Make Sodium Metal Without Electrolysis Using Domestic Chemicals

Make Sodium Metal Without Electrolysis Using Domestic Chemicals

Warning: Sodium is a highly reactive and flammable metal. This reaction also uses high temperatures and generates flames. Fire safety protocols must be in place. Dioxane is carcinogenic. Wear gloves when handling it and work outside or in a fume hood. Greetings fellow nerds. Sodium metal is quite famous among chemists as well the general public for being flammable and highly reactive. You’ve seen me use it in previous videos to dry solvents. In another video i showed how to make a sodium magnesium oxide aggregate that could be used as a sodium substitute in many cases. Now in recent years getting sodium metal has become much easier. Online sellers can be sought out who will now sell sodium to individuals. But making actual sodium metal without electrolysis remains one of the holy grails of amateur chemistry, until now. First, we need some sodium magnesium oxide aggregate. I already showed how to do this in a previous video so i won’t go over it in detail. Please check the link in the video description for how this is made. But briefly, we react a sodium hydroxide with a stoichiometric excess of magnesium metal in an air free environment. It is here where we are actually making sodium metal. Upon cooling this produces sodium magnesium oxide aggregate that we chisel out. This is actually the easy part though. The much harder part is separating the sodium in the aggregate from the magnesium oxide slag. To do this, take this aggregate and grind or blend it into a powder. Doesn’t need to be too fine, but the pieces should be small enough that solvent can thoroughly seep in. The consistency of coarse sand is acceptable. And here it is after being poured into an erlenmeyer flask. Now for the secret ingredient, add about 2 to 3 times its volume of 1,4-dioxane which i showed how to make in a previous video. I’m using purified dioxane but you can use dioxane that was dried over sodium hydroxide. The small amount of moisture and contaminants will decrease yield slightly, but the process will still work. Now drop in a stir bar and set up a distillation apparatus around the flask. Turn on the stirring, the cooling water and the heating and distill off the dioxane. Keep the stirring high to break apart the aggregate and liberate the sodium metal. At first the sodium will be a fine droplets but as you continue distilling off the dioxane the sodium is forced together and will coalesce into larger spheres. Now i know another youtuber nighthawkinlight showed how to separate the sodium by dropping it in mineral oil floating on water. While this does work to some extent the yields are very low and the sodium produced is not nearly as pure or as high quality as the method i’m showing here. Now some people have suggested to distill the sodium out of the aggregate. You can try but considering sodium has a boiling point of 883 degrees celsius the process is not exactly trivial. You need to make your own steel and copper distillation apparatus and working with those temperatures isn’t easy even for a professional lab. The great thing about this dioxane method is that the temperature needed is just 101 degrees celsius, which is easy to handle for the amatuer with proper ground glass distillation equipment. 101 degrees celsius is much easier to work with than 883 degrees celsius. Anyway, as you progress and continue distilling off the dioxane the sodium will start pushing up against the surface of the dioxane as it runs out of room. Those bubbles you’re seeing here aren’t actually bubbles. They’re spheres of molten sodium. I must admit when i first saw this and realized what it was i was ecstatic. Getting sodium without electrolysis has been on my personal achievement list. Okay, now at this point slow down stirring to a minimum but don’t stop. We need the stirring reduced to gently push the sodium together and encourage coalescence, but not strong enough to break apart the globules again. The original job of the stirring to break apart the aggregate and release the sodium is complete. Now keep distilling the dioxane all the way to dryness. You can see a huge globule of sodium in the back there as it picks up many of the smaller sodium globules. Anyway, as the slag of magnesium oxide gets drier and drier turn off the stirring completely. This is so the stir bar doesn’t start breaking the globules now that there is no solvent to to lubricate its movement. Now that globule in the back is probably too big to fit through the neck of the flask. If you want you can willfully restart stirring to break up large globules into smaller globules. I’m not going to do this though since the sizes are unpredictable. I’d rather have them in large sizes that i can manually cut down later. But the choice is yours. And here it is after all the solvent has been distilled off. The sodium looks perfect right now because it’s molten and because the flask is currently full of dioxane vapor and very little air. So the surface hasn’t been oxidized yet. Now you might be wondering why i’m using dioxane specifically for the sodium separation step, and would more common solvents like toluene, xylene or mineral oil work. For some reason dioxane works best for this. I and many other amateurs have tried all kinds of solvents over the years and we could only get minor quantities of sodium out. Not enough to be useful. Here i am repeating the experiment with xylene. As you can see, there was almost no separation. Maybe a few tiny globules hitting the glass but nothing recoverable. Now it might seem that dioxane’s higher density to sodium is the key, allowing the sodium to float out. While I do agree this is tremendously beneficial, i don’t think that alone is what allows separation. I and many other amateurs have tried other high density solvents like tetralin, napthalene, biphenyl, methyl napthalene and so on. But none of them afforded separation this good. Anyway turn off the heating and let the sodium solidify. And here we are with the sodium at room temperature. Now slip in a metal strip or stirring rod and dislodge the sodium nuggets. The reason why I let the distillation go dry is that it makes the slag of magnesium oxide cake together instead of remaining as a powder. This makes it much easier to separate and only a small amount of the slag will come out when we pour out the sodium. Once the sodium is out, pour it out of the flask. Looks i was right and the large sodium nugget was too big to fit through the neck. I’m not disappointed though. Better to have a successful reaction problem than a total reaction failure. If this happens to you just slip in a metal knife or strip and cut the sodium in half or smaller. Let me pull the sodium out and there we go. Dump out the pieces. Now it’s just a matter of picking them out. These big pieces are obvious. But if you’re having trouble telling the sodium from the slag then squeeze them with your tweezers. Sodium is soft and deforms easily like clay. But the magnesium oxide is hard like a rock and if it does give it will break instead of deforming. And there we have it, sodium metal. It’s already