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No, you CAN'T make sodium!


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#41 chilled_fluorine

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Posted 25 September 2012 - 08:54 PM

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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#42 elementcollector1

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Posted 24 October 2012 - 03:44 PM

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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#43 John Cuthber

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Posted 24 October 2012 - 06:04 PM

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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#44 ajkoer

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Posted 21 April 2013 - 09:16 PM

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, 22 April 2013 - 04:46 AM.

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#45 ajkoer

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Posted 22 April 2013 - 02:07 PM

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.

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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, 22 April 2013 - 09:09 PM.

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#46 elementcollector1

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Posted 22 April 2013 - 11:59 PM

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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#47 ajkoer

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Posted 25 April 2013 - 12:48 AM

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....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....i/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...008616329847...

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, 25 April 2013 - 04:52 AM.

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#48 elementcollector1

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Posted 25 April 2013 - 06:02 AM

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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#49 ajkoer

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Posted 25 April 2013 - 01:48 PM

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, 25 April 2013 - 02:05 PM.

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#50 elementcollector1

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Posted 25 April 2013 - 05:40 PM

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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#51 ajkoer

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Posted 26 April 2013 - 12:06 PM

As I noted previously Mg powder is observed to be inferior, but large metallic Mg surface is needed. So one possible path is a temperature dependent reduction process directly as a result of say MgO on some metal surface. For example, see the full paper "Theoretical study of the decomposition of HCOOH on an MgO(100) surface" at http://www.qcri.or.j...011/07/p236.pdf to quote:
 
"It is well known that metal oxides are good catalysts for a variety of chemical processes [1– 26]. For example, methanol, formaldehyde and formic acid readily decompose on MgO catalysts [7,8,10–12]."
 
See also related full paper "Chemical reactivity of oxygen vacancies on the MgO surface:
Reactions with CO2, NO2 and metals at: http://www.captura.u....pdf?sequence=1 where the author notes, to quote: "It was proposed that the interaction is dominated by an electrostatic mechanism of electron transfer and that this is strictly connected to the oxygen vacancies." relating to impurities on the MgO surface. See also comments at "Acid-base reactions on model MgO surfaces" at http://link.springer...00767207#page-2.
 
So, in the current context, it may be that a particular KOR may be reduced forming potassium in a metallic phase principally due to MgO and heating based on my repeated below Wikipedia comment (link: http://en.wikipedia....AlkoxideSection ) 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..."
 
Also possible is via chemisportion occurring on the metallic Mg (but not strongly supported in the literature).
 
Now, there is apparently much literature focus, which is great news, for a nano MgO surface reactions (which can be introduced by other cheap routes, perhaps eliminating Mg metal altogether, but replacing with something to remove water and provide support for the MgO, perhaps Al foil and a drop of Iodine, Zn,...).  
 
Also, my limited research on the reaction mechanism occurring on the MgO surface is complex via the formation of some intermediaries and I would not be surprised if we actually saw some 'novel' chemistry on the MgO surface itself (as I have see in atmospheric gas/solid reactions at very ambient temperatures, for example, see full paper "Water Chemisorption and Reconstruction of the MgO Surface" at http://arxiv.org/pdf...-th/9508001.pdf). To quote:
 
"The presence of surface hydroxyl groups on MgO powders
exposed to H2O has been demonstrated by infra-red
spectroscopy7,9,10. Hydroxyls are clearly distinguishable
from physisorbed molecular water by the HOH bending mode
which disappears above 100◦C, while the OH stretching mode
persists even above 500◦C. Furthermore
there is complete monolayer coverage of the surface by
hydroxyls, as shown by microgravimetry measurements7
Despite these observations, the most reliable theoretical
calculations predict that water molecules do
not dissociate on the (001) surface."
 
and:
 
"In summary, water demonstrably chemisorbs onto
MgO but trustworthy calculations show that H2O molecules
should not dissociate on the only known stable
surface."
 
