Sciencemadness Discussion Board - Sodium by chemical methods

Sodium by chemical methods

Theoretic - 24-6-2003 at 07:29

I heard that charcoal would reduce soda on strong heating to give CO and Na, which are both volatile

, so the equilibrum would be shifted in favour of products. Would anyone try?

vulture - 24-6-2003 at 08:08

I think this would yield the same products as with potassium, being C6O6Na6 (in fact hexasodium salt of hexahydroxybenzene) and NaOC#CONa (# = triple bond).

These compounds tend to be explosive.

Polverone - 24-6-2003 at 13:12

Actually, the explosive compounds won't form when making sodium instead of potassium. See the Muspratt (http://bcis.pacificu.edu/~polverone/muspratt.html) entries on sodium and potassium. In any case, the production of alkali metals by reduction with carbon seems at least as difficult and hazardous as the production of white phosphorus from phosphates.

GZAust - 11-7-2003 at 07:38

Polverone, I don't think so, as far as difficulty or hazard.

Well not if you have access to a furnace or kiln.

Make a vessel shaped like an upright lid-less cyllender. with a long nipple to attach an inert gas inlet.

pump up the inert gas when your transferring the sodium metal to a safe storage solvent.
haha that isnt dangerous at all!

opening a bottle of 30% HCl is more dangerous in fact.

[Edited on 11-7-2003 by GZAust]

Organikum - 11-7-2003 at 09:45

As it sounds like you have done this GZAust, would you share some more details on the procedure?

(opened a bottle of 33% (!) HCl today several times - thats no bigger problem agreed)

What the book says

BromicAcid - 15-7-2003 at 10:54

I have a book detailing all the different methods for making the Alkali Metals, electrolytic processes are the first listed but there are three other ways, here they are, word for word taken from "Comprehensive Inorganic Chemistry Vol6" John F. Suttle D. Van Nostrand Company Inc 1957
**********************

2)Thermal Reduction Methods: The thermal reduction methods in general utilize carbon or a carbide as the reducing agent (G.L. Putnam, Ind. Eng. Chem. 30, 1138; 1938)

6 NaOH + 2 C ---> 2Na2CO3 + 2Na + 3H2

A mixture of rubidium chloride or cesium chloride with calcium carbide heated to 700 - 900 C in vacuo gives a 75% yeild of the alkali metals. (V.D. Polyakov and A. A. Fedorov, J. Applied Chem. USSR 13, 1833-8; 1940) With sodium chloride the temperature of 950 C is used. (P.V. Gel'd et al., J. Applied Chem. USSR 20, 800-8; 1947) The production of potassium is reporded using silicon or calcium carbide as the reducing agent at a temperature o 1000-1150 C.

2KF + CaC2 ---> 2K + CaF2 + 2C

Part of KF may be substituted by K2CO3 or K2SiO3 without any loss in yeild.

2K2CO3 + 3Si + CaO --> 4K + 2C + 3(2CaO*SiO2)

These methods usually require good vacuums at high temperatures.

3) Metal Replacement: The replacement of an active metal from its salt by a less active metal may appear to be contrary to the oxidation potentials of the two metals but can be accomplished sucessfully in many cases because the difference in volatility of the two metals at elevated temperatures will displace the equilibrium. (J.H. DeBoer, P. Clausing and G. Zecher, Z. anorg. allgem. Chem. 160, 128-32; 1927) Rubidium and cesium can be displaced from their salts by iron, (L. Hackspill and H. J. Pinck, Bull. soc. chim. 49, 54-70; 1931) the temperatures necessary to accomplish this depends upon the salts involved. For example, iron displaces the alkali metals from sulfates and arsenates at their melting point, from thiocyanates at 650 C, from borates and phosphates at 1300 - 1400 C.

Potassium and sodium-potassium alloys are prepared by the reduction of molten KCl with sodium. (Chem. Eng. News 33 (7), 648; 1955) Cesiuim is obtained from pollucite by reduction with calcium.

4) Decomposition: Alkali metal compounds, such as the hexacyanoferrates (II), cyanides, and azides can be decomposed into the alkali metal by heating.

4KCN ---> 4K + 4C + 2N2

*********************
When it says carbon I don't think it means charcoal, it means straight industrial coke.

Theoretic - 24-7-2003 at 06:57

Alkali-metal cyanides=>metal+carbon+ +nitrogen?

What sort of temperatures does THAT need? Why aren't cyanides then decomposed at the moment of formation?

BromicAcid - 24-7-2003 at 18:32

I have no clue as to the temp, I copied the book word for word, I just tried to look up potassium cyanide in two of my chemistry books to see if there is some decomposition temperature listed, there is a melting point listed as 634C but no boiling point or decomposition point. I did a quick google search but didn't get any relevent results but if it does decompose as stated then the temp must be somewhat higher then 634C I wouldn't want to work with cyanides at that temperature.

blip - 24-7-2003 at 23:21

Quote:

Why aren't cyanides then decomposed at the moment of formation?