starting to oxidize on air and becoming covered in white sodium oxide. Now some of you might be asking instead of distilling would it be preferable to reflux the dioxane and the aggregate and separate the sodium that way. I tried it and it does work but i didn’t like the overall process. The sodium doesn’t float in dioxane even though it has lower density because it picks up a lot of slag as it cools. Additionally when you pour it out you have to deal with the full brunt of the dioxane vapor and then you have to manually pick out the sodium. Here it is. It looks like i got better yield but this is actually the same amount of sodium as my other run. It’s just spread out in smaller pieces. As you can see the pieces are of much lower quality with considerable magnesium oxide slag sticking to the surface. So on the whole i prefer distilling to dryness and picking up the larger pieces. Having the slag cake together like this is much more convenient for me. But you can use refluxing if you want. Okay, once you’ve picked out the sodium i recommend putting dioxane back onto the slag and running the distillation one more time to squeeze out the last remaining bits of sodium. This is optional though since it’s not a lot. You can actually recycle the dioxane you just distilled so if you’re very careful you can recycled the same dioxane for a great many runs of sodium. Now after recovering the sodium you’ll need to neutralize the slag now that it’s caked into the flask. Get a tall container of water and push the flask of slag down into it with a rod. I don’t recommend using your hand directly in case there are leftover exploding bits of sodium. We’re doing this to destroy the sodium so the slag is safe to toss out in the trash. Last thing you want is your trash on fire. We’re using this much water to absorb the large amount of heat generated by the reaction. Leave it in there for a few hours to make sure all the slag is destroyed. Okay now back to the sodium. Now there are still bits of magnesium oxide slag sticking to them and some of the spheres are rather small. We can remove the slag as well as coalesce the smaller pieces although this is not strictly necessary for most purposes like solvent drying. Anyway, to clean and coalesce them, get some in a vial or beaker and cover them in mineral oil. Then add in a small amount of isopropyl alcohol, a few drops to start. Now heat up the mixture until the sodium melts. As the mixture heats up it will bubble as the sodium reacts with the alcohol. The surface coating of sodium oxide and magnesium oxides will also react to produce various alkoxides that slough off the sodium. As the sodium melts they’ll pull together in spheres. They should get shiny but if the bubbling dies down and the sodium still dull then add in another shot of alcohol or manually stir it with a stir rod. It helps to turn on the hotplate stirrer even if there is no stir bar. The paramagnetism and eddy currents in the sodium will cause the spheres to spin and help to shake off the slag. You can also poke the spheres with a rod and coax them to coalesce. You can use any alcohol and i actually tried a few different ones. The best ones i found for making the cleanest spheres were tertiary alcohols like t-butanol and t-amyl alcohol. Lighter primary alcohols like ethanol still work but you need to stir them manually a lot more. Once the sodium is cleaned and coalesced, turn off the heating and let it cool. And there is our sodium sphere. And there it is, pure sodium metal without electrolysis made with domestically available chemicals. Now store the sodium in mineral oil or other inert solvent. You can even use the heptane we obtained from starter fluid in an earlier video. Otherwise the sodium will slowly react with air and get covered in sodium oxides and hydroxides. If you store it for too long they’ll be destroyed completely. Now if you don’t want large spheres and actually want small spheres of sodium then the cleaning process is much simpler. First scrape off as much of the slag as possible by hand and then boil the sodium in dioxane. You can even stir the mixture to encourage break up of the sodium. This will produces smaller spheres of sodium. They’re are useful for reactions where you need large surface area, but aren’t good for long term storage since larger surface area means they also decay faster as they react with air and moisture. Anyway, i’m going to cleanup and coalesce the rest of my sodium. As mentioned before this is optional for most purposes. You can simply cut off the slag on the larger pieces if you really need fresh clean sodium. For purposes like drying solvents the slag has no effect. I’m going to clean all my sodium as i want a nice trophy of my accomplishment. And here is all my sodium from 3 different runs stored under mineral oil. That strange texture on the surface is actually crystallization lines. To do that indicates the sodium is exceptionally pure. You might be wondering what the yield is. For an initial reaction of 40g of sodium hydroxide and 30g of magnesium metal i was able to obtain a final recovered yield of 9.5g or about 41% based on sodium hydroxide. While some of the loss is from irrecoverably small sodium particles in the slag, i think most of the loss is from the initial production method. Air can get in and destroy the sodium during the initial cool down and i’m sure a lot of air destroyed the sodium during the grinding step. I didn’t cover it in dioxane since my blender is plastic and would have been destroyed. But I’m not too worried about the overall performance. I think being able to get 41% yield out of an air sensitive reaction done in a soup can is in itself amazing. This method gives the amateur chemist access to sodium even more easily and quickly than electrolysis of molten sodium hydroxide. I was able to run the whole process in less than a day. And all the chemicals including dioxane can be found or made from domestically available sources. The dioxane is made from antifreeze, the sodium hydroxide is from drain cleaner and the magnesium can be obtained from fire starters. So on the whole, sodium is now accessible to the amateur chemist. Now some might ask if you can use dioxane for the catalyzed magnesium reduction method of making alkali metals like potassium. I tried doing it with sodium and it didn’t work. But then again this reaction is notorious for being difficult to reproduce so my failures may not be indicative of actual plausibility. Anyway, now that we’re on the topic of other metals like potassium. Could the high temperature reduction method also work for potassium and other alkali metals? Yes it would, and i’ll make a video on those metals as well. But for now, i hope you enjoyed my video on how to make sodium without electrolysis using domestically available chemicals. Thanks for watching. Special thank you to all of my supporters on patreon for making these science videos possible with their donations and their direction. If you are not currently a patron, but like to support the continued production of science videos like this one, then check out my patreon page here or in the video description. I really appreciate any and all support.

