hermanntrude

No, you CAN'T make sodium!

64 posts in this topic

Could you not electrolysise molten sodium chloride? i know it would be a hugely high temperature, but could you not theoretically do it?

 

Yes. It is in fact the preferred method of mass scale sodium production. They use doping agents to lower the melting temperature though. However, I doubt that anyone does this at the lab scale, liquid sodium chloride is some really nasty stuff.

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and anyway, there must be some type of way of extracting sodium from compounds, or how would we get hold of any of it?

 

oh wait, couldn't you reduce sodium compound with carbon monoxide?

 

Yes. It is in fact the preferred method of mass scale sodium production. They use doping agents to lower the melting temperature though. However, I doubt that anyone does this at the lab scale, liquid sodium chloride is some really nasty stuff.

ahh yes thanks, i thought there must be some way of doing it

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and anyway, there must be some type of way of extracting sodium from compounds, or how would we get hold of any of it?

 

oh wait, couldn't you reduce sodium compound with carbon monoxide?

 

CO isn't a strong enough reducing agent. [math] E^o [/math] for the reaction [ce] Na^{+} + e^{-} -> Na^{0} [/ce] is about -2.71 V relative to the standard hydrogen electrode. One would need a really strong reducing agent as the Na metal that would form will be very reducing itself.

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In the immortal words of Barack Obama, YES WE CAN. I've made sodium metal several times, the result is impure but works.

In case anyone wants to know how, it was a NaOH and Mg thermite reaction.

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There are at least three other ways to make sodium metal, that do not involve electrolysis.

 

Aluminum metal can actually reduce sodium hydroxide to sodium metal. Now I realise that many of you chemists will immediately say this is impossible, because "sodium is a more reactive element than aluminum", so here is a link showing pictures of the reaction, but with potassium being made from magnesium, instead of sodium from aluminum. http://sites.google....allic-potassium

 

It also works if solid NaOH is ignited with Al powder in a metal container, and a lid placed over it to prevent reaction with air. The sodium cannot be obtained in a pure state by this method, however, as it is mixed with slag. But it is still charactaristically reactive with water.

 

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

 

but note that

AlCl3 + (3)Na --> (3)NaCl + Al

 

It is possible to prepare sodium metal by cautiously heating sodium azide in the absence of oxygen.

 

(2)NaN3 --> (2)Na + (3)N2

 

Although lithium can burn in nitrogen, both sodium and potassium nitrides are very unstable. Sodium nitride decomposes into elemental sodium, giving off nitrogen gas, at only 87°C.

 

Comparing the decomposition of other metal nitrides

 

Similarly, the explosive decomposition of copper azide also results in the separation of the constituent elements, but this reaction happens for very different reasons.

 

Cu(N3)2 --> Cu + (3)N2

 

But the same reaction with iron (which is dangerous) will result in iron nitride.

 

(3)Fe(N3)2 --> Fe3N2 + (8)N2

 

The iron nitride can be decomposed to elemental iron and nitrogen gas above 800°C.

 

Fe3N2 --> (3)Fe + N2

 

The decomposition of calcium azide is similar to that of iron.

Ca(N3)2 decomposes above 110degC, explosively so over 140degC. The Ca3N2 that forms only decomposes at 1600degC, at which point the elemental calcium simultaneously vaporizes out with the nitrogen.

 

 

Creative Way to Make Elemental Potassium?

 

An idea for chemical preparation of elemental potassium, which does not require electric current. It would be impractical, but very creative. Not sure if all the reactions would work.

 

Ca3N2 + (6)KCl --> (3)CaCl2 + (3)K2 + N2

 

Distilling calcium nitride with potassium chloride in with steel-walled distillation may cause potassium to boil out. This proposed reaction would make use of Le Chatelier's principle. Although potassium boils at 759°C, it is possible that molten potassium could be produced below this temperature.

 

 

(6)CaCl2 + Ti3N4 --> (2)Ca3N2 + (3)TiCl4

 

The titanium nitride (m.p. 2930°C) would be crushed into a fine powder and distilled under intense heat with calcium chloride. Titanium tetrachloride (TiCl4) is a liquid which boils at only 137 °C.

