Mix the MgCl2 and KCl togther in a ball mill. Add a small amount of silica sand to it as well. Melt down. Continue heating to bright red heat in a steel can. Next use a graphite rod as the anode. 7 to 15A is about right for a soup can container. The magnesium will sink to the bottom of the can. The high temperature also prevents formation of an alloy of k and mg.❤
Hi Scrap Science, Nice video and encouraging results. Magnesium dust or sponge can be quite reactive with air or water. Dust can be pyrophoric and may lead to a FAE while opening the drum. Jus like Al, with water, Mg reacts and it can be accelerated by warm water (sometimes due to self warming). Mg + H2O --> H2(g) + MgO(s) MgO + H2O --> Mg(OH)2(s) Some water gels of Al (and probably Mg too) can be actuated by primers and H2O2 mixes are devastating. Magnesium rods are used to reach very high temperature underwater because it reacts with water to produce MgO that burns hot and bright underwater generating H2 gas arround. When trowing the molten metal and MgCl2 onto rusted iron...it may decompose a bit the Mg metal due to "magneto-thermic" reaction, Mg is very close to aluminium and it is thus reducing metal but more reactive than Al. 3Mg + Fe2O3 --> 3MgO + 2Fe Reagards, PHZ (PHILOU Zrealone from the Science Madness forum)
@6:15 perhaps you could use a small fine-meshed steel sieve for the electrode put deeply inside the molten salts (and saving it from oxygen and CO2 and CO, that may have reduced the yield ?). Could you coalesce the metal chunks under a borax layer ?
Not a super long time ago, I found out that there are rocks in my area that are extremely magnesium rich, which I hope to eventually separate from the rocks and convert into custom castable alloys.
I think one of the problems is, once the molten metal falls, it's then falling on the anode and thus is in an oxidizing environment. Maybe put an insulator, like a silicate dish, on the bottom of the graphite crucible to collect the liquid metal?
Should have blendered the oxides and such after removing from salts, the heated to 7-800c with a bunch of glass or boric acid etc for flux, and metal should reform?
8A on such a small wire is probably close if not more than the limiting current, which would also explain that your metal is reactive .... if you namely go beyond limiting current for your magnesium reaction you will get potassium deposition or alloy deposition Also, I think it wont work to do molten magnesium in your cruicible and the cruicible being the anode .... as the magnesium would drop down and then reaoxidize on the anode For that the crucible should be the cathode and another graphite piece(/rod) as the anode
I'm pretty confident that we're not dealing with any potassium metal here. If we've made potassium under these conditions, even in an alloy, it goes against everything potassium tends to do in high temperature chloride melts. It is extremely difficult to deposit potassium from any molten salt, so this would be quite a feat. But yes, the crucible being the anode was not ideal here - I wasn't expecting the metal to sink. I address a couple of ways to get around this at around 12:12.
I have a crazy idea I'd like to try: I think it should be possible to dissolve MgCl2 in acetone or ethanol (less so than in water but still) and you should be able to extract the Mg from there by electrolysis. For a reason I am unable to explain for having forgotten it, the electrolysis doesn't work in water but could work in these other solvents. It is hugely recommended against because the electrolysis will also produce tons of weird organic compounds and everything might catch on fire. But it should be a pretty fun experiment.
No it does not work, but it does do something a lot more interesting if you are using an alcohol. It makes organic magnesium halides aka griginard reagents. ❤
7:10 could you maybe have a bit of potassium contaminating your magnesium that would explain the tame but existing reaction of your product with water?
Ohh, that makes me wonder if it would be useful for thermochemical dioxane process to make sodium, especially since it reacts with water, so could also do away with the starter bit of sodium/lithium as a drying agent.
It seems that ultrasonication (about 25 years ago) was used into the department nextdoor (physical chemistry and corona chemistry) where I was making my endwork study - report (photochemistry)) can be used so that cavitation bubbles exploding onto the surface of Mg metal makes its apparent surface so big that the metal can then be used as such for Grignards without drying, ether nor dioxane. I know that ultrafine Al dendritic dust (made from exploding Al wire under high amperage and vacuum, then exposed to slow flow of cold air so softly passivated) can be compressed into a spongeous form. This dust has a much higher specific impulse for rocketery than conventionnal german black and it makes a sponge that can absorb TNT forming a core for implo-explosion-detonation; the nanometric Al dust generated into the TNT detonation makes a FAE that suck up all O2 and N2. PHZ (PHILOU Zrealone from the Science Madness forum)
Could it be that what you see reacting with water is in fact a tiny percentage of metallic Potassium that you managed to reduce in the same electrolysis bath?