100 thoughts on “Make Sodium Metal Without Electrolysis Using Domestic Chemicals

  1. I was just a little bit curious after seeing a different video somewhere, thought I would skip… but it's so interesting, I ended up both watching the full thing and I subscribed!!

    … "Last thing you want is your trash on fire"… Classic! 🙂

  2. NICE!!!! holly smokes this redefines home chemistry! :)) I don't see why you couldn't use the aggregate as is to make sodium ethoxide , I think the magnesium would be inert, not to sure on separation, probably not to hard with acetone

  3. I would like some of your video because I live in the Yukon in summer times and in the jungle of the Philippines in winter. I am a gold digger 90% of time I don't have access to chemical. I see that you make some of the reagent that I need to purify my gold and some extraction of platinum. It is hard for me to be online because of my love to nature. Could it be possible to have some of your video download so I could do your I could had them to my chemistry files on my laptop. my facebook Jean Claude Emond

  4. hi @NurdRage ive been thinking in using Aluminum powder instead of Magnesium powder, would it work the same way as the Magnesium ?

  5. Hey NurdRage, i also gived this Setup some try, BUT my 1.4 Dioxane is just cooking away without seperating the sodium from the slag… om my Setup i got more luck with hot mineral oil ~150°C, adding some drops of Isopropyl alcohol gently periodically.

  6. Love your video, holy grail it is indeed. I have question though, does aluminum instead of magnesium works also with this reaction??

  7. Is it possible to use mineral oil instead of dioxane. mineral oil has slightly more density that sodium and it is pretty heat resistant. Dioxane is VERY toxic and flammable do could mineral oil be an option?
    I haven't tested it myself because I am still researching on more less toxic chemicals to use so I don't put myself into danger.

  8. Pretty sure this has something to do with a chelating effect where the dioxane ring has the sodium atom in it's center; or maybe an octahedral complex with three dioxanes around it. Interesting!

  9. I think there is another reason for the low yield, which is sodoun could form vapor at the first step. The evidence is the fire which was supposed to be hydrogen flame was yellow caused by the combustion of sodium, while the real hydrogen flame is blue.

  10. Would adding a small amount of alcohol to the dioxane in the distillation step help separate the sodium from the magnesium oxide?

  11. WOW!!.. I have to say you're Amazing!! As a Person! Not to mention the Knowledge and Experience you obviously demonstrate! But For the most IMPORTANT!! Part you give your knowledge so freely, THANK YOU! I have only ever Subrcibed to 2 Channels in all my years on youtube. You will be #3. and #1 was a channel for my kids. Just Well done and I wish you all the luck in your life.

  12. Could one expect even better results with some solvent heavier than dioxane and with higher BP, especially for method with no metallothermy involved? 1,3,5-trioxane maybe? Or some other long-chain ether. Could crown ethers fit?