 

 

(3)TiI4 + (16)NH3 --> Ti3N4 + (12)NH4I

 

I think titanium tetraiodide (b.p. 377 °C) could be reacted with anhydrous ammonia gas to form titanium nitride and ammonium iodide. I am not sure if the NH3 could be bubbled into molten TiI4, or if the TiI4 would need to be in the vapor phase, with the intense heat required for the reaction. The reaction would be expected to procede because TiI4 is very acidic, and because the titanium-nitrogen bonds are stronger than titanium-iodide. Wikipedia claims that TiCl4 "with ammonia, titanium nitride is formed"; this is not surprising since TiCl4 reacts with water to form titanium dioxide and hydrogen chloride.

 

Titanium tetraiodide melts at 150 °C. It can be prepared from easily obtainable materials:

(3) TiO2 + (4) AlI3 --> (3)TiI4 + (2)Al2O3

Edited by Anders Hoveland
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Could sodium be used as a battery like lithium

Probably only by someone able to understand what "off topic" means

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Probably only by someone able to understand what "off topic" means

lol IK but I didn't want to start a new thread and its fairly close to the thread topic. Edited by dragonstar57
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I think this thread's title should be changed to "Yes, you CAN make Sodium!", since there exist several different routes to preparing the element sodium, without electric current, two of which have been the subject of much successful experimentation by amateur home chemists.

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Heck, if people want boom, why dont they get stuff like NI3. We made it at our school lab and its pretty cool. Explodes on touch.

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Heck, if people want boom, why dont they get stuff like NI3. We made it at our school lab and its pretty cool. Explodes on touch.

sounds safe and like something a home chemist would want to handel

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I have interest in sodium as an organic chemist. Strong bases like sodium amide and alkoxides require sodium metal for any decent yeilds of the reagents. My idea is based on the mercury process for production of sodium hydroxide. I would be insterested in feedback as the method doesn't involve high temperatures or any reactive metals other than the produced sodium which I feel makes it safer (mercury is a much more predictable hazard than highly reactive metals exposed to who knows what at 1000+ degrees C).

 

I am reluctant to go into detail as I havent had the chance (mercury is hard to come by) to test it and don't want anyone hurting or poisoning themselves trying to copy it.

 

So just basically, the THEORY behind the method is to:

 

  • Form the sodium in the mercury as per the mercury cell process for sodium hydroxide production.
  • Collect the amalgam
  • Remove mercury from the amalgam by vacuum leaving behind sodium metal
     
    The sodium would be contaminated with mercury- but seeing as im using this for reagents thats not so much of a problem as the mercury probs wouldn't react.

I want to re-inforce that this is just an idea that I want feedback on, it may not work and it may be extremely dangerous particularly for amature chemists.

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setup.jpg

 

In above picture, 6.12 g KOH (flakes), 3.12 g Mg, 50 ml Shellsol D70, and 1.02 g 2-methyl-2-butanol are being heated in a sand bath to 200 °C, such a reaction (conducted in the absence of oxygen) can produce globules of metallic potassium after 2 hours.

 

 

 

I have been thinking about the possibility of reducing sodium hydroxide with aluminum foil. So a calculation of the expected enthalpy of formation of such a reaction may be helpful, to get some idea as to whether such a reaction would be expected to be favorable.

 

The enthalpy of formation for Al2O3 is -1669.8 kJ/mol, while the value for Na2O is -414.2 kJ/mol.

 

As Al2O3 contains 3 times as many oxygen atoms per mol, 3 times 414.2 equals 1242.6, which is still less than 1669.8, so aluminum has more affinity for oxygen than sodium. And indeed an exothermic thermite reaction between sodium hydroxide and aluminum powder can produce sodium.

http://www.youtube.com/watch?v=908rjHQ5mmc

 

The enthalpy of formation for AlCl3 is -705.63 kJ/mol, while the value for NaCl is -411.12 kJ/mol.