Also, see: http://www.chem.tamu...1_jp983729r.pdf
----------------------------------------------------------------------------------------------------
 
Now, with respect to your proposed use of Calcium, here is an interesting reference comparing MgO and CaO catalyst in "Natural Gas Conversion VI" edited by T.H. Fleisch, J.J. Spivey, Enrique Iglesia, page 213. Link: http://books.google.... by MgO&f=false
---------------------------------------------------------------------------------------------------
 
With respect to cutting the relative Mg use in half, I am not sure if that is the best test, but here is the logic:
 Non-aqueous reaction cutting relative Mg input:

2 KOH + 2 ROH --> 2 ROK +  2H2O

Mg + 2 H2O --KOH--> MgO + H2O + H2* (g)

2 ROK +  H2* + Heat --MgO--> 2 ROH + 2 K (s)

Net:

2 KOH + Mg ---> MgO + 2 K + H2O + H2* --> Potassium decomposition/reduced yield on water contact.
 
So this equation's indicated amount of Mg is 1/2 of the available # of moles of KOH, or actually ROH if it is less than KOH per the 1st equation in the chain. This amount of Mg is then added to the amount of Mg required, per below,  to remove all water from the aqueous first phase:
 
Mg + H2O --KOH--> MgO + H2* (g)
 
where I would assume 20% of the KOH, for example, is actually water. The hypothesis is that this total amount of Mg is insufficient for a successful Potassium production run.
 
Interestingly, note that in the usual preparation of Potassium having too much water initially with sufficient Mg around is not necessarily a bad thing as it just produces more MgO (which I claim may have a positive role in the non-aqueous phase).


Edited by ajkoer, 26 April 2013 - 08:54 PM.

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#52 ajkoer

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Posted 3 May 2013 - 06:33 PM

Came across an interesting patent "Method for the production of anhydrous potassium tert.butoxide US 4577045", link: http://www.google.co...tents/US4577045

 

 Some extracts of some possible important points as potassium tert.butoxide is presumed formed and decomposed  to Potassium in my reaction chain:

 

  "a) using cyclohexane or hexane as withdrawing agent, (b) using the tert.butyl alcohol in such an excess with respect to the aqueous potash lye and the withdrawing agent that in the bottom of the column a 10 to 18 wt.-% solution of potassium tert.butoxide is present, and the content of tert.butyl alcohol in the gas mixture situated in the center of the column is between 50 and 90 wt.-%, and © distilling out a mixture of withdrawing agent, tert.butyl alcohol and water at temperatures between 65" and 75.

 

 "Preferably, however, the tert.butyl alcohol is to have a water content of less than 0.1% by weight". So having too much water is problematic to the formation of potassium tert.butoxide."

 

 "Lyes having KOH contents of about 50% by weight can be used; the KOH content can be even lower, but then correspondingly larger amounts of water have to be distilled out. For this reason the use of potash lyes of KOH contents under 30 weight-percent is not recommended."

 

 "In the column, the reaction product that forms therein and is dissolved in the tert.butyl alcohol/hexane or cyclohexane mixture is washed into the bottom of the column, which is kept at the boiling temperature throughout the reaction. The potassium tert.butoxide is then in the bottom in the form of a 10 to 18% solution in pure, anhydrous tert.butyl alcohol."

 

  "The tert.butyl alcohol is not only reacting agent, it serves simultaneously as solvent for the potassium butoxide obtained, up to 90% by weight of total tert.butyl alcohol amount. Less than 1% by weight of total tert. butyl alcohol is part of the aqueous phase. Therefore the amount of the tert.butyl alcohol needed for the whole process should be sufficiently great that a 10 to 18% by weight, preferably a 10 to 15% by weight, solution of the potassium salt will be present in the bottom of the distillation column."

 

 Now this last comment is interesting because without sufficient excess tert.butyl alcohol, the potassium salt, lying at the base, may be capable of thermal decompositon (to K?) due to uneven heating. In general, note the relatively low temperature (under 100 C) use to prepare potassium tert.butoxide, and the possible consequences of a higher temperature (around 200 C) suggested in the Mg/KOH/tert.butyl alcohol path to Potassium.  