It's possible, but it would be surprising. Potassium is notoriously difficult to deposit from its molten salts, especially chloride melts, since the reaction is thermodynamically unfavourable and the metal is generally soluble in the melt itself. If we had in fact made a magnesium-potassium alloy, I'd also have expected it to give a slight purple flame when burning it - which we didn't see. Again, it's possible, but the evidence seems to suggest that it's unlikely.
You could melt the sponge stuff in a non-reactive atmosphere. Of course the usual option is argon but it would be interesting to see a fuel like hydrogen to be used since it doesn’t oxidize.
That you got anything at all is impressive. Can you imagine being in the time of Davy, and relying on nothing but an old battery that's been used for lectures and experiments for years?
I'm honestly not sure. Generally, mixing chloride and hydroxide ions in a melt you want to electrolyse is probably a bad idea. The oxidation products can end up generating chlorate and other oxyhalogen species, but this may be at a high enough temperature that it's not a concern. Possibly worth a try.
When anhydrous, magnesium sulfate will decompose before melting. Even if it didn't, the anode reaction would be rather complex when working with molten sulfates. I don't know exactly what would happen, but I doubt it would be well-behaved...
I bought a magnesium fire starter. I played with the metal, filed it, played with powdered and flaked (burned and such). The main block ended up in a cup of water, can't remember what I was doing, but anyhoo, it became covered in this white stuff, looking like an oxide layer on steroids, but the black and whiteness of it let's me know that spongy stuff is all magnesium. It is very reactive, basically with more surface area its almost like activated aluminum (like mercury treated). A simple fire test proves this, as you see with the beads. You have far more magnesium than those little beads. That Slag is pretty pure magnesium.
If the metal sinks off the cathode as it forms, it will sink onto the anode where it will inhibit the oxidation of chloride by oxidising itself away. Assuming you don’t put some sort of insulating vessel inside the molten salt, or maybe some sort of much denser chloride salt that the magnesium metal will float atop. Best off performing the electrolysis at a lower temperature and attempting to remelt it separately. Didn’t think about using the graphite as the cathode, that’s sensible. Bet that dendritic sponge burns well though. Maybe you could dissolve the salts off in particularly high pH water?
I was kind of surprised to see that the sponge doesn't even burn. The salts within the sponge completely coat the stuff when you heat it up, and I couldn't ever get it to ignite. Dissolving with very basic solution is something I hadn't thought of though. Might work! If I ever try this again, I'll have to give it a go.
I wonder if some potassium fluoride would help the oxychloride mixture dissolve/electrolyze? Also be interesting to use a graphite rod anode in the middle, letting metal form on and sink to the bottom of the crucible. I wonder if carbide formation is relevant here? Mg2C I think isn't very common, or stable, but with direct elemental synthesis, who knows. A welded steel crucible cathode might be mitigation for this (Fe and Mg are pretty much insoluble/immiscible).
I'm definitely no expert on what fluorides would do. It's definitely possible, but I don't have much desire to work with fluorides... And yeah, magnesium carbide is also a possibility. But I'm not really sure on that front either, honestly...
The reversible double carbon reduction reaction against magnesium oxide would be a better way to do this. That or the one way reaction with aluminum. Both require a lot of heat energy, but both are very precise and precisely controllable.
HEY ive been looking for a diy thing for magnesium. I was considering it for a looping oxidation furnace. I eventually decided id rather use lime since calcium cyanamide is a byproduct, aka easy nitrogen fixation.
Wouldn't the black stuff be Magnesium Carbide? It would explain it being reactive with water, producing acetylene. Perhaps there is reaction with the crucible.
It's possible! In that case though, I don't understand why it's still forming on the cathode like it does at 4:30, since that's nowhere near the graphite surface and shouldn't have any carbon to react with.
Yep, a 50:50 mole fraction of MgCl2 and CaCl2 is a eutectic with a melting point of around 620-630 C. I can't guarantee it would work to make magnesium, but it's quite likely. I used a mixture that corresponded as close as I could get it to the low melting point region of the graph I show at 1:17. About 40% MgCl2 by weight.
I think you have made magnesium anhydride, which releases hydrogen when exposed to water. If it’s hydrogen the bubbles should ignite when exposed to a flame.
I can guarantee I'll make at least a couple of videos on the topic of perchlorate synthesis at some point. I'm still on the research stage for that project, but it will happen one day.
Great hearing that. Im also currently looking into said process and i am astounded at how much more complicated and difficult the perchlorate process is. Kind regards from Denmark@@ScrapScience
Another method is Silicon to purify Calcium and Magnesium metal. It could work since it binds to iron, etc if needed. But I assume it is quite pure Magnesium. Yes, Oils are too low for the project Cool video. Patent info on it said Silicon may work. However high temperatures are required. Also works for Magnesium.
