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Author: Subject: unconventional sodium
Hermes_Trismegistus
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[*] posted on 1-12-2003 at 14:54
Well Gosh darn it all to heck!


If it has worked for you in practice, I would be a fool, (not an uncommon state for me to be in) to say that it is not possible.

If Organikum has a digital camera available, I sure we would all love to see his in action.

Unfortunately, up here in Canada, Na is a scheduled chemical so all my posts regarding a Down's cell have been theoretical (of course!)

Looking forward too! ;)


However I still beleive that the more traditional Downs cell to hold alot of promise, not least of which is the evolution of Cl gas.

I see the immediate and best use of the Cl gas to be to bubble it through a large volume of brine solution made up of Iodized salt.

2I- + Cl2 => 2Cl- + I2

The NaCl left over would be extremely pure and suitable for use in the down's cell and a VERY nice byproduct would be the evolution of Iodine.

:D


Orrrrrrrr.....

Let's say we stick with NaOH as a source of sodium and we still wanted to use NaCl as a source AND isolate Iodine in the process......(greedy aren't I :D)

We COULD electolyze strong brine to evolve the Cl gas, run it through more brine to collect our Iodine....

then use the NaOH produced in the above reaction to liberate poor unfortunate Na from the grips of the evil duo OH-



:D:cool::D

It is true, that in both reactions, a great deal of HCl would be available :P


Now I know what you're thinking.....All that HCl and WHAT TO DO?!?:(

I was thinking since we're playing with on the far left of the table anyway.......:P

why not drip the HCl onto some CaCO3
CaCO3 + 2HCl => CaCl2 + H2O + CO2

then add a little of the precious sodium:P

CaCl2 + 2Na => Ca + 2NaCl


and PRESTO!!:P

Calcium metal and ultrapure non-iodized salt (useful in some obscure reactions)


:P:P



[Edited on 2-12-2003 by Hermes_Trismegistus]




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[*] posted on 2-12-2003 at 11:54
IODINE ??????


you might inform yourself how much iodine is in iodinized salt. Tip: less, much much less.

And Na is not allowed for privates here where I live also. So you will have to miss pictures on this forever because I wouldnt make something fOrBiDDeN never. :P

Tintetrachloride is not forbidden here and doesnt belong in this thread but my distorted mind NEEDS to get offtopic at least a little bit.

The NaOH process is older than the NaCl process btw. For being much easier. Chloride is conveniantly made by bleaching powder and muriatic acid, washed with water and dried with H2SO4 for further abuse.
The Cl2 from a Downs cell is mucho shitty contaminated and completely USELESS for utmost everything my dear Hermes.

But you should try it by any way as I believe you will learn it only by the hard way.

Teaser you. ;)

And non-iodinized ultrapure salt is available for "dishwasher regeneration" cheap and without any hassle.

btw. if you add your precious sodium to some water you will get INSTANT NaOH! And vulture has an receipt for turning GOLD to LEAD! ;););)

[Edited on 2-12-2003 by Organikum]




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[*] posted on 5-12-2003 at 08:41
Sodium thermite


I mixed aluminum powder with solid NaOH and heated it in an iron crucible with a blowtorch until the NaOH melted. Inicially small sodium-orange sparks showed here and there. Eventually a strong reaction happened at once, with orange fire and white smoke. I believe most of the sodium evaporated and oxidized to white sodium oxide fumes, but in the crucible there was a gray concrete-like leftover. It showed strong effervescence when thrown in water, but I can't say it's not just fine aluminum and excess NaOH reacting. It absorbs humidity from air and becomes a bubbling gunk.

I find the sodium salt in nitrobenzene electrolysis a very interesting idea. Why would it be so dangerous? Any ideas? Would it be the risk of fire/explosion in a poisonous media or some cancerigenous mutagenic byproduct? The first option I can manage, but not the second.
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[*] posted on 5-12-2003 at 08:56


I've read detailed accounts of experiments done in nitrobenzene solution electrolysis of alkali metals. Short summary: the cathodic deposits are impure, not terribly high yield, and explosive.