 

As 3 times 411.12 equals 1233.36, sodium has more affinity for chlorine than aluminum. And indeed, the reduction of aluminum chloride by elemental sodium was first done by H. Sainte-Claire Deville, although H. C. Ørsted had previously used potassium instead.

 

 

But of course the interaction with the alcohol would affect the enthalpy of formation, increasing the affinity of sodium for oxygen. A quick estimation of this effect can be made by comparing the enthalpy of formation for sodium hydroxide, which is no doubt even more favorable than sodium alkoxides (sodium alkoxides vigorously hydrolyse with water).

 

NaOH -425.93 kJ/mol

H2O -285.83 kJ/mol

Na2O is -414.2 kJ/mol

 

So the hydration of sodium oxide to anhydrous sodium hydroxide should release 151.83 kJ for each mole of Na2O reacted.

Na2O + H2O --> 2 NaOH

 

So it can be inferred that the presence of tert-butanol would not significantly affect the affinity of sodium for oxygen, meaning that the reduction of a sodium alkoxide by aluminum should still be energetically favorable.

 

 

The competing affinities between sodium and aluminum for fluorine apparently is more complicated:

It will be noted that when aluminum fluoride is in excess to that contained in cryolite (NaF)6Al2F6, aluminum does not reduce sodium fluoride, and on the other hand, when sodium fluoride is in excess, aluminum does reduce sodium fluoride.

Metallurgical and Chemical engineering, Volume 11, p178 (1913)

Edited by Anders Hoveland
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If you wish to remove mercury from sodium, distillation is your best option. Mercury boils at 357 C, while sodium boils much higher at 883 C. A good hotplate can easily achieve 360 C, and a ground-glass setup would probably help. Plus, you get your mercury back!

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If you wish to remove mercury from sodium, distillation is your best option. Mercury boils at 357 C, while sodium boils much higher at 883 C. A good hotplate can easily achieve 360 C, and a ground-glass setup would probably help. Plus, you get your mercury back!

 

I've only seen boiling mercury once, and even now, it makes me nervous just thinking about it.Yes, you CAN make sodium! I've done it myself. I would know.

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I've only seen boiling mercury once, and even now, it makes me nervous just thinking about it.Yes, you CAN make sodium! I've done it myself. I would know.

 

Try boiling bromine. Much, MUCH worse.

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Try boiling bromine. Much, MUCH worse.

No.

 

Boiling mercury is very dense- it sloshes around and, if/when it bumps, it breaks the glassware. The glass is half way to it's maximum temperature limit and that reduces its strength.

Then it runs down into the heater and boils off.

The vapour floods out and contaminates the area: that area stays contaminated until someone puts a lot of effort into a clean up.

 

The bromine is much less likely to break the glass. If it does, a fair bit of it will run down onto the heater and will, quite probably react with it. So your heater's dead.

Very sad, give it a decent funeral.

The remaining bromine is dispersed into the atmosphere and diluted to a point where it's harmless.

 

Seriously, which site would you be happier visiting an hour later, a bromine spill or a mercury spill?

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If you have metallic K, then just add NaCl making Sodium in situ (of course, still doing this safely, is the issue). But, assuming one could address this problem, how would one safely make K? Someone mentioned Sciencemadness which is currently running a very long thread on making Potassium at much safer ambient temperatures via Mg turnings and KOH and (for this forum as to not violate rules with respect to recipes for dangerous substances like metallic K) called the alcohol ROH. This could also give insight as to a low temperature Sodium preparation. Since the Potassium reaction temperature is around 300 C (correct me, it is a 53 page thread), my take on the reaction mechanism:

1. Aqueous phase since as much as 27% of KOH could be water:

Mg + 2 KOH + H2O --> [K2MgO2 + H2O] + H2* (g)

K2MgO2 + H2O <---> 2 KOH + MgO

Net:
Mg + H2O --KOH--> MgO + H2* (g)

2. Non-aqueous reaction:

KOH + ROH --> ROK + H2O

Mg + H2O --KOH--> MgO + H2* (g)

ROK + 1/2 H2* <---> ROH + K (s)

Net:

KOH + Mg ---> MgO + K (s) + 1/2 H2*

------------------------------------------------------

Comments: Now, Wikipedia actually cites the laboratory preparation for KOR from the un-named alcohol as follows:

K + ROH → ROK + 1/2 H2 (g) + Heat

where ROK species is, itself, noted as being a strong, non-nucleophilic base in organic chemistry.