That's essentially what I was doing at around 2:50. No matter how you heat it though, there's no way to remove all of the water without some kind of decomposition of the salt occuring (without complex drying techniques, that is).
@@magnuswootton6181 I don't know about it being the cheapest, but magnesium allows are very lightweight, and used extensively in the aerospace industry.
Hi, you need to dehydrate the mgcl2 first (or buy it). Magnesium production from chlorides is all about the dehydration - which is not easy to do. You can’t just boil the water off as you make various magnesium oxide compounds and hcl gas. Once dehydrated you should be able to make a lot of nice metal with your setup.
Correct. I discuss this a little at 1:30-2:00. Due to the difficulty in dehydrating the salt, I opted to try without it. The yield would, of course, be much better by doing this, but it appears unnecessary when we only want a small sample of the metal.
Gotcha. You seemed a little uncertain and I just wanted to mention that it is certainly the difference between what you got and really nice amounts of metal.
No there really isn’t to my knowledge. It’s pretty intense any way you go about it. You can slowly boil it to ~170 c (I think, I would need to verify the exact temp) which gives you the dihydrate form. You will still have problems but it will be much better (that’s how Dow did it at Freeport). You can also buy anhydrous but it’s expensive.
Are you referring to the fact that I spilled the molten salt in the furnace last episode? I completely forgot to mention that... Yeah, I was luckily able to clean it all out and the furnace still works perfectly. :) If you're just referring to the crucible I used in this video, everything easily washed out with a little water.
@@ScrapScience I was referring to the crucible from this video. Making magnesium is very tempting and maybe I will try to replicate your method but I don't want to damage my crucible. It is quite expensive for 60mg of the metal ;). I like your videos because it's like watching pioneer chemists on their way to discover new elements. You don't use highly specialized equipment (although it's getting better over time :) that is available for other chemistry amateurs. Thank you and I wish you a happy New Year :).
Try dissolving MgCl2 and KCl in water and boil the solution down to crystallize carnallite (KMgCl3·6H2O). It should dehydrate easier than the MgCl2 and give better yields. Also the magnesium sponge can potentially be a good source of highly reactive magnesium (for alcohol catalyzed sodium synthesis or similar experiments).
The only trouble with the magnesium sponge is that it's filled with salt from the melt, making it largely unreactive outside of aqueous solutions (which dissolve those salts away, of course). I tried burning it and it wouldn't even ignite, since the salts kept coating and protecting the metal within. Any ideas for removing the salts and leaving the magnesium sponge intact? Maybe it would still work for those kinds of reactions anyway? I'm unsure...
@@ScrapScience Only thing I can think of is putting it inside a blender along with mineral oil and dump the resulting mix in water to get rid of the salts. I would probably directly use it in the sodium production.
@@ScrapScience just an idea what if you put it into coffee grinder; theoretically the chloride shouldnt be a problem anymore ( but on the other hand it can become pyrophoric not sure)
Thanks for this video very well presented. I would have tried melting all the black stuff under argon or nirogen gas to melt the magnesium, that would have incresed your yield of Mg metal.
Yeah, if we could melt the magnesium at temperatures that mineral oil could tolerate, this would be a great method for coalescing the metal. Sadly, magnesium's melting point is way too high.
When anhydrous, magnesium sulfate will decompose before melting. Even if it didn't, the anode reaction would be rather complex when working with molten sulfates. I don't know exactly what would happen, but I doubt it would be well-behaved...
I'm pretty sure one of the experiments in my 80's chemistry set proved that Magnesium in ribbon form does react with water. Think it had to be hot water. This is from memory, I'm trying to hunt out the manual in my terribly organised bookshelf.
Not surprising considering that little bead was around 10mg you are looking at about 250J pf which at sumch small piece and considering the process is adiabatic around 22%-25% of the energy is the form of visible light which means around 60J in light for about four seconds that means about 15W of light on average that is hell bright of light hou know
I didn't explain it very clearly, but that was kind of the idea at the end of the electrolysis run. I tried making a whole bunch of the spongy stuff and then stopped the electrolysis and raised the heat to get it all to melt together. It seems like the sponge was resistant to melting together, but maybe it just needed to get hotter for longer... I don't know.
Bravo.......screw magnesium.......it makes hydrogen .......when u put in water........measurements how much yield......also reactor with gas shield.....cheers
Hi, I've been interested in chlorate production for quite a while now. Finally I found mmo anodes at a reasonable price in the EU. But I have a question and I think you are the person who might know the answer. Question is: what's better? Plate or mesh? Both anodes are in same price. From NurdRange's vid I know that KClO3 crystals can grow in the holes of the mesh blocking electron flow, and that's what I'm afraid of. I guess mesh has a larger surface area so I don't know which one to pick. Hope you know the answer
Crystals growing in the holes is not a huge problem, provided you’re not running your anode at a very high current density. Even with plate anodes, crystals can still easily deposit on the electrode surface. Even the surface area of the plate vs mesh is not significantly different most of the time. All up, it’s pretty much a choice that doesn’t matter. Pick the one you think would look the best I suppose.