 

Also, in this preparation of potassium tert.butoxide, hexane acts as withdrawing agent to remove water via distillation in place of Mg. So, could one use hexane together with MgO (and possible H2) at higher temperatures to this preparation to form metallic Potassium (Sodium)? A speculation for would be experimentors.


Edited by ajkoer, 3 May 2013 - 07:21 PM.

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#53 Genecks

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Posted 4 May 2013 - 07:46 AM

 First off, I've often considered this thread to target the amateur audience who wants sodium.

 

...Also, in this preparation of potassium tert.butoxide, hexane acts as withdrawing agent to remove water via distillation in place of Mg. So, could one use hexane together with MgO (and possible H2) at higher temperatures to this preparation to form metallic Potassium (Sodium)? A speculation for would be experimentors.

 
Lay down some mechanisms, and I might consider giving this some more thought. There are multiple disadvantages that are described due to limiting and excess reagents. However, I'm not sure if using Na-OH would give a similar affect. The energy requirements may be different in attempting to make sodium tert-butoxide.
 

The potassium tert.butoxide solution is withdrawn from the bottom of the
column continuously, preferably through an overflow, and then the
potassium
salt is isolated in a manner known in itself, e.g., by distilling out
the alcohol in vacuo. It precipitates as a while, finely
granular, hygroscopic powder of high purity.

 


 
I found a nice video a moment ago on the Internet. I have not tried the experiment, but it looks interesting. It appears to be a simple way for an individual to make small amounts of sodium.
 

 

 

 

Also, sodium can be used as a catalyst:

http://pubs.acs.org/...021/j150552a012

The Catalytic Properties of Supported Sodium and Lithium Catalysts

J. Phys. Chem., 1957, 61 (6), pp 756–758
DOI: 10.1021/j150552a012
Publication Date: June 1957

Edited by Genecks, 4 May 2013 - 08:22 AM.

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#54 elementcollector1

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Posted 12 May 2013 - 01:55 AM

We've seen (and mentioned) that video several times throughout the thread. It is an interesting one, but an ever-present problem is how to collect the sodium metal at the end.


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#55 Enthalpy

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Posted 29 September 2013 - 06:11 PM

I've just read that sodium and potassium can be obtained by electrolysis of the chloride dissolved in propylene carbonate.

http://en.wikipedia....ylene_carbonate

It's a mass-produced compound (Huntsman has a doc) used in lithium batteries.


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#56 KOBTstronginteraction

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Posted 6 October 2013 - 01:10 AM

Are we talking about creating the element sodium? Or creating NaCl?


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#57 John Cuthber

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Posted 6 October 2013 - 09:28 AM

Are we talking about creating the element sodium? Or creating NaCl?

Have you read the title of the thread?


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What's this signature thingy then? Did you know Santa only brings presents to people who click the + sign? -->

#58 chemadict

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Posted 27 October 2013 - 03:30 PM

Well, you can buy it, lol


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#59 mns

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Posted 6 August 2014 - 05:47 PM

It appears that sodium metal can be obtained from the anhydrous reaction of sodium hydroxide and magnesium. The popular video of this synthesis on youtube is pretty irresponsible, though. The video in question shows the combination of powdered sodium hydroxide and magnesium, which leads to an explosive reaction after ignition with a fuse. If too much material is used, this could be very dangerous. Powdered solids react very quickly due to their high surface area, and should be avoided.

 

The reaction is likely:

 

2Mg(s) + 2NaOH(s) -> 2MgO + H2(g) +2Na

 

While metallic sodium cannot be electrolytically produced in aqueous conditions, it can in the absence of water. I have not been able to find standard reduction potentials for sodium or MgO in anhydrous conditions, but the demo video I mention above seems to show that sodium has been produced, indicating the redox potential must be positive. If you use standard reduction potentials tabulated for aqueous conditions, a negative redox potential is obtained, but it is not directly relevant to the anhydrous reaction above.

 

The sodium metal produced is really not that dangerous compared to burning magnesium. Burning magnesium is the much bigger hazzard here, since it burns at very high temperature (hot enough to weld steel), can't be extinguished with water, and produces a bright enough flare to cause blindness. You should never look directly at burning magnesium.