The carbon and nitrogen atoms have a triple bond, like in diatomic nitrogen. Also, the alkali metal cation likely wouldn't readily accept another electron without a good deal of energy.

blip - 25-7-2003 at 19:44

Decomposition of sodium azide wouldn't be so bad if you had a way to contain the large, abrubt gas output yet prevent the sodium from reacting with the containment vessel (I make this sound too easy).

akinmad and xoo1246's idea is <a href="http://www.sciencemadness.org/talk/viewthread.php?tid=518" target="_blank">here</a> and <a href="http://www.sciencemadness.org/talk/viewthread.php?action=attachment&tid=518&pid=7637" target="_blank">here</a>.

NH4NO3  210°C > 2H2O + N2O
Basic reaction: NaNH2 + N2O <strike>  </strike>> NaN3 + H2O
2NaN3  275°C > 2Na + 3N2

[Edited on 7-26-2003 by blip]

additional note

Polverone - 27-7-2003 at 21:48

I agree that this method is very impractical on a lab scale. However, it's worth noting a modification: someone (forget who, this is from old Usenet posts) in the late 1800s greatly improved the process. By mixing iron filings and pitch, then heating the mass, they were able to obtain carbon adhering to iron that wouldn't float in the NaOH. This considerably increased the speed and lowered the cost of the reaction, but the modified process was displaced by electrochemical production not long after.

Nick F - 28-7-2003 at 06:55

I've been thinking about getting some sodium from sodium azide, since I have a very large amount of it (more than I could ever use for primaries), but I have not yet done any experiments.
275*C, blip? That shouldn't be too hard then. Candle wax has a bp of 250-300*C, so unfortunately it probably couldn't be done under that, but I'm sure there must be a common substance that could be used to protect the sodium as it is formed.
Then there is the problem of all the gas, which could possibly spray a near-boiling hydrocarbon all over the place if it's generated a little too quickly. And if I'm heating it over a bunsen...
Piles of sodium azide don't seem to burn too vigorously, but it might be different if it's gradually heated.

I suppose you could just heat the azide in a vessel with a narrow opening (easily made from plumbing bits), and cap it off before it cools, or use the cooling gas to suck in oil to protect the metal. Then you haven't got the fire hazard and only a small amount of air would be present to react with the sodium formed. Hmmm... maybe you could gradually pour azide down into the opening. This would mean you could control how vigorously the gas was made by the rate of addition.

Another thing I was wondering about was using sodium azide as a reducing agent to make phosphorus. The azide decomposes exothermically, to make a load of hot sodium, which would then reduce the phosphate. The nitrogen produced forms an inert atmosphere to protect the phossy, and it also carries it over into the recieving vessel and helps to "flush out" all the remaining phossy vapour in the reaction vessel, making it safer to open up afterwards in order to prepare the next batch.
But I was too worried about an explosion to try anything!

[Edited on 28-7-2003 by Nick F]

Marvin - 28-7-2003 at 21:01

xoo, The internet is always a suspect source of information, that site is a prime example. It looks fantastic, the descriptions of the practical aspects of the chemical processes are right, the heats of reaction are probably accurate, but its been written by someone with a grade school understanding of chemistry. He is taking things he knows about, information from books and making up the actual chemistry according to his understanding of the subject, which is rather unwise at the best of times. In particular he is making the common grade school mistake of assuming that CO2 is always the end product of carbon used as a reducing agent.

I particulaly like the following statement, after reviewing his site,

Quote:

The direct reduction of iron (III) oxide with carbon is sufficiently endothermic that it does not proceed, but reduction with carbon monoxide is possible:

2Fe2O3 + 3C 3CO2(g) + 4Fe, DG0 = +301.3 kJ/mole

Fe2O3 + 3CO(g) 3CO2(g) + 2Fe, DG0 = -29.4 kJ/mole

This goes against neerly 3000 years of practical iron smelting and I find this particulally funny, not to mention rather ironic, considering his website is called '3rd millenium online'.

CO is the end product of most of these reductions and any CO2 produced as above would immediatly be reduced to CO by the carbon in the mixture.