*You find the most awkward way to do the simplest jobs. The Graphite Crucible should have been your negative and the Carbon rod be the positive. Minimum temp 1400F. at 3.3 volts and 20 amps. Result: 24 grams MG metal in bottom of Crucible in about 10 minutes. Use Magnesium Oxide and Sodium Fluoride and just keep adding more MgOx as needed. Very similar to Aluminum making process.*
I dunno mate, you always make it very difficult to take your comments seriously, lol. A 1500% error in your original claim doesn't make me very keen to follow your advice...
Potassium is notoriously difficult to deposit from molten salts like this, especially mixtures containing chloride. If we actually deposited potassium alongside our magnesium product, that would be a surprising result.
Are you by any chance making some kind of acid that's accelerating the reaction between water and magnesium? Your reaction produces chlorine gas, and your starting products contained water. Is it possible that you're making some hydrogen chloride somewhere in the process? What's the pH of the solution after dunking the magnesium sludge into it?
Whoops, sorry for the short reply, that was supposed to be a much longer response. Out of interest, what makes you say that there's potassium in the product? Potassium is significantly more difficult to extract as a metal under these conditions.
@ it doesn’t matter the difficulty it matters that it’s possible. If it’s possible then it’s probable it just comes down to what degree. How much potassium. Then also I don’t think you’re accounting for the heat and its thermal reaction being added to the process. Think of hydrochloric acid… it’s not good for dissolving au…. But it is able to attack enough to produce an oxide layer that it can attack slowly, but if you take that same hcl and add heat it increases the rate of ionic exchange.
@ so going by the slow reaction and the fact it’s to h20 given then process used the most obvious thing to think is that you carried over some potential potassium. Until proven otherwise I think that’s the logical conclusion…. But I’m also smart enough to know you’re smarter than me that’s simply a country boys thoughts… the only other thing I can think of is a little sodium metal was made…. But I like the potassium view.
Mix the MgCl2 and KCl togther in a ball mill. Add a small amount of silica sand to it as well. Melt down. Continue heating to bright red heat in a steel can. Next use a graphite rod as the anode. 7 to 15A is about right for a soup can container. The magnesium will sink to the bottom of the can. The high temperature also prevents formation of an alloy of k and mg.❤
Hi Scrap Science,
Nice video and encouraging results.
Magnesium dust or sponge can be quite reactive with air or water.
Dust can be pyrophoric and may lead to a FAE while opening the drum.
Jus like Al, with water, Mg reacts and it can be accelerated by warm water (sometimes due to self warming).
Mg + H2O --> H2(g) + MgO(s)
MgO + H2O --> Mg(OH)2(s)
Some water gels of Al (and probably Mg too) can be actuated by primers and H2O2 mixes are devastating.
Magnesium rods are used to reach very high temperature underwater because it reacts with water to produce MgO that burns hot and bright underwater generating H2 gas arround.
When trowing the molten metal and MgCl2 onto rusted iron...it may decompose a bit the Mg metal due to "magneto-thermic" reaction, Mg is very close to aluminium and it is thus reducing metal but more reactive than Al.
3Mg + Fe2O3 --> 3MgO + 2Fe
Reagards,
PHZ
(PHILOU Zrealone from the Science Madness forum)
@6:15 perhaps you could use a small fine-meshed steel sieve for the electrode put deeply inside the molten salts (and saving it from oxygen and CO2 and CO, that may have reduced the yield ?). Could you coalesce the metal chunks under a borax layer ?
Not a super long time ago, I found out that there are rocks in my area that are extremely magnesium rich, which I hope to eventually separate from the rocks and convert into custom castable alloys.
always a good day when scrap science uploads !!
I think one of the problems is, once the molten metal falls, it's then falling on the anode and thus is in an oxidizing environment. Maybe put an insulator, like a silicate dish, on the bottom of the graphite crucible to collect the liquid metal?
....looking at this I have no idea what I was thinking. Magnesium would react with a silicate dish to pull out the elemental silicon....
Great idea. Happy that you managed to pull it off
Should have blendered the oxides and such after removing from salts, the heated to 7-800c with a bunch of glass or boric acid etc for flux, and metal should reform?
Nice work! Really enjoying your metals series. Some are tougher than others, glad you were successful!
Is it possible to extract the lanthanide metals like lanthanum and neodymium from their compounds using electrolysis?
Yep! We'll hopefully be trying some of those later on in the molten salt electrolysis series!