 

The sodium slag produced in the video should not be recovered in mineral oil over water due to the combustion hazard. Instead, the metal and slag can be separated in a pyrex beaker by heating to around 100 degrees C when mixed with mineral oil.Note that a hotplate should be used, not an open flame. Mineral oil can burn. The sodium is very low melting, and will pool at the bottom of the mineral oil, with the magnesium and sodium oxides rising to the top of the metal pool. The  metal oxides might even float to the top of the mineral oil where they can be skimmed off. This is a relatively safe procedure since the molten sodium is protected from moisture by the mineral oil. This is a common technique for collecting sodium into one large plug from the small oxide encrusted pieces that tend to accumulate in a sodium reagent bottle after lots of use. I employed this technique quite a bit in grad school.

 

Mercury should not be used at all. Mercury is a potent and persistent neurotoxin that is hard to clean up. If you ever do use mercury, make sure that you have elemental sulfur on hand to scavenge any spills. Mercury is not easily separated from sodium. The two metals make an amalgam, which is a mercury alloy.

 

Last, there's really no reason to do this reaction. It's very dirty, as you can see in the video. If you reaclly want some sodium, you can buy it on amazon (http://tinyurl.com/ow5cyww), although it is very expensive.

 

Keep in mind, though, that hobby chemistry is a pretty dangerous pastime. This kind of activity needs to be conducted with proper safety equipment. At the very least, it should be done outside, away from flammable materials, using impact resistant polycarbonate safety glasses and with a working CO2 or chemical retardant fire extinguisher at the ready. A proper chemistry lab has chemical resistant work surfaces, extinguishers, fire sprinklers, fume hoods, emergency respirators, safety glasses, blast shields, lab coats, heavy gloves, etc. Even with all of that safety infrastructure, a chemistry lab can still be a very dangerous place.

 

In my grad school research group, a grad student working in the 1970's had been blinded by shrapnel produced by the explosion of a perchlorate salt. The student in question had been wearing safety goggles. The glass shrapnel penetrated the glasses and took out both eyes. As a result, my research advisor had banned use of perchlorate salts. Many near misses occurred when I was in school. I remember a student had heated a flask of ether that he held in his hand with a heat gun (produced a blast of air hot enough to melt lead). This was a very dumb thing to do since the autoignition temperature of ethyl ether is only 190 degrees C, and this was a grad student at UC Berkeley, one of the top ten chemistry grad schools in the world. Even brilliant chemists can do very stupid things.  The super heated ether ignited, the student panicked, dropping the flask. The flask shattered and the ether spread to cover the floor of the entire lab with fire. Luckily the ether was quickly consumed, but not before the sprinklers went off. I once dropped a bottle of highly reactive phosphorous compound that filled the lab with toxic green smoke. We had to evacuate, and when I got to the hallway, I realized that the phosphorous compound had melted almost completely through my heavy leather work boots. I had to throw the smoking boot back in the lab. If we didn't have proper ventilation, we might not have survived. I was once sitting outside taking a break from lab work, and watching birds fly over Latimer Hall. It was at that time, a burst of opaque, brilliant yellow smoke issued forth from one of the exhaust stacks. The smoke had been pulled into a fume hood somewhere in the building, and had been too much for the air scrubbers to absorb before being exhausted out of the roof. A seagull flew through the cloud of yellow smoke, and tumbled down to the ground. I ran over to the bird to find it was dead and had an overwhelming odor of geraniums and sulfur. I have no idea what the chemicals involved were, but I was coughing and nauseous the rest of the day.

 

Without all of the safety equipment and industrial hygiene experts at Cal, the death toll to the chemistry grad students would be pretty high. Don't do this stuff without the safety gear and training. Blowing stuff up is not worth the extreme risk to life and limb. If you really want to blow stuff up or play with fire, learn how to do it right. Many metropolitan areas will have organizations that will teach fire arts. For instance, in the San Francisco bay area, there's "The Crucible". 


Edited by mns, 6 August 2014 - 05:54 PM.

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#60 johnsonchester03

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Posted 19 January 2016 - 08:49 AM

Thanks for this post. This is one of my question.


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