On the other hand, electrolysis of lithium chloride in pyridine gives clean lithium metal rapidly and easily, but it's not so easy for me to get or work with pyridine. Amyl alcohol and acetone are also supposed to work for the LiCl electrolysis, though not as well.
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[*] posted on 5-12-2003 at 14:31


Last I heard, lithium reacted with alcohols (and I'm not sure about pyridine of acetone too).
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[*] posted on 5-12-2003 at 17:00


It does react with alcohols, but the reaction with amyl alcohol isn't especially fast. In any case, you're continually working "uphill." If enough current is delivered, metal will form faster than the alcohol can dissolve it. But high current density seems to lead to malformed and impure metal deposits. This is why pyridine is especially nice, because it can be used with iron wire as cathode and relatively low current density/voltage.
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[*] posted on 5-12-2003 at 17:24


Tacho, that's very interesting.
What is the reaction taht is *supposed* to happen when you react NaOH with Al? (or other salts of Na, is that also possible?)
Else, how about doing this in a large container that is filled with an inert gas such as Argon or helium, to protect Na? Problem obviously seems to be be that this are Na vapors, and they have to condense... to collect them!

Polverone, on another note... this reaction of Li in pyridine, is it, in a modified form, trnasferable to other alkali metal salts such as NaCl/KCl?




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[*] posted on 5-12-2003 at 19:17


Slightly off topic but still on lithium "Lithium has also been isolated by electrolysis of somlutions of lithium salts in non-aqueous solvents. A solution of lithium chloride in anhydrous pyridine give satesfactory results with a cathode current desity of 0.002 - 0.003 ampere per sq. cm. at 14 volts (Kahlenberg, J. Physical Chem. 1899, 3, 603). Solutions of the chloride in alcohols have also been used, but are less satisfactory because the lithium reacts slowly with the alcohol (Patten and Mott, ibid. 1904, 8, 170)"

Darn, it's a photocopy, but I didn't copy down from where so no citation. I wonder if lithium reacts slowly enough with ethanol to facilitate electrolysis of lithum chloride in that medium, it's solubility is 24.28 g/100g solvent at 20C, I know it wouldn't be possible to get sodium in this way though seeing as how it is the alkali metal of choice to produce alkoxides. If you could make lithium though, you could get sodium, you know:

Li(l) + NaCl ----> LiCl + Na(g)

The equilibrium would be to the right with the sodium boiling at 892 C compared with lithium at somewhere around 1317 C (lithium vapor is red, beautiful!). But then again that would be impossible both due to the fact borosilicate glass is not a good idea for that temp, plus when liquid lithium touches glass, watch out!

"Liquid lithium is the most corrosive material known. For example, if a sample of lithium is melted in a glass container, it reacts spontaneously with the glass to produce a hole in the container, the reaction being accompanied by the emission of an intense, greenish white light." (Descriptive Inorganic Chemistry 3rd edition, Geoff Rayner-Canham; Tina Overton)

BTW: For the link above relating to the molten salt database I didn't mean for people just to look at my examples, but to actually type in NaCl or whatever sodium salt you have availible and look at all the eutectic mixtures, 652 of them if you type in NaCl!




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[*] posted on 5-12-2003 at 20:01


Ethanol reacts too fast for it to work for lithium electrolysis. At one point xoo1246 posted a fascinating patent about producing alkali metals by heating their hydroxides with magnesium in a high-BP hydrocarbon. Unfortunately, he appears to have erased all his posts! If anybody has the patent noted I'd appreciate it if you could post it here.
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[*] posted on 6-12-2003 at 09:21


Chemoleo,
I didn't think much about the reaction.
I tried Al powder based on the refs. that said that sodium was produced reacting iron powder with NaOH. Al is more reactive then Fe in this case. If they had cheap Al in those days, they would probably use it.

Inert gas can be provided by a plastic bag with steel wool that has been soaked in dilute acid. If left overnight, only nitrogen is left. I then squeeze it into the "chamber".

I plan to do it soon.
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[*] posted on 6-12-2003 at 14:31
Sodium at last!


I repeated the sodium thermite experiment and made metallic sodium this time.

I put about 3g of NaOH in a small stainless steel cookie form. Then added half a teaspoon of powdered Al. Some fizzing happened because the hydroxide had already absorbed water from air and was attacking the Al. I heat with the blowtorch and the hydrogen from the fizzing burns with orange fire. But don’t be mislead, this fire is not the reaction.

I cover the cookie form with a stainless steel dish for condensation(thanks chemoleo) and keep heating until the pan steel turns glowing red.

.|________________| dish
..........|_______| cookie pan (5cm diameter)
............... () fire

After some seconds, a really strong orange fire comes out from where the cookie pan touches the dish. The gas is cut, the reaction goes by itself. There was not enough air inside, so it has to be the Al stealing the oxygen from the hydroxide.