8A on such a small wire is probably close if not more than the limiting current, which would also explain that your metal is reactive .... if you namely go beyond limiting current for your magnesium reaction you will get potassium deposition or alloy deposition
Also, I think it wont work to do molten magnesium in your cruicible and the cruicible being the anode .... as the magnesium would drop down and then reaoxidize on the anode
For that the crucible should be the cathode and another graphite piece(/rod) as the anode
Can't use graphite for reducing alkali metals, it would eviscerate his crucible in no time. Likely the same for alkali earth metals
I'm pretty confident that we're not dealing with any potassium metal here. If we've made potassium under these conditions, even in an alloy, it goes against everything potassium tends to do in high temperature chloride melts. It is extremely difficult to deposit potassium from any molten salt, so this would be quite a feat.
But yes, the crucible being the anode was not ideal here - I wasn't expecting the metal to sink. I address a couple of ways to get around this at around 12:12.
That is so funny. I literally use the exact same tea spoons for my molten salt/metal experiments. I love it
I have a crazy idea I'd like to try: I think it should be possible to dissolve MgCl2 in acetone or ethanol (less so than in water but still) and you should be able to extract the Mg from there by electrolysis. For a reason I am unable to explain for having forgotten it, the electrolysis doesn't work in water but could work in these other solvents.
It is hugely recommended against because the electrolysis will also produce tons of weird organic compounds and everything might catch on fire.
But it should be a pretty fun experiment.
No it does not work, but it does do something a lot more interesting if you are using an alcohol. It makes organic magnesium halides aka griginard reagents. ❤
@@christopherleubner6633 uuu :)
7:10 could you maybe have a bit of potassium contaminating your magnesium that would explain the tame but existing reaction of your product with water?
I wonder if the spongy high surface stuff would be good to make Grignard reagents?
Ohh, that makes me wonder if it would be useful for thermochemical dioxane process to make sodium, especially since it reacts with water, so could also do away with the starter bit of sodium/lithium as a drying agent.
It seems rather unreactive outside of aqueous conditions since it's filled with the salts from the melt. If we could remove the salts though...
It seems that ultrasonication (about 25 years ago) was used into the department nextdoor (physical chemistry and corona chemistry) where I was making my endwork study - report (photochemistry)) can be used so that cavitation bubbles exploding onto the surface of Mg metal makes its apparent surface so big that the metal can then be used as such for Grignards without drying, ether nor dioxane.
I know that ultrafine Al dendritic dust (made from exploding Al wire under high amperage and vacuum, then exposed to slow flow of cold air so softly passivated) can be compressed into a spongeous form. This dust has a much higher specific impulse for rocketery than conventionnal german black and it makes a sponge that can absorb TNT forming a core for implo-explosion-detonation; the nanometric Al dust generated into the TNT detonation makes a FAE that suck up all O2 and N2.
PHZ
(PHILOU Zrealone from the Science Madness forum)
Could it be that what you see reacting with water is in fact a tiny percentage of metallic Potassium that you managed to reduce in the same electrolysis bath?
It's possible, but it would be surprising. Potassium is notoriously difficult to deposit from its molten salts, especially chloride melts, since the reaction is thermodynamically unfavourable and the metal is generally soluble in the melt itself.
If we had in fact made a magnesium-potassium alloy, I'd also have expected it to give a slight purple flame when burning it - which we didn't see.
Again, it's possible, but the evidence seems to suggest that it's unlikely.
I enjoyed your videos. It is also good to post a video on the production method of ammonium persulfate
You could melt the sponge stuff in a non-reactive atmosphere. Of course the usual option is argon but it would be interesting to see a fuel like hydrogen to be used since it doesn’t oxidize.
That you got anything at all is impressive. Can you imagine being in the time of Davy, and relying on nothing but an old battery that's been used for lectures and experiments for years?
Can i mix it with NaOH to lower melting point or it would get wlectrolyzes first?
I'm honestly not sure. Generally, mixing chloride and hydroxide ions in a melt you want to electrolyse is probably a bad idea. The oxidation products can end up generating chlorate and other oxyhalogen species, but this may be at a high enough temperature that it's not a concern. Possibly worth a try.
@@ScrapScience okay thanks for response
I was waiting for this! Great video!
P.s. Could you please make a video about the production of persulfates?
At some point, yes! It's on my list of future projects.
Small prills are super easy to filter out with using a fine screen.
What about epson salt?, ir si cheaper, can i extract nagnessium from there ando sepárate ir from the sulphurim that😀 can be useful as well?
When anhydrous, magnesium sulfate will decompose before melting. Even if it didn't, the anode reaction would be rather complex when working with molten sulfates. I don't know exactly what would happen, but I doubt it would be well-behaved...