So, what I am postulating here is that activated hydrogen is formed (most likely via chemisorption, discussed more below) and that with excess active H2* (from the initial dehydration), pressure (balloon employment), and heat applied to the reaction chamber, that the ROK formation reaction, cited above, is reversed to some extent releasing K (in agreement with Le Chatelier's principle to remove stresses relating to temperature, pressure and/ or concentrations):

ROK + 1/2 H2* (in excess) + Heat (applied) ---> ROH + K (s)

Now, more on this chemisorbed hydrogen most likely created via the presence of MgO. Here is an abstract (source: "H2 chemisorption and consecutive UV stimulated surface reactions on nanostructured MgO", in Phys. Chem. Chem. Phys., 1999,1, 713-721. To quote from the abstract:

"MgO nanoparticles obtained by chemical vapour deposition (CVD) were exposed to H2 and subsequently to UV irradiation and/or molecular oxygen at room temperature. A combined IR/EPR study reveals the role of low coordinated surface sites and anion vacancies in the diverse reaction steps. The hydride groups emerging from the initial H2 chemisorption processes (heterolytic splitting) play an active role in the consecutive reactions. They provide the electrons which are required for the UV induced formation of surface colour centres and for the production of superoxide anions (redox reaction). Both the colour centres and the superoxide anions are EPR active. The hydroxy groups resulting from H2 chemisorption do not actively participate in the consecutive reactions. Together with the OH groups formed in the course of colour centre formation they rather play the role of an observer. They undergo specific electronic interactions with both the colour centre and the superoxide anion which are IR inactive (or IR inaccessible) surface species. They may, however, be observed by IR spectroscopy via the specifically influenced OH stretching vibrations. This proves the intimate interplay between IR and EPR spectroscopy as applied to the surface processes under investigation. As a result, two paths were found for the three consecutive surface reaction steps: H2 chemisorption, colour centre formation and superoxide anion formation. In the first one a single, well defined surface area element is involved, namely a low coordinated ion pair, the cation of which is a constituent of an anion vacancy. In the second path a diffusion controlled intermediate step has to be adopted in which the electron required for the colour centre is transported by an H atom travelling from a hydride group to a remote anion vacancy. In either case there is clear experimental evidence that the finally resulting superoxide anions are complexed by the colour centre cations."

See also "Theoretical aspects of H2 and CO chemisorption on MgO surfaces", Surface Science (May 1982), 117 (1-3), pg. 571-580, to quote from the abstract:

"Preliminary ab initio calculations at the SCF level and beyond are reported for the chemisorption of H2 and CO at the (001) surface of MgO. It is concluded that the dissociative chemisorption of H2 requires the presence of defects and that at anion vacancies, V− centres and self-trapped holes the overall process is exothermic in each case. It is predicted to be non-activated at anion vacancies and possibly the same at the other two defects. Binding energies are calculated for the interaction of CO with a non-defective (001) surface of MgO and at impurity ions therein. They range from 2.5 kcal/mole at Al3+ to 20.8 kcal/mole at Cu2+ and are shown to be highly sensitive to lattice relaxation of the defective surface."

where there is an interesting reference to the role of impurities in the MgO.

Now, there are also so studies citing the reaction of between hydrogen and magnesium, but mostly as fine Mg powder (or nano). However, Mg surface attacked by KOH, may be more amiable to H2. Probably, an important speculation is that absence MgO, no or reduced chemisorbed H2 formation, no reduction reaction and no K is produced! Also, less than completely pure Mg, KOH or KOH, could produce detective surfaces on the MgO, increasing yield.

I would speculate that MgO dust on Mg turnings may provide a good contact point for gaseous H2 and, with infrequent stirring to add ROK, help to form potassium. Neither frequent or very infrequent (in agreement with the patent instructions) would be advisable.