Everything gets really hot, with glowing red spot in the center of the dish. I put drops of water on top of the dish to cool it. They fizz jumping around.

When I removed the dish, some gray stuff was stuck where I expected the sodium to condense . Just some grains, but was a good sign. This gray stuff quickly turned white, the way I would expect sodium to oxidize. Very good sign. I drop drops of water on the grey-now-white stuff and…orange sparks! Sodium!

So…

1- Yes, it is some kind of NaOH +Al -> AlOH +Na themite like;
2- This could be improved to become a practical way of making Na;

The Hydroxide absorbs water really fast, reacting with the Al and becoming a bubbling gunk, so don’t mix them much , be quick and use coarse NaOH.

Next I will do some tests with copper pipes, if they don’t get corroded too fast, it’s the way to go. Steel pipes are too thick and would be hard do heat enough to start the reaction.

I also bet the grey stuff left in the cookie pan has lots of elemental sodium.

I have a headache now, maybe some ultraviolet in that fire. My googles are transparent.
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[*] posted on 6-12-2003 at 15:42


An interesting experiment. Seems like a fairly effective way of making DIY sodium (I presume that's what you have got; I guess it could be some finely divided Al that's reactive enough to burn or some such. Come to think of it you have the elements to make NaAlH4.)
Most goggles will block UV quite well, BTW.
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[*] posted on 6-12-2003 at 16:35


That was a facinating experiment, and by the way where did the sodium condense?
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[*] posted on 6-12-2003 at 19:57
AWESOME!


I think I will try this during Christmas holidays.
First I will grind the NaOH, in a glass without access to air, or more importantly, moisture.
Then, I will mix stoichiometric amounts. I have to think of the reaction equation first, however, to figure that one. Either Al2O3 is produced, and hydrogen, or, which is probably more likely, Aluminumhydroxide/oxide ...
I wonder whether this can be set off in a thermite like manner, like with Mg/NaClO3 etc. Will test this in 2 weeks from now or so!
More importantly, tacho, maybe fill the steel dish with water to increase the cooling effect, and hence the effect of condensation of Na vapours onto it!
If you have enough, why dont you scrape it off, and analyse it a little? i.e. add ethanol to it, and see if it reacts, while the Naethanolate forms (hydrogen bubbles should evolve)! Or, collect it, and heat it in a SEALED tupe with N2 (you can get rid of the oxygen by incubating a bit of steel wool with a small amount of acid (i.e. HCl) in a closed container, while shaking it.. there are many methods, anyone got additional info on it?)
Anyway, in an inert gas atmosphere, Na should melt quickly, more easily than Lead! This is a definite test!
Admittedly, evaporated NaOH will melt early too, so, evaporation onto the dish toghether with Na will yield a grey substance indeed.
I guess we have to find a way for separating NaOH from the Na.
I remember that Na dissolves in liquid ammonia NH3, but thats hardly applicable to us!
Anyway, great experiment, cant wait to try it myself!

Edit: how about you use some high container, like a small can, for heating your NaOH/Al, which is covered by a steel dish? this way you have a crude temperature gradient, where you collect the most volatile material (i.e. Na, NaOH) at the top!

Edit2: what is DIY sodium?????

[Edited on 7-12-2003 by chemoleo]




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[*] posted on 6-12-2003 at 22:40


DIYsodium= do it yourself sodium
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[*] posted on 7-12-2003 at 04:50


Aluminium hydroxide is not stable on heating. Why do you think it is more probable than the oxide?
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[*] posted on 8-12-2003 at 02:27


NaAlH4 is a possibility, unfortunately. Does it give orange sparks with water?

Chemoleo, do not grind the NaOH or mix it too well with the Al. It absorbs water from air and its enough to begin a very exotermic reaction. NaOH will have to melt long before the "thermite" begins, which bring us to another caveat: Don't use glass. You have to heat the mix to glowing orange before the reaction begins. I think that's way beyond glass capabilities (sp?). Besides, molten NaOH is a nice solvent for glass.

I will give it a try in larger scale, but it has to be outside the house, and it's raining these days. I don't have to say that this is all quite dangerous in a larger scale. Molten NaOH is... well, nasty. And, yes, I will put water in the dish.

You can give it a try in very small scale first, with a stainless steel spoon, pliers, a blowtorch, gloves and googles. Just to see how it happens.


