attainable in a flower pot. Good quality ceramic crucibles are generally rated up to 1400C.

Incidentally, I do not believe you can use NaOH for a reduction with another metal; I seem to recall that it must be a non-volatile halide.

As for a condenser, a nice stainless steel pipe would be very useful, even fused quartz or perhaps with some creativity a borosilicate chamber could be used. The main difficulty in condensing sodium is the inert atmosphere. (Regulating pressure and the like.)

tumadre - 25-7-2008 at 15:53

Salt will saturate a flower pot and corrode electrical heating elements, so make sure its physically separated or use gas.

Picric-A - 26-7-2008 at 14:43

ok this is a pretty crap drawing but it pretty much shows what i am thinking.
The condenser on top be attached to the cruicible somehow to make an airtight seal and will have water flowin through it fast to cool down the distilled sodium.
The cruicible will have two 'inlet and outlet' taps either side to flush out with inert gas (eg. helium which can be bought form party supply stores) but then sealed when the distilling takes place so the sodium vapour doesnt get carried away.

soidum prep.jpg - 22kB

ShadowWarrior4444 - 26-7-2008 at 17:36

You must not seal the chamber completely when heating; the increased pressure may cause a breach in addition to reducing the amount of sodium that will boil off. Also, in this particular case the sodium would likely condense and stay as a liquid due to its close proximity to the heat source--this may cause it to drip back down. (Upon contacting the PbCl, it will revert.)

You will most likely need a separate chamber for sodium condensation, and a useful way to usher the sodium into this chamber could be putting the inert gas outlet in it. I also seem to recall designs that include a trough of sorts for allowing the sodium to drip into a collection chamber.

[Edited on 7-26-2008 by ShadowWarrior4444]

Picric-A - 27-7-2008 at 02:22

well i didnt show it well in te diagrm but the cruicible would be about 30cm high but yes the pressure would be a sight problem...
maybe if you first flushed the vessel out with an inert gas, then you pulled a slight vaccum on it?
or if that fails, non-return valves?
You would have to keep it airtight or else oxygen would diffuse in and it would result most likely with an explsoion

ShadowWarrior4444 - 27-7-2008 at 13:05

Quote:

Originally posted by Picric-A
well i didnt show it well in te diagrm but the cruicible would be about 30cm high but yes the pressure would be a sight problem...
maybe if you first flushed the vessel out with an inert gas, then you pulled a slight vaccum on it?
or if that fails, non-return valves?
You would have to keep it airtight or else oxygen would diffuse in and it would result most likely with an explsoion

Sodium will not explode in an oxygen environment, only burn.

The somewhat 'golden' rule of heating things up is making sure they are *not* pressure sealed. You would much rather have some of the sodium burn than the entire vessel breaching to spray sodium and lead everywhere.

Oxygen will not diffuse in if you have a slow current if argon, it is much 'heavier' than air. A one way valve on the exit may be useful to conserve argon, but a vacuum will be created when the chamber cools and so must be repressurized with argon before opening.

Ancillary: A vacuum distillation of sodium may be useful, but thus far I have not heard of it being used as an industrial practice; there is perhaps good reason for this. As in other vacuum distillations, you’ll need to rigorously check for leaks and have argon standing by for repressurization.

[Edited on 7-27-2008 by ShadowWarrior4444]

Tin man - 12-9-2016 at 14:24

I would try heating NaCl with molten aluminium intead of lead, because aluminium has a higher boiling point and is way less toxic. I would also think Al would be a better reducing agent.

blogfast25 - 12-9-2016 at 16:04

Quote: Originally posted by Tin man

I would try heating NaCl with molten aluminium intead of lead, because aluminium has a higher boiling point and is way less toxic. I would also think Al would be a better reducing agent.

That couldn't work because AlCl₃ is highly volatile.

In the past sodium metal has been used to reduce AlCl₃ though:

http://www.azom.com/article.aspx?ArticleID=1530

[Edited on 13-9-2016 by blogfast25]

Cryolite. - 12-9-2016 at 20:13

That is true, but I wonder if sodium hydroxide and aluminum would work. Sort of like a magnesium thermite reaction, but more controlled and with aluminum instead. The aluminum oxide formed is pretty much nonvolatile, so that's a plus.

The issue is that a lot of anhydrous metal chlorides are pretty volatile, due to them being covalent compounds.

Tin man - 12-9-2016 at 20:14

Oh, I should have thought about that. What about zinc powder and NaOH. Just cause zinc powder can easily be made by electrolysis of sodium zincate solutium, adding to the OTCness.

[Edited on 13-9-2016 by Tin man]

elementcollector1 - 12-9-2016 at 20:15

Quote: Originally posted by Cryolite.

That is true, but I wonder if sodium hydroxide and aluminum would work. Sort of like a magnesium thermite reaction, but more controlled and with aluminum instead. The aluminum oxide formed is pretty much nonvolatile, so that's a plus.

The issue is that a lot of anhydrous metal chlorides are pretty volatile, due to them being covalent compounds.

