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No, you CAN'T make sodium!


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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.jp/lab/wp-content/uploads/2011/07/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.uchile.cl/bitstream/handle/2250/7121/Florez_Elizabeth.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.com/article/10.1007%2FBF00767207#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/mtrl-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."


"In summary, water demonstrably chemisorbs onto
MgO but trustworthy calculations show that H2O molecules
should not dissociate on the only known stable

Also, see: http://www.chem.tamu.edu/rgroup/goodman/pdf%20files/343_jcpb_103_99_3391_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.com/books?id=I0SCzBkz7IoC&pg=PA213&lpg=PA213&dq=reduction+of+CH3OH+by+MgO&source=bl&ots=bhoppFUF9m&sig=oFkDYcYkIe2kAlmz8Q8qFJy3IMk&hl=en&sa=X&ei=vKl6UdHiAsfj4AO80YHYCA&ved=0CDwQ6AEwAw#v=onepage&q=reduction%20of%20CH3OH%20by%20MgO&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)


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


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
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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
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Yes you actually can. If you mix Sodium Hydroxide with Magnesium or Aluminum powder and ignite it there will be a thermite. If you pour the slag into a container of mineral oil and water the sodium goes to the top...

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I have a question. Sodium bicarbonate is slightly soluble in aceton. It is more soluble in methanol but that will blind you. Would a sodium bicarbonate in acetone solution be electricly conductive? It might be a dumb question, but I guess it's worth asking

Btw Nile Red has a nice vid about how alkaline metals can be synthesized, but people who are trying to do it with salt and a 9V battery probably won't have a 3rd tier alcohol laying around xD

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It seems it is a sad factor for mee...Coz i like BOOOOMS and ya know Na is very expensive in Sri Lanka.... Ohhhh

Yes you actually can. I dont know why this is such a large thread when you CAN make sodium. You can make a NaOH thermite and it will yield sodium in the slag.

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On 29/10/2009 at 4:01 AM, hermanntrude said:



On 22/4/2009 at 6:28 AM, hermanntrude said:

Sodium hydroxide was another suggestion I read on the forums, and while it certainly has a lower melting point, we're talking about a strong base here... do you REALLY want a molten puddle of fuming, spitting sodium hydroxide in front of you? REALLY? No. You don't.


Well I was actually trying to do this and it worked, but I would have done better with carbon electrodes (I used a tin can, Cables, Graphite electrodes, etc. But I only made 3 grams, not much for the time (5 hours) I spent planning and making the practice.

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I thing think that in 320 C H2O (in electrolysis of NaOH; this is how sodium was made for first time) will react realy fast with Na creating e protective layer of H2. The same as reaction of copper with 70% and 50% nitric acid.

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Actually, sodium is fairly easily prepared by electrolysis.  Commercially, it used to be made by electrolyzing using a mercury-metal negative electrode, which collected the sodium formed in an amalgam.  ("amalgam" is the term for an alloy where one component is mercury metal).   The mercury is then  pumped away, and is then distilled off, leaving the sodium metal behind.

Sodium can also be prepared by electrolying a molten salt, sodium chloride might do, although this would involve very high temperatures to melt the salt.


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On 11/12/2019 at 7:11 AM, jim bell said:

Sodium can also be prepared by electrolying a molten salt, sodium chloride might do, although this would involve very high temperatures to melt the salt.


Probably it's possible to use wind energy to heat(using Faucault currents, so inconstant wind power isn't a problem) the salt. The very good news is that this molten salt has a great heat of fusion(around 250-300 kwh/m3) value so it can work even if wind is not constant.

Edited by harlock
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I completely agree with the author of the topic. It is not enough for the substance to react with water. High-precision laboratory research is required, which is practically impossible in the case of amateur chemists. And it turns out that almost everyone can become Tony Stark from the Mavel universe

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well you could make sodium without electrolysis by using magnesium and sodium hydroxide with a menthol catalyst


2Mg + 2NaOH ----Menthol, 200C--> 2MgO + 2 Na + H2

to make sure nothing explodes keep apparatus airtight and use mineral oil as a solvent ig, also keep bubbler for H2 and don't use water because it'll probably backflow

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Actually, you can make sodium, but it's important to exercise extreme caution when doing so, as sodium is a highly reactive metal. Sodium is not typically produced by individuals in a home setting due to its extreme reactivity with moisture and air. However, it is manufactured on an industrial scale.

The most common method for producing sodium involves the electrolysis of molten sodium chloride (table salt) in a process known as the Downs cell. This method is used in industrial settings to produce sodium for various applications, such as in the production of chemicals and metals.

If you were to attempt to produce sodium at home, it would be very dangerous and should not be done without proper training and safety measures in place. The reaction of sodium with water is highly exothermic and can lead to fires or explosions, so it is not recommended for amateur chemists or individuals without the necessary knowledge and equipment.

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