Na2O is indeed produced when a mixture containing carbon and sodium carbonate is heated, but this is almost certainly due to Na2CO3 + C => Na2O + 2CO, not Na2CO3 => Na2O + CO2. His equation for the ultimate production of Na is more spectaulaly wrong though, Na2O + CO => CO2 + 2Na is laughable becuase sodium will actually burn in a CO2 atmosphere. Speculating using the equations he has produced as guesses of what is going on wont lead anywhere. The real shame, is that its a potentially very good site the reactions are frequently wrong for the processes. The 'right' reactions, it should be pointed out, arnt what is going on 100%, but they provide a workable understanding of a reaction. Real chemists spend massive amounts of other peoples money finding out what reactions are really going on in these industrial processes, and to what extent. The guesses on the site wont get anyone anywhere though.

Alkali metals from carbonates, all except lithium can be done by reduction of carbonates with carbon (producing carbon monixide) or iron (producing iron oxide) in iron retorts (Mellor describing the Deville process) at probably somewhere between 1000 and 1500C (my guess). Iron being one of the few things that will survive the heating. It also mentions that for the production of a very fine mixture of carbonate and carbon can be done by heating sodium or potassium tartrate.

An iron tube bent at 45 degrees alows production of sodium from salt and iron, the sodium condensing in the other half of the tube with the iron forming ferrous chloride (Gmelin). This is a 'white heat' process, so not much less than above. Suggested is a steel tube from a shotgun(!) with specific luting which I dont have written down. It does say that the tube only usually survives 3 runs of this process before it burns through and that the amount of sodium/potassium produced is very small.

A warning about salt eutectics to reduce melting points for reactions/electrolysis. This is a cunning method, but a sodium/potassium salt eutectic is very likley to produce an alloy as its product, with the formation of the metal eutectic providing the energy difference between sodium and potassium redox. This alloy is very hazardous, autoignition in air, liquid at room temp etc. Stick to single alkali cation eutectics.

Calcium will reduce the halides in a vacuum at 400 to 500C producing alkali vapour. Calcium itself can be produced by electrolysis of CaCl2, which is at least easy to get, in a graphite crucible with an iron electrode. Its worth considering mainly becuase the Ca can be produced solid, so you avoid a lot of the problems with alkali cells producing liquid metal. Its the only sane way for rubidium/caesium.

Ultimatly the easiest way for Na/K, will be electrolysis of molten hydroxide though. If there was an easy smart way to the metals, people would be using it.

Nick, a few thoughts, Firstly I think sodium azide decomposes at a fairly low temperature far lower than reduction would happen at, so youd really be forming sodium then reducing later on. Secondly, if you reduce with either the azide or the metal, why on earth would the reaction form free phosphorous instead of sodium phosphide? Assuming you have any P gas coming off partway through the reaction I think the sodium/sodium azide is just going to burn in it.

rikkitikkitavi - 29-7-2003 at 00:33

I think that it is worth mentioning that the process of production of Na,K is quite well described in Muspratt . Even the english text in it self is enjoyable reading (it was written more than 150 years ago)

Marvin, if electrolysing CaCl2, will not the hydrate decompose when heated , giving off HCl leaving CaOx*CaCl2y ? Just like MgCl2?

/rickard

Nick F - 29-7-2003 at 03:03

"Anhydrous calcium chloride" is easy to get as a water-absorbing substance for use in dehumidifiers. At my local shop it's £1.99/500g. Although I suppose it's possible that it is actually a mixture of the oxide and chloride, produced by heating the hydrated chloride, given an inaccurate name. If it is in fact CaCl2 then that could be used.
But I didn't think that the hydrate did that to a very significant degree? I know Mg and especially Be chlorides do but I thought calcium chloride was sufficiently non-covalent for it to be dehydrated by heat without problems..?

Marvin, yes I know the azide decomposition alone wouldn't be enough, it would also be heated from outside but I didn't make that clear. But yes, phosphide formation would be a problem, I hadn't thought that far ahead (edit: I had thought up to the point in the reaction were it explodes and sprays burning azide everywhere)!

[Edited on 29-7-2003 by Nick F]

No one else tried this actually

metalresearcher - 18-10-2010 at 12:40

Seven years later ..... (last post in this thread was in 2003)

Nice posts it turns out that there are more chemical methods to produce Na metal of which the Na2CO3 method is the best.
So I am not so really mad ... but AFAIK I am the first one who really tried this one.