Now why is Mg dust no good for this reaction? My speculations, first, the reaction rate would be too fast (also more heat) and limit gaseous contact. Also, absence the Mg turnings, less support for formed MgO thereby limiting the H2 contact.

In closing for disclosure, the K creation is based on a patent and a questionable (at least in my opinon, but not Sciencemadness's moderator) reaction chain, depicted very differently from my opinion:

Mg + 2 KOH + H2O --> [K2MgO2 + H2O] + H2 (g)

K2MgO2 + H2O <---> 2 KOH + MgO

Net Aqueous phase (same reactions):

Mg + H2O ---KOH---> MgO + H2 (g)

Non-Aqueous:
2 KOH + 2 ROH ---> 2 KOR + 2 H2O
Mg + H2O ---KOH---> MgO + H2 (g)
2 KOR + Mg ---> 2 K + Mg(OR)2
Mg(OR)2 + H2O --> MgO + 2 ROH

Net reaction:
2 KOH + 2 Mg ---> 2 MgO + 2 K(s) + H2

or:
KOH + Mg ---> MgO + K (s) + 1/2 H2

where one of the main disagreement is that Magnesium shavings replaces Potassium in KOR:

2 KOR + Mg ---> 2 K + Mg(OR)2

occurring, no less, at ambient temperature (under 350 C, with K boiling at twice this) with no apparent rationale based on Le Chatelier's principle, as far as I can see. Observations suggest this is not, in effect, in the reaction chain as it has been observed that fine Mg powder does in fact produce a much lower yield (which should support, not detract, from this reaction's effectiveness), while the same observation as to the best Mg size supports the concept of surface formation on Mg for MgO to increase yield. This same observation generally suggests a more complex chemical and physical role for the Mg metal itself. With respect to documentation of this reaction, other than a patent's quasi unbalanced reaction, there is no, not even one, reputable source for this questioned reaction anyway on the internet or elsewhere. What do you think?

[EDIT] For the solid to solid reaction, I conceive that relative lattice energies and that at high tempeatures (over 1,000 C), the more volatile nature of either Na or K versus Mg, may, per Le Chatelier's principle explain Mg replacing Potassium (or Sodium), but this is not as strong an argument at a lower temperature. Also, my reaction dynamics fits many of the observed particularies (like upon changing the size of Mg employed), negative influence of faster reaction rate (limiting H2 contact), role of temperature in the reaction, best stirring frequency, and why the preparation is generally problematic (seen in select hydrogenation reactions). Also, this should not be described solely as a hydrogenation reaction, or solely as a Magnesium (or MgO) reduction reaction, but as a dissociative chemisorption of H2 activated through a Mg/Mg complex formation on a MgO surface.

Edited by ajkoer
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For those who still believe in the reactions posted previously on the road to K and Na:

 

KOR + Mg ---> 2 K + Mg(OR)2

 

Mg(OR)2 + H2O --> MgO + 2 ROH

 

Please explain to me why Mg(OR)2 gets to react with water before Potassium, and even if, it is a one pot reaction, there still should still be something observed. In other words, this accepted reaction is ignored entirely:

 

2 K + H2O --> 2 KOH + H2 (g)

 

because, if you accepted this reaction occurs at all, you would have more Hydrogen than observed, less K and some Mg(OR)2 lying around, but not reported. And why is the reverse reaction not occurring, namely:

 

2 K + Mg(OR)2 --> 2 KOR + Mg

 

More bad news is a search for verification uncovered the following,"Reactions of Magnesium and Titanium Alkoxides. Preparation and Characterization of Alkoxy-Derived Magnesium Titanate Powders and Ceramics". To quote the complete abstract:

 