I bought a magnesium fire starter. I played with the metal, filed it, played with powdered and flaked (burned and such). The main block ended up in a cup of water, can't remember what I was doing, but anyhoo, it became covered in this white stuff, looking like an oxide layer on steroids, but the black and whiteness of it let's me know that spongy stuff is all magnesium. It is very reactive, basically with more surface area its almost like activated aluminum (like mercury treated). A simple fire test proves this, as you see with the beads. You have far more magnesium than those little beads. That Slag is pretty pure magnesium.
If the metal sinks off the cathode as it forms, it will sink onto the anode where it will inhibit the oxidation of chloride by oxidising itself away. Assuming you don’t put some sort of insulating vessel inside the molten salt, or maybe some sort of much denser chloride salt that the magnesium metal will float atop. Best off performing the electrolysis at a lower temperature and attempting to remelt it separately.
Didn’t think about using the graphite as the cathode, that’s sensible.
Bet that dendritic sponge burns well though. Maybe you could dissolve the salts off in particularly high pH water?
I was kind of surprised to see that the sponge doesn't even burn. The salts within the sponge completely coat the stuff when you heat it up, and I couldn't ever get it to ignite.
Dissolving with very basic solution is something I hadn't thought of though. Might work! If I ever try this again, I'll have to give it a go.
There was an wet labor elektrolysis with an double salt of ammoniummagnesiumsulfate.No exact descriptions. But writen in an Chemistry book from 1920
Are there any research papers out there for extracting magnesium from magnesium chloride?
I know there has been some work on this.
I wonder if some potassium fluoride would help the oxychloride mixture dissolve/electrolyze?
Also be interesting to use a graphite rod anode in the middle, letting metal form on and sink to the bottom of the crucible.
I wonder if carbide formation is relevant here? Mg2C I think isn't very common, or stable, but with direct elemental synthesis, who knows. A welded steel crucible cathode might be mitigation for this (Fe and Mg are pretty much insoluble/immiscible).
I'm definitely no expert on what fluorides would do. It's definitely possible, but I don't have much desire to work with fluorides...
And yeah, magnesium carbide is also a possibility. But I'm not really sure on that front either, honestly...
Couldn't you put the powder un a centrifuge?
The reversible double carbon reduction reaction against magnesium oxide would be a better way to do this.
That or the one way reaction with aluminum.
Both require a lot of heat energy, but both are very precise and precisely controllable.
HEY ive been looking for a diy thing for magnesium. I was considering it for a looping oxidation furnace. I eventually decided id rather use lime since calcium cyanamide is a byproduct, aka easy nitrogen fixation.
Is that furnace 10A or 15A? :O I have a few too many failed plaster of paris furnaces... ; _ ;
This one is 10 A I believe. One rated for 15 A would require outlets that I don't even have in my house, I'm pretty sure.
OO ty ty! I couldn't quite tell from the IEC plug in the back ;) Now I've got to go price a 10A furnace for myself ;p
Wouldn't the black stuff be Magnesium Carbide? It would explain it being reactive with water, producing acetylene. Perhaps there is reaction with the crucible.
It's possible! In that case though, I don't understand why it's still forming on the cathode like it does at 4:30, since that's nowhere near the graphite surface and shouldn't have any carbon to react with.
What about the chlorine gas this electrical process generates?
What about it?
@@ScrapScience Epic
Will calcium chloride work instead of potassium chloride?
Also, what ratio of magnesium chloride to potassium chloride do you use?
Yep, a 50:50 mole fraction of MgCl2 and CaCl2 is a eutectic with a melting point of around 620-630 C. I can't guarantee it would work to make magnesium, but it's quite likely.
I used a mixture that corresponded as close as I could get it to the low melting point region of the graph I show at 1:17. About 40% MgCl2 by weight.
you will wake up tomorrow
Conflicting messages throughout this comment section already…
You wouldn't know if I didn't.
Unless you don't.
You won't as there is no tomorrow.
You might not wake up tomorrow, carbon-monoxide always wins.
I think you have made magnesium anhydride, which releases hydrogen when exposed to water. If it’s hydrogen the bubbles should ignite when exposed to a flame.
Hi, just wondering is there ever going to be the synthesis of perchlorates proposed by you in the video about sodium chlorate?
I can guarantee I'll make at least a couple of videos on the topic of perchlorate synthesis at some point. I'm still on the research stage for that project, but it will happen one day.
Great hearing that. Im also currently looking into said process and i am astounded at how much more complicated and difficult the perchlorate process is.