My page: www.metallab.net/Na.php
A thread on this: http://www.sciencemadness.org/talk/viewthread.php?tid=14491

[Edited on 2010-10-18 by metalresearcher]

verode - 14-11-2010 at 14:11

you get sodium with NaOH and Magnesium powder
Mg + 2NaOH = 2 Na + MgO + H2
or even try with aluminium powder
you may start the reaction with KMnO4 (few) and glicerine or magnesium ribbon

verode - 14-11-2010 at 14:15

Quote: Originally posted by verode

you get sodium with DRY NaOH and Magnesium powder
Mg + 2NaOH = 2 Na + MgO + H2
or even try with aluminium powder
you may start the reaction with KMnO4 (few) and glicerine or magnesium ribbon

bbartlog - 14-11-2010 at 16:48

Quote:

or even try with aluminium powder

That won't work; you will get sodium aluminate, to wit:

2Al + 2NaOH + 2H2O → 2NaAlO2 + 3H2

Magnesium however will do the job.

condennnsa - 14-11-2010 at 18:59

what? where did you come up with the H2O? dry molten NaOH passivates aluminum and doesn't react further.

bbartlog - 14-11-2010 at 21:40

You're right, that's in solution... does not pertain to the dry reaction. My bad. Nonetheless it won't work dry, the aluminate will still be produced rather than elemental sodium.

metalresearcher - 15-11-2010 at 04:36

Yes, NaOH and Al *do* react. Recently I put NaOH pellets and Al snippets in a 5cm long and 2.5cm diam steel tube (cosed on one side) and heated it to dull red (600oC). I saw yellow light from the tube and sometimes yellow sparks appear in the tube which means the Na metal is formed.

Wrapping DRY NaOH pellets in Al foil and heating it also works.

[Edited on 2010-11-15 by metalresearcher]

blogfast25 - 15-11-2010 at 08:25

Hmmm... I wonder how many more threads on 'Chemical sodium' we're gonna here see in the future. This one isn't really contributing anything to the topic that hasn't already been said to exhaustion elsewhere on SM.

Search SM and yee shall find!

ScienceSquirrel - 15-11-2010 at 08:31

Sodium hydroxide is not dry unless it has been held in a molten state for a long time.
It is extremely hygroscopic which is why it is not used as a primary standard for quantitative work and is usually employed in excess as a reagent.
If you leave some on the bench in a little dish for a few hours it will turn into a caustic puddle!

condennnsa - 15-11-2010 at 09:03

yea, i think that all these chemical methods on an amateur scale always produce a majority of sodium hydride, which also reacts with water similarly to sodium metal.
The only reliable way to make good sodium remains Len1's castner cell.

verode - 15-11-2010 at 10:40

Quote: Originally posted by bbartlog

Quote:

or even try with aluminium powder

That won't work; you will get sodium aluminate, to wit:

2Al + 2NaOH + 2H2O → 2NaAlO2 + 3H2

Magnesium however will do the job.

you never heard Aluminotermia
I know Al react making aluminate but not if is dry and you
make to burn it not water at all

verode - 15-11-2010 at 10:42

you make Al burn with Mg burning or with glicerine and KMnO4

hissingnoise - 15-11-2010 at 10:49

Quote:

you never heard Aluminotermia

Nope! Any relation to thermite?

verode - 15-11-2010 at 11:08

Quote: Originally posted by hissingnoise

Quote:

you never heard Aluminotermia

Nope! Any relation to thermite?

the same

verode - 15-11-2010 at 11:10

yes you must see
http://www.youtube.com/watch?v=i5IjDyyqIZE

verode - 15-11-2010 at 11:11

Quote: Originally posted by verode

yes you must see
http://www.youtube.com/watch?v=i5IjDyyqIZE

it's very very hot

verode - 15-11-2010 at 11:12

yu may also try with NaF and Al or Mg but it's more difficult
Sorry my english

blogfast25 - 16-11-2010 at 05:35

All the aluminothermic and magnesiothermic attempts at making ‘chemical sodium’ that I’ve seen (and that’s quite a few) are half-failures at best. One obtains a mixture of little sodium with a lot of slag which has proven inseparable up to and including now. At best you’ll have a few grains of contaminated sodium that do indeed test positive with water…

If you do want to successfully make gram amounts of quite pure sodium metal, melt a good dollop of NaOH in a heat resistant ceramic (or even glass) bowl, stick some graphite or steel electrodes in the melt and run a current of 5 – 10 A through it. Small globules of molten sodium will metal appear at the cathode, which can be scooped up with a gauze (steel) wire ‘spoon’ and dropped into light mineral oil where they will settle and solidify for later use. If you want to go all sophisticated use an argon or CO2 blanket to cover the ‘cell’ (and protect the sodium).

This is how Davy discovered the metal in 1807 (minus the Ar or CO2 – I saw a facsimile of the used apparatus at the British Museum, it was extraordinarily primitive) and what’s good for him should be good enough for us!

Alternatively build len1’s Castner Cell…

blogfast25 - 16-11-2010 at 05:43

Quote: Originally posted by verode

yu may also try with NaF and Al or Mg but it's more difficult
Sorry my english

In the case of 2 NaF + Mg --> 2 Na + MgF2 the reaction is likely to run so hot (heat of reaction) that the formed sodium will evaporate (well, 'zapped off' is probably a better term): sodium's BP is only 883 C. This would therefore have to carried out in a bomb reactor...