"The interaction between magnesium and titanium alkoxides is studied in order to chose the best precursors for synthesis of MgTiO3. No reaction between magnesium and titanium methoxides and isopropoxides occurs. The solubility diagrams for Mg(OR)2-Ti(OR)4-ROH, R = Et,-Bu at 20°C are studied. Magnesium ethoxotitanates of variable composition MgnTi4-n (OEt)16-2nċ2nEtOH (n=2.0-0) which are structural analogs of Ti4(OR)16 (R = Me, Et) are isolated. This is a quite unusual example of statistical distribution of heteroatoms in molecular structures of metal alkoxides. Among the systems of metal alkoxides with simple aliphatic radicals only Mg(OBu)2-Ti(OBu)4-BuOH gives a convenient precursor for the synthesis of MgTiO3. A simple scheme of preparation of magnesium titanate from the alkoxide solutions is suggested. The phase purity of MgTiO3 is to a considerable extent dependent on the hydrolysis conditions. The alkoxy-derived magnesium titanate is obtained in the form of a uniform fine powder, it can be sintered into dense ceramics in the temperature range of 1140–1220°C which is 150–200°C lower in comparison with the conventional powders."

 

Note, the second sentence, 'No reaction between magnesium and titanium methoxides and isopropoxides occurs", and further, when they do react, double salts and no titanium metal precipitation. So barring a reaction with a higher alcohol with a very weak bond, the prospective of the professed reaction really needs a good source.

 

The bottom line is, you just can't make up bad stuff, it just gets worse for you.

----------------------------------------------------------------------------------------------------------

 

The good news is that if you reject this increasing unlikely reaction, and even partially accept my path, there is hope in replacing Mg in a ambient temperature synthesis of K and/or Na.

 

In fact, perhaps, Aluminum foil (contains Si and Fe impurities which are good here perhap activated with a drop of Iodine) together with dry MgO from say:

 

MgSO4 + 2 NH3 + 2 H2O --> Mg(OH)2 + (NH4)2SO4

 

Mg(OH)2 --Heat--> MgO + H2O (g)

 

may even work!

 

While the particular ROH employed is still a problem, but more research may suggests a more accessible and cheaper substitute.

 

Power to Chemistry!

Edited by ajkoer
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Dear AJKOER:

Just because your reactions met with negative feedback on one forum, does not mean you can post them on the other. There's a reason we don't believe in this 'nascent hydrogen'.

Seriously, woelen even suspended your account because you wouldn't stop postulating. Does that send any message to you? Maybe one about hijacking threads with your theories?

Besides, this thread is about sodium. So unless you have a post to make about sodium, move your post to a new thread, titled 'Nascent Hydrogen in K Production' or something, not an already existing thread about a similar element.

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Elementcollector1:

There is no 'nascent hydrogen' mentioned here. Activated H2 via chemisportion, and its many associated recents references (I only gave two) are easy found in the same very highly regarded journals I cited (note the copyright from the American Chemical Society). So forget all the non-truths to support a baseless position (asserted absolutely, no less, despite numerous possible chain reaction failings and, even more problematic, that proposition cannot explain a single experimental observed particularity).

 

Now, the open question in this Na forum, is not the formation of Sodium without electrolysis (that has been done), but can it be performed safely. Now, assume you have a reaction preparation that does form K at ambient (or safe temperatures). Does adding NaCl, mean that you now have a safe path to Sodium? Or, can you just use NaOH in place of KOH? My understanding of the reaction path for K, still being explored and apparently debated, indicate paths for which this would not be favored.

 

So, back to path exploration, here are the results of more research, and I would appreciate some honest feedback (but fear of that moderator may inhibit your honesty, I understand):

First, per Wikipedia (link: http://en.wikipedia.org/wiki/AlkoxideSection ), under heading "Thermal stability" of metal alkoxides, to quote:

"Many metal alkoxides thermally decompose in the range ~100–300 °C. Depending on process conditions, this thermolysis can afford nanosized powders of oxide or metallic phases. This approach is a basis of processes of fabrication of functional materials intended for aircraft, space, electronic fields, and chemical industry: individual oxides, their solid solutions, complex oxides, powders of metals and alloys active towards sintering."