Technically it does, but separation's a chore.
https://www.youtube.com/watch?v=seSg_GWj1b0

Tin man - 12-9-2016 at 20:24

Would getting to 883°C( sodium's BP) be possible with a propane torch?

[Edited on 13-9-2016 by Tin man]

blogfast25 - 12-9-2016 at 23:11

Quote: Originally posted by Tin man

Would getting to 883°C( sodium's BP) be possible with a propane torch?

Not really, no. Vacuum will depress the BP though.

Both K and Na have been obtained by reactive distillation of suitable Na/K compounds with reducing agents like C or Fe (and others), at high temperatures. None of these methods are particularly suited to hobbyists, though...

[Edited on 13-9-2016 by blogfast25]

Tin man - 13-9-2016 at 06:39

Could one use a charcoal furnace designed to melt aluminium? Aluminium's melting point is 660°C.

blogfast25 - 13-9-2016 at 07:01

Quote: Originally posted by Tin man

Could one use a charcoal furnace designed to melt aluminium? Aluminium's melting point is 660°C.

Yes.

Dan Vizine - 21-9-2016 at 06:27

On eBay you can purchase SS 304 cylinders similar to the familiar "lecture bottle" and smaller. Prices can be as low as $40.
Reactor (w/o leaks) issue solved.

You can purchase nickel flakes in graphite as sealant. Leak issue solved.

Any form of compressed air feeding a charcoal fire will give you copious amounts of heat. Hard coal is even better. 40-odd years ago unintentionally melted a 1/2" thick solid Nichrome crucible that way. A vacuum cleaner supplied air. It was the foundry in a 5 gallon pail once described in Popular Science ( back when it was actually a useful magazine).

Using He destined for party balloons is not too good an option. Some has oxygen in it and none of it is pure. Bloxygen (argon) may be worth looking into, I have only used commercial tanks. I have no idea exactly how pure it is.

A vacuum is very desirable. Vacuum pumps are $$$, but eBay provided me with a Hitachi "Cute Vac" pump for $75 (two stage, oil-filled, vane-type). The main shaft seal leaks and someday I'll replace it, but it takes days of continuous running to lose a liter of oil ($15). A vacuum pump is one of the most vital pieces of equipment in the lab. Unless a pump has been run with low oil and been burned and scored inside (only a fool would try to sell that mess), most pumps can be rebuilt with new vanes, springs, seals and filters with re-build kits to give near-new performance.

When condensing, beware of a plug forming in the condenser. SS is a good choice. Transition to glass with a rubber stopper. You'll want a hot air gun handy when and if you distill Na out of the reactor, you never know where or if a plug will occur until you have run several reactions.

As blogfast25 noted, these reductions with less-active reductants require more demanding conditions. Unless this is meant as absolutely pure science, just to see it work, this is a tough road to travel.

Incidentally, K is much safer to work with from certain standpoints. K can be heated to 300 C in air without burning, Na burns spontaneously at half that temperature. Keeping an open condenser (free of plugs) is easier, too.

aga - 21-9-2016 at 06:53

Here, on one occasion, charcoal & a hairdyer on the end of a 1m steel pole created enough heat to melt copper.

~1085 C

Tin man - 22-9-2016 at 10:18

Sorry if this is a stupid question, but would potassium require a higher or lower temperature to form? assuming the same reducing agent was being used.

blogfast25 - 23-9-2016 at 08:10

Quote: Originally posted by Tin man

Sorry if this is a stupid question, but would potassium require a higher or lower temperature to form? assuming the same reducing agent was being used.

Analogous Na and K compounds tend to have very similar Enthalpies of Formation. For that reason it's to be expected that using the same reducing agent will require similar temperatures.

But since as these metals are then obtained by distillation and K has a higher BP, K may require higher temperatures.

[Edited on 23-9-2016 by blogfast25]

Meltonium - 27-9-2016 at 17:10

Instead of using argon as an inert atmosphere, could you use a CO2 generator, or would that react?

Also, you could use a vacuum aspirator to create the low pressure; however, I'm not sure if the presence of the water may be risky business.

Metacelsus - 27-9-2016 at 18:46

Carbon dioxide will react with alkali metals at these temperatures. (Have you ever seen the demonstration where burning magnesium is put into a block of dry ice? It would be like that.)

m1tanker78 - 28-9-2016 at 17:02

Quote: Originally posted by Meltonium

Instead of using argon as an inert atmosphere, could you use a CO2 generator, or would that react?

Liquid sodium reacts with dry CO₂ to form varying degrees of Na₂CO₃, C, and CO. The reaction initiates between 250C and 300C^[1]. Depending on the temperature, a glassy intercalation of the solid products + sodium can be obtained. The reaction produces enough CO that it can be 'flared' off in air and produces a beautiful tenuous light blue flame.

Short answer: Stick to argon.

[1] Miyahara, S. (2011). "Experimental Investigation of Reaction Behavior Between Carbon Dioxide and Liquid Sodium." Nuclear Engineering and Design #241.

Tin man - 28-9-2016 at 21:03

Nitrogen doesn't react with sodium to my knowledge. It would be a lot cheaper than argon.