Can we post pictures in this forum?
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[*] posted on 8-12-2003 at 02:54


About what temperature do the reactants have to be heated to?
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[*] posted on 8-12-2003 at 06:06


The references that said sodium was obtained reacting iron powder and NaOH state temperatures above 1000ºC. I had to heat the crucible (ss cookie pan) to orange, I believe that’s around 1000ºC. But I gess only a portion of it has to be so hot, since the reaction generates a lot of heat.

check:

http://intro.chem.okstate.edu/ChemSource/Alkalimetals/alkmet...

“Gay-Lussac and Thénard (1808) showed that molten caustic potash (KOH) or caustic soda (NaOH) brought into contact with red-hot iron turnings produced the respective alkali metal as a distillate.
Castner (1886) produced sodium on a large scale by heating NaOH with iron and carbon at a temperature of 1000 °C:


Another thing:

The cement-like residue reacts quite violently with water, but generating only hydrogen (?), no sparks. As unionized has pointed out, NaAlH4 is a possible result here. This residue could be a strong reducing/hydrogenating agent. Could anybody suggest an easy test for strong reduction?

I know this is not electrochemistry, but here is where sodium production is being discussed.
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[*] posted on 8-12-2003 at 06:45


If we ignore the formation of NaAlH4 for the moment, the theoretical reaction would be:

Al + 3NaOH --> Al(OH)3 + 3Na

2 Al(OH)3 --> Al2O3 + 3H2O
-----------------------------------

Overall:

2 Al + 6 NaOH --> Al2O3 + 6 Na + 3H2O

Depending on the temperature though, you may get AlO(OH), which yields Al2O3 with further heat (I was referring to this unionised).
Anyway Tacho, maybe try this with weighed stoichiometric amounts! This in a sealed container, with a small pressure outlet, in an atmosphere that is preferably devoid of O2 etc.
I am sure this wold work nicely!
Anyway, pls test your grey substance with ABSOLUTE ethanol - it should evolve H2 if it contains free Na. If it is NaAlH4, it probably shouldnt (altho i dont know), as Ethanol cant be reduced further...

PS hey Tacho, I played lots with molten NaOH and know of its hygroscopicity, so dont worry for my safety... yet... until I dump 500 g of metallic sodium into my bath tub :D:D
PS yes you can post pictures, look for the 'new feature for sciencemadness' thread.

[Edited on 8-12-2003 by chemoleo]




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[*] posted on 8-12-2003 at 10:58
I think Tacho is on to something here, but...


Quote:
Originally posted by chemoleo
the theoretical reaction would be:

Al + 3NaOH --> Al(OH)3 + 3Na

2 Al(OH)3 --> Al2O3 + 3H2O
-----------------------------------

Overall:

2 Al + 6 NaOH --> Al2O3 + 6 Na + 3H2O


[Edited on 8-12-2003 by chemoleo]


Al2O3 + 6 Na + 3H2O as products?!?

wouldn't (6Na)2 + (3H20)2 ==> (6Na(OH)) + 3H2 ???

Since Chemoleo loves poetry so much I feel the need to express this another way..:D

------
Said a globule of sodium, so sweet
To a droplet of water petite,

“Let’s unite in the air, rearrange ourselves there, and give off a great deal of heat!”

-----------
***Thanx Tacho!***
P.S.

"By 1890, Castner developed a large scale electrolytic method for preparing sodium using a cylindrical iron pot with an iron cathode and a nickel anode"

Now, I've had a little trouble trying to puzzle a way out (in my mind) to electrically insulate the anode's and cathodes from the stainless steel container.

I beleive if I was going to construct a down's cell using carbon anode/cathode's I would have similar trouble with sealing the electrodes to the vessel, both physically and electrically.

It has occured to me to use a ceramic plug to hold the electrodes but then I think that the expansion of the Stainless steel pipe bottom might crack the ceramic inserts spilling molten salt and sodium on the garage floor followed by a little chlorine gas.

In case anyone is wondering, It would seem to me to be easy to use a foot long peice of 3" stainless pipe from online metals.com as my reaction vessel. If I was going to attempt to build a Na Cell.
--------
Maybe the Castner heat/charcoal process is a little more realistic than dealing with the juice.

The first Castner Process... 6NaOH + C==>2Na+3H2+2NaCO3 seems quite workable using ceramic crucibles and a single pole (sublimator type)condensor.


[Edited on 8-12-2003 by Hermes_Trismegistus]




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[*] posted on 8-12-2003 at 12:41
good point... with a FATAL flaw :D:D


Lol Hermes aka Bob,

you have a point there, thanks for pointing it out! Sadly though, your poetry seems to distract you from thinking this through :D - do what you do best!