Now, this is applicable to both Na and K metal alkoxides. Also, in a recent thread at ScienceMadness, Nicodem warned Blogfast about stability issues asociated with alkoxides per his personal experience. However, the possible decomposition products cited here by Wikipedia are of particular interest, including a metallic phase (the liberation of metallic Potassium?) and/or nanosized K2O. This is in further support ot the reverse formation proposition of KOR to K that I have proposed previously.

Next, in the particular case of nanosized Potassium oxide formation, K2O (from the decomposition of KOR) could be attacked with H2 via chemisorption (on nanostructured MgO, see, for example, "H2 chemisorption and consecutive UV stimulated surface reactions on nanostructured MgO", in Phys. Chem. Chem. Phys., 1999,1, 713-721., see also "Theoretical aspects of H2 and CO chemisorption on MgO surfaces", Surface Science (May 1982), 117 (1-3), pg. 571-580.) This may not be the case for Na2O as there is no supporting literature on chemisorption.

But, even disgarding this path, there is still a possible hydrogenation reaction, although occurring rarely (per Wikipedia, link: http://en.wikipedia.org/wiki/Hydrogenation) for reactions below 480 °C between H2 and organic compounds in the absence of metal catalysts. However, with respect to rare exceptions, Wikipedia also states under the topic "Metal-free Hydrogenation", that to quote: "Hydrogenation can, however, proceed from some hydrogen donors without catalysts, illustrative hydrogen donors being diimide and aluminium isopropoxide. Some metal-free catalytic systems have been investigated in academic research. One such system for reduction of ketones consists of tert-butanol and potassium tert-butoxide and very high temperatures.[24]".

I would observe that perhaps this is related to the formation of K2O, which is apparently employed in mixed oxide catalyst for hydrogenation (see, for example, "Promotion effect of K2O and MnO additives on the selective production of light alkenes via syngas over Fe/silicalite-2 catalysts"). This path is particular to K and, as such, excludes the formation of Sodium again.

In any event, my opinion is that the simple thermal decompostion of potassium tert-butoxide (simple and perhaps now the best explanation, and also for the Na salt), by itself, or in the presence of activated H2 per chemisportion on MgO, or via hydrogenation in the presence of K2O, may be some alternate explanations with peer reviewed references.

Now, I want to be fair, some more research on the direct reaction of Mg with KOR (or NaOR), and there is nothing supporting the direct replacement reaction by Mg as proposed liberating K (or Na). For example, see "Reactions of Magnesium and Titanium Alkoxides. Preparation and Characterization of Alkoxy-Derived Magnesium Titanate Powders and Ceramics". To quote the entire abstract:

"Abstract
The interaction between magnesium and titanium alkoxides is studied in order to chose the best precursors for synthesis of MgTiO3. No reaction between magnesium and titanium methoxides and isopropoxides occurs. The solubility diagrams for Mg(OR)2-Ti(OR)4-ROH, R = Et,-Bu at 20°C are studied. Magnesium ethoxotitanates of variable composition MgnTi4-n (OEt)16-2nċ2nEtOH (n=2.0-0) which are structural analogs of Ti4(OR)16 (R = Me, Et) are isolated. This is a quite unusual example of statistical distribution of heteroatoms in molecular structures of metal alkoxides. Among the systems of metal alkoxides with simple aliphatic radicals only Mg(OBu)2-Ti(OBu)4-BuOH gives a convenient precursor for the synthesis of MgTiO3. A simple scheme of preparation of magnesium titanate from the alkoxide solutions is suggested. The phase purity of MgTiO3 is to a considerable extent dependent on the hydrolysis conditions. The alkoxy-derived magnesium titanate is obtained in the form of a uniform fine powder, it can be sintered into dense ceramics in the temperature range of 1140–1220°C which is 150–200°C lower in comparison with the conventional powders"

Link: http://link.springer.com/article/10.1023%2FA%3A1008616329847...

Interestingly here, when products are produced, mixed salts (and not deposits of titanium) occurred. Other searches on Springer's articles also only produced references to oxide formation reactions.