Kind regards from Denmark@@ScrapScience
Another method is Silicon to purify Calcium and Magnesium metal. It could work since it binds to iron, etc if needed. But I assume it is quite pure Magnesium. Yes, Oils are too low for the project Cool video. Patent info on it said Silicon may work. However high temperatures are required. Also works for Magnesium.
could you have baked the Magnesium chloride at 400 for a few hours to drive water off?
That's essentially what I was doing at around 2:50. No matter how you heat it though, there's no way to remove all of the water without some kind of decomposition of the salt occuring (without complex drying techniques, that is).
Could you do a video on potassium persulfate
One of these days I'll get around to it. It's on my list of projects to work on, but it's unlikely to come any time soon.
I never seen that anywhere. I think magnesium is a very cool metal.
its the cheapest and most light weight of all metals!
@@magnuswootton6181 I don't know about it being the cheapest, but magnesium allows are very lightweight, and used extensively in the aerospace industry.
Hi, you need to dehydrate the mgcl2 first (or buy it). Magnesium production from chlorides is all about the dehydration - which is not easy to do. You can’t just boil the water off as you make various magnesium oxide compounds and hcl gas. Once dehydrated you should be able to make a lot of nice metal with your setup.
Correct. I discuss this a little at 1:30-2:00. Due to the difficulty in dehydrating the salt, I opted to try without it. The yield would, of course, be much better by doing this, but it appears unnecessary when we only want a small sample of the metal.
Gotcha. You seemed a little uncertain and I just wanted to mention that it is certainly the difference between what you got and really nice amounts of metal.
Do you happen to know of any convenient ways of dehydrating the stuff? I've never seen any that seemed viable in a home setting.
No there really isn’t to my knowledge. It’s pretty intense any way you go about it.
You can slowly boil it to ~170 c (I think, I would need to verify the exact temp) which gives you the dihydrate form. You will still have problems but it will be much better (that’s how Dow did it at Freeport).
You can also buy anhydrous but it’s expensive.
Nice video. First time to see a magnesium extraction. Have you managed to clean your crucible?
Are you referring to the fact that I spilled the molten salt in the furnace last episode? I completely forgot to mention that...
Yeah, I was luckily able to clean it all out and the furnace still works perfectly. :)
If you're just referring to the crucible I used in this video, everything easily washed out with a little water.
@@ScrapScience I was referring to the crucible from this video. Making magnesium is very tempting and maybe I will try to replicate your method but I don't want to damage my crucible. It is quite expensive for 60mg of the metal ;).
I like your videos because it's like watching pioneer chemists on their way to discover new elements. You don't use highly specialized equipment (although it's getting better over time :) that is available for other chemistry amateurs.
Thank you and I wish you a happy New Year :).
I dunno that fine grain magnesium might make for an interesting thermite reaction. Just saying
Try dissolving MgCl2 and KCl in water and boil the solution down to crystallize carnallite (KMgCl3·6H2O). It should dehydrate easier than the MgCl2 and give better yields. Also the magnesium sponge can potentially be a good source of highly reactive magnesium (for alcohol catalyzed sodium synthesis or similar experiments).
_"Carnallite may not only be fluorescent but is capable of being phosphorescent"_ very interesting
The only trouble with the magnesium sponge is that it's filled with salt from the melt, making it largely unreactive outside of aqueous solutions (which dissolve those salts away, of course). I tried burning it and it wouldn't even ignite, since the salts kept coating and protecting the metal within.
Any ideas for removing the salts and leaving the magnesium sponge intact? Maybe it would still work for those kinds of reactions anyway? I'm unsure...
@@ScrapScience Only thing I can think of is putting it inside a blender along with mineral oil and dump the resulting mix in water to get rid of the salts. I would probably directly use it in the sodium production.
@@ScrapScience just an idea what if you put it into coffee grinder; theoretically the chloride shouldnt be a problem anymore ( but on the other hand it can become pyrophoric not sure)
Thanks for this video very well presented. I would have tried melting all the black stuff under argon or nirogen gas to melt the magnesium, that would have incresed your yield of Mg metal.
Have you thought of Purifiying it with Mineral Oil sort of what you would do with Sodium? Or is that a totally different method? Cool video indeed.
Melting point is way too high for mineral oil.
Yeah, if we could melt the magnesium at temperatures that mineral oil could tolerate, this would be a great method for coalescing the metal. Sadly, magnesium's melting point is way too high.
You should do this is magnesium sulphate it's very cheap and common from Epson salt
When anhydrous, magnesium sulfate will decompose before melting. Even if it didn't, the anode reaction would be rather complex when working with molten sulfates. I don't know exactly what would happen, but I doubt it would be well-behaved...
I'm pretty sure one of the experiments in my 80's chemistry set proved that Magnesium in ribbon form does react with water. Think it had to be hot water. This is from memory, I'm trying to hunt out the manual in my terribly organised bookshelf.