Ok, the reaction of the products will occur, as you said this is

2Na + 2H2O --> 2 NaOH + H2.

There you failed to notice the fatal flaw, which is
HYDROGEN IS PRODUCED! You are effectively removing hydrogen out of the system, leaving oxygen only!!

So what does that mean in terms of reaction equations?

1] 2 Al + 6 NaOH --> Al2O3 + 6 Na + 3H2O

2] 2 H2O + 2 Na --> 2 NaOH + H2

As you are eliminating water in this process, multiply 1] times 2, and 2] times 3.

This gives

1'] 4 Al + 12 NaOH --> 2 Al2O3 + 12 Na + 6 H2O

2'] 6 H2O + 6 Na --> 6 NaOH + 3 H2


Overall, after eliminating H2O, we then get the final version:

--------------------------------------------------------
4 Al + 6 NaOH --> 2 Al2O3 + 6 Na + 3 H2
--------------------------------------------------------

!!!!!!!!

The beauty of chemistry ey? :D :D

This is why the reaction DOES work!!

Let me rewrite that poem:

**********************************
Said a globule of sodium, so sweet
To a droplet of water petite
I will eat you up no worry,
but it farted HYDROGEN and it was sorry!!

**********************************

:D:D:D


[Edited on 8-12-2003 by chemoleo]




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smile.gif posted on 8-12-2003 at 19:12


Df for NaNO3 and Al2O3 are -424.8 & -1669.8 KJ
NaNO3    > Na(s) + 0.5 N2 + 1.5 O2 - 424.8 KJ (in thermodynamic you can use any theorical equation)
1.5 O2 + 2 Al    > Al2O3 + 1669.8 KJ
Na(s)    > Na(g)    DHsubl= +108 KJ
1669.8-424.8-108=1137 so
NaNO3 + 2 Al    > Na(g) + 0.5 N2 + Al2O3 + 1137 KJ :o

Na2CO3 Df= -1130.9 KJ/mol
1669.8-1130.9-108=430.9
Na2CO3 + 2 Al    > 2Na(g) + C + Al2O3 + 430.9 KJ

NaOH Df= -426.7 KJ/mol
((1/3)*1669.8)-426.7-108=21.9
NaOH + 2/3 Al    > Na(g) + 1/3 Al2O3 + 0.5H2 + 21.9 KJ




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[*] posted on 8-12-2003 at 20:20
Yep it works nicely...


5.4 grams Al powder.
8.6 grams NaOH prills.

Substances placed into an empty Campbells soup can. Can was crimped slightly so pliers could grasp it. An empty tuna can served as a lid and was crimped slightly to grasp the soup can (didn't want it flying off when the H2 ignited).

Holding onto the can with pliers, the container was held over the flame of a bunsen burner for about 2-3 minutes when a nice hissing sound began to grow louder. As the hissing was increasing in intensity, a burst of hydrogen ignited and proceded to burn from around the edges of the tuna can lid. The container which was brought to a dull red by the burner was now a bright orange. The reaction subsided within 30 seconds or so and I let the container cool with the aid of some snow (leaving the lid on of course).

I looked inside and saw a grey, fused mass. This was chipped out of the container with minimal effort and isolated. A small portion was grasped with forceps and placed into a jar of water resulting in vigorous fizzing. I could see the mass falling apart (into NaOH and Al2O3 no doubt) at the bottom of the jar with many bubbles of H2 coming off. The rest of the mass was placed into a pyrex jar for storage. I took the reaction cans outside and poured water on them and got some nice fireballs from the sodium adhering to the sides.

The product weighs 9.2 grams and has gotten coated with a faint white haze. I can see many small spherical, metallic globs sticking out of the grey chunks which must be sodium with an oxide coating. I am also not storing the product under kerosene because I don't care about the oxidation. I am also aware of the non-stoichiometric experiment. This was slightly intentional because I wanted the sodium to be trapped in a lot of Al2O3/Al.

Now to design a larger scale process and then a purification step involving melting the sodium and somehow removing the Al2O3....

Awesome.
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[*] posted on 8-12-2003 at 20:49
purification


Why couldent you just heat the unupurified sodium/Al2O3/Al mix(in an inert gas) untill the sodium melts and then pour off the molten sodium (the Al2O3 and Al shoud sink due to their greater density).
Can anyone see a problem with this?
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