With time, and observations, my basic argument appears only to get more supporting paths and, so far, a total lack of rationale for the other side at the very ambient temperature recommended for this reaction. But, if I have missed something, please cited it, and unlike the others pushing their position, I welcome open supported discussion, which so far your comments are lacking, both in accuracy (no nascent H2 here, just well researched paths to activated hydrogen in journals of physical chemistry), and no apparent attempt to educate yourself by researching and citing references. Remember, good science always wins in the end.
--------------------------------------------------------------------------------------------

[EDIT] I have recently come to a somewhat uncomplimentary view, based on my research, as to how this Potassium reaction was discovered. It owes it genesis, I suspect, to a failed preparation of KOR. The KOR is formed, and subsequently, perhaps unexpectingly, decomposes, liberating K on occasion. Is this translatable to Na? Wow, what a revelation if I am correct!

Edited by ajkoer
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Ahaha, no. Not the last part, by any means - no one has found a way to make the tertiary-alcohol synthesis work for Na. And I'm not sure about the decomposition part - that would require testing with pure potassium tert-butoxide or some such, as opposed to the usual reaction mix.

 

As for the story, I admit it'd be humorous, but unless I see some citations, I'll think of it as nothing more than a story. Chalk it up to scientific skepticism, but something you just came up with in your head does not make itself true.

 

Also, isn't it a bit redundant to synthesize Na metal from K metal? Most experimenters want Na for the drying power or the reaction in water, and K is better in both respects. As for element collectors, they might go after Na from K, but would probably go with Na from Mg or electrolysis first.

 

The formation of sodium from electrolysis has a few problems with a simple, yet difficult solution. The first is the 'splattering' that tends to happen when NaOH is melted, and this is by no means fun. The second is the sodium immediately oxidizing at the surface once formed, resulting in decreased or no yield. Both of these could be solved with an inert atmosphere - but setting that up can be a challenge in itself. Sodium from Mg via the reaction between NaOH and Mg metal can be viable only if, once again, there is no oxygen or moisture in the reaction container. Nighthawkinlight, who seems to have 'pioneered' this method, simply used a closed steel pot, which could be modified by an inlet and outlet tube and hooked up to an argon tank (the lid would also possibly need to be modified to seal better).

 

'Remember, good science always wins in the end'

I attempt to educate myself through actual testing - show me a test to prove some calculations, I'll do it and post the results. There are plenty of papers out there that can claim anything they want, but I would much rather see what they're talking about firsthand. To be honest, I did 'jump the gun' on the nascent/chemisorbed hydrogen, and for that I apologize, having no background on the subject but your previous posts and subsequent reactions.

 

I'd be interested to look at the reaction between NaCl and K, but would first have to check if the reaction is energetically favorable in terms of enthalpy. If it does work, and one can narrow down a good source and repetition of tert-butanol and K formation, then it would be a much simpler and possibly less dangerous path to sodium than I have seen. I wonder if such a reaction could work for lithium, or even rubidium and cesium as well...

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Actually and unfortunately, perhaps what I suggested isn't a story after all, more of a re-wording (or a re-write as it were). Stating (as per Wikipedia) ""Many metal alkoxides thermally decompose in the range ~100–300 °C. Depending on process conditions", means to me wrong temperature and process conditions and your synthesis of KOR is toast. And to further confuse the hapless would be chemist, he is presented with some potassium (or, perhaps the oxide).

 

Solution: just re-write the failed preparation as a patentably path to K (or whatever) instead. The only major issue, of course, is that you are most likely clueless as to exactly the how and why, and thus are purposefully vague as to reaction mechanics, or worse, make-up a replacement reaction as a potential explanation. Sound familar? Funny, but only until someone really takes the reaction mechanism seriously.

Edited by ajkoer
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Again, maybe that is what happened, but it's not scientific to make up a story - which you pointed out in the story.

And besides, making up a reaction-mechanism, however wrong it may be, is the first step to pinning down the real thing. Who knows, maybe this hypothetical chemist got it on the first try. Maybe not. We'll need to think of some tests to prove that the originally proposed reaction, as it stands, is wrong, and that a different reaction would be better. I'm interested in cutting the Mg content in half, but I still don't understand what that would do. Another test would be to simply heat up KOR under solvent and see if any K forms.

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