Can l make mg with Mgso4
Anhydrous magnesium sulfate decomposes before melting, so it's unlikely.
Nice spot to film the intro. Do you mind sharing where that is?
Tasmania is the one! I'm afraid I won't be able to be any more specific than that though.
@@ScrapScience Fair enough! I should visit the country some time, it seems like time well spent.
Not surprising considering that little bead was around 10mg you are looking at about 250J pf which at sumch small piece and considering the process is adiabatic around 22%-25% of the energy is the form of visible light which means around 60J in light for about four seconds that means about 15W of light on average that is hell bright of light hou know
Maybe just try to collect more of that spongy stuff and melt it in solid blob?
I didn't explain it very clearly, but that was kind of the idea at the end of the electrolysis run. I tried making a whole bunch of the spongy stuff and then stopped the electrolysis and raised the heat to get it all to melt together. It seems like the sponge was resistant to melting together, but maybe it just needed to get hotter for longer... I don't know.
The comments are getting weird .
We now have Schrodinger's chemist .
Bravo.......screw magnesium.......it makes hydrogen .......when u put in water........measurements how much yield......also reactor with gas shield.....cheers
Well done !
Super! Thank you very much!
Hi, I've been interested in chlorate production for quite a while now. Finally I found mmo anodes at a reasonable price in the EU. But I have a question and I think you are the person who might know the answer. Question is: what's better? Plate or mesh? Both anodes are in same price. From NurdRange's vid I know that KClO3 crystals can grow in the holes of the mesh blocking electron flow, and that's what I'm afraid of. I guess mesh has a larger surface area so I don't know which one to pick. Hope you know the answer
Crystals growing in the holes is not a huge problem, provided you’re not running your anode at a very high current density. Even with plate anodes, crystals can still easily deposit on the electrode surface.
Even the surface area of the plate vs mesh is not significantly different most of the time.
All up, it’s pretty much a choice that doesn’t matter. Pick the one you think would look the best I suppose.
Thats really cool shit, never seen that anywhre
great video
*You find the most awkward way to do the simplest jobs. The Graphite Crucible should have been your negative and the Carbon rod be the positive. Minimum temp 1400F. at 3.3 volts and 20 amps. Result: 24 grams MG metal in bottom of Crucible in about 10 minutes. Use Magnesium Oxide and Sodium Fluoride and just keep adding more MgOx as needed. Very similar to Aluminum making process.*
You might benefit from a little research into Faraday's laws of electrolysis...
@@ScrapScience You should know better than to use a Small Cathode and plus a LARGE Anode in combination in molten salt electrolysis.
I dunno mate, you always make it very difficult to take your comments seriously, lol.
A 1500% error in your original claim doesn't make me very keen to follow your advice...
@@ScrapScience Let your EGO slip away. Only then can you learn from a Master Chemist.
Cool!
Of course it’s reacting with water you put potassium in it. I’m no chemist but it seems obvious to me
Potassium is notoriously difficult to deposit from molten salts like this, especially mixtures containing chloride. If we actually deposited potassium alongside our magnesium product, that would be a surprising result.
Are you by any chance making some kind of acid that's accelerating the reaction between water and magnesium?
Your reaction produces chlorine gas, and your starting products contained water. Is it possible that you're making some hydrogen chloride somewhere in the process?
What's the pH of the solution after dunking the magnesium sludge into it?
You had potassium in it
Yes
@ really cool vid and thought journey nonetheless
Whoops, sorry for the short reply, that was supposed to be a much longer response.
Out of interest, what makes you say that there's potassium in the product? Potassium is significantly more difficult to extract as a metal under these conditions.
@ it doesn’t matter the difficulty it matters that it’s possible. If it’s possible then it’s probable it just comes down to what degree. How much potassium. Then also I don’t think you’re accounting for the heat and its thermal reaction being added to the process. Think of hydrochloric acid… it’s not good for dissolving au…. But it is able to attack enough to produce an oxide layer that it can attack slowly, but if you take that same hcl and add heat it increases the rate of ionic exchange.
@ so going by the slow reaction and the fact it’s to h20 given then process used the most obvious thing to think is that you carried over some potential potassium. Until proven otherwise I think that’s the logical conclusion…. But I’m also smart enough to know you’re smarter than me that’s simply a country boys thoughts… the only other thing I can think of is a little sodium metal was made…. But I like the potassium view.
make oxigenated wather
i know you can't make clickbaits nor would but small yields decreases viewers
let him cook
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Hi!
en.wikipedia.org/wiki/FFC_Cambridge_process Reducing metal oxides to metal in molten CaCl2 or it's eutectics (NaCl/KCl/CaCl2)
Love it! In fact, it's already on my list of future video ideas. It's a very cool process.