  13. Nice Video man, but isn't distilling dioxane extremely dangerous bcause it can lead to the formation of peroxides which can explode?

  14. This is such fascinating and skilled work! I've featured this process in my podcast about sodium on the Episodic Table of Elements, and included your video in the show notes at . Thank you for sharing this amazing educational resource!

  15. That was excellent! And to YouTube I say, leave this channel alone. It's educational and it's an awesome channel!

  16. It's so weird to me that the dioxane is completely unreactive with sodium. That no part of the dioxane molecule, the sodium atom will ever attack. Because water, burnt hydrogen, is presumably fairly stable, since it releases so much energy when you burn hydrogen, you'd think it would be hard to break apart that molecule and that anything that could, would do something to some part of a large, presumably more fragile molecule, especially one which also contains hydrogen AND oxygen, and is definitely a fragile enough molecule far enough from its energy rest state to be highly flammable. So why is water so vulnerable to sodium but dioxane completely immune? Is it just because there are no bonds between hydrogen and oxygen, and all the hydrogen and oxygen atoms are bonded directly to carbon? I'd still expect some sort of interaction to happen, especially since it's a flammable liquid. Just by being very flammable should make it reactive, but then I already know sodium doesn't react with the alkanes and that's why they store it in mineral oil, but that's kind of weird too, even if it doesn't contain oxygen in its molecule.

  17. For those of us without a professional lab hotplate stirrer or $500 to spend on a hotplate, can we use a regular cooking hot plate with some stirring with a stirring rod instead for distilling dioxane? Would a better option be making a crude homemade hotplate stirrer like in here: or buying a cheap hotplate stirrer from ebay

  18. i dont think i can make sodium metal D:
    too much chemicals that i cant get at my highschool age… but i can make potassium 😀

  19. In this video you say "Make sodium metal without electrolysis using domestic chemicals" way too many times… but you deserve it!!! 🙂

  20. Oxygen-carbon bond length=112.8pm, multiplied by two because of the two bonds=225.6pm… strangely close to the radius of a sodium metal atom; 227pm (values found on wiki). Maybe the dioxane is working as a cage like chelate? I wonder what would happen if you used 1,3,5-trioxane……or 1,4-dimethyl-piperazine

  21. I went ahead and gave this method a try today and was able to make 1 gram of sodium metal! This was the first time I've ever actually been able to make sodium on my own; all the other methods I've tried over the years have failed. So I'm absolutely thrilled to have finally been able to make sodium for the first time! There's just something magical about a metal that's soft as butter and when reacts with water melts under its own heat to a perfect sphere, and produces hydrogen that surrounds itself and immediately combusts into a wizzing fireball that just like that disappears!

  22. The oxygen in the dioxane ring will not oxidize the sodium, but has just enough pull to free it from hanging around the magnesium oxides. Loved the use of a sparkler to make the oxide too.

  23. I think you can use THF or toluene instead of dioxane, also the main idea depending on different melting points between Na and Mg

  24. what happened to the excess magnesium metal? does it slough off with the oxides when you do the cleaning step?

  25. I was gonna say a joke about sodium, but I was like Na the people here are probably too dumb, I'm super smart and they wouldn't even get it. I have the IQ of like 120. So if you want to hear it, then just ask, but you wont understand it lol

  26. Nurdrage. Would boiling the aggregate in mineral oil work? Not as good as dioxane but still better that the water/mineral oil process. Would be great if you can reply. Thanks.

  27. y not just set the sodium hydroxide and magnesium on fire, then just put it in mineral oil and clean it with the alcohols?

  28. Waiiiiit a minute… You're an AMATEUR chemist?… Well thats pretty cool. I thought you were a damn pro! Makes me feel like i might be able to get somewhere by just reading books and forums. 🙂
    Would really like to learn more about you, out of curiosity.

  29. Nice Video documented Sodium metal synthesis fromout a chemical reaction possible to make it at home – Using Magnesium and Sodiumhydroxide. 😀
    It is a nice Method to be tried by Hobby Chemists.
    I thought about if that Sodium metal produced like shown here is pure enough to construct or beeing used in a sodium vapour lamp – a light source often beeing used in spectroscopy or optics to get the sodium D Line. 😀
    Such a lamp was first used by german Physicists in the 1920´s like Max Born (1882-1970) or Friedrich Hund (1896-2002) in the famous work "Optics", first published in 1933.

    I abonnated your channel – it´s quite nice to watch some videos like your in now a days…:-) Greetings, Silvia

  30. Thank you, doctor, teacher, and friend, for your very easy to follow, educating and ultimately inspiring video lessons. I look forward to every video u post, and often watch my favorites multiple times.

  31. So you’re an amateur chemist with a PhD in chemistry? A PhD degree in chemistry is kinda the definition of “professional”, isn’t it?

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