max.yevs Posted March 20, 2009 Posted March 20, 2009 is it possible to extract magnesium from epsom salt (MgSO4) by electrolysis? i know that it is very hard to oxidize sulfate ion and in solution, water would much rather get split up then magnesium sulfate... which is why im thinking, what if i put table salt in the solution along with it, i would assume 2NaCl + MgSO4 > Na2SO4 + MgCl2, which would then make magnesium plate the cathode when electrolysing.... i assume that if magnesium hydroxide tries to form, it would just get electrolized also would this work? i'm very interested to handle magnesium.
UC Posted March 20, 2009 Posted March 20, 2009 No, all you will produce in aqueous solution is magnesium hydroxide. You need molten salts of some kind or a mercury cathode to get metal.
max.yevs Posted March 20, 2009 Author Posted March 20, 2009 yes, that is one of the things i'm thinking might happen... but even if magnesium hydroxide does form, it will get ripped back apart... i think this really depends on magnesium's reactivity when electrolysing sodium chloride, the sodium does not even dream of plating the cathode, it will make NaOH...sodium is too reactive so I don't know, is water a better oxidizer then a stream of electrons? and does magnesium react with cold water, unprovoked? the reactivity series chart says it only reacts with acids, but idk... Merged post follows: Consecutive posts mergedi did some research (and i know i should have done this before asking) and in fact electrolyzing MgCl2 from seawater is the major process for producing magnesium... so if anything wouldnt work its 2NaCl + MgSO4 > Na2SO4 + MgCl2 but i think it should. Merged post follows: Consecutive posts mergedgood thing about this is that magnesium sulfate is available at any pharmacy or superstore (epsom salt) and pure magnesium should be kind of fun to play around with (i hear it burns quite well) ill post if it works Merged post follows: Consecutive posts mergedthen again, maybe they meant commercially its by electrolysing molten magnesium chloride, not mgcl2 solution... will have to test this... i think with enough electricity its possible to do it with solution...
hermanntrude Posted March 20, 2009 Posted March 20, 2009 the trouble is that magnesium chloride and sodium sulfate are both soluble in water. What that means is that there is no reaction. all that happens is both your reactants will dissolve leaving you with four different ions floating about in solution. If one of the products were insoluble, a precipitate would be seen and u could isolate it.
Theophrastus Posted March 20, 2009 Posted March 20, 2009 I would have suggested a single displacement reaction, however, magnesium is quite reactive, and the only more reactive metals, I can think of, are the alkali metals, as well as some of the alkali earth metals. an easy way to obtain magnesium, is through the body of pencil sharpeners, as those are either made from steel or magnesium, and from that point on, you can simply judge it by the weight. The blade, however, is always made of steel, so its best you remove it with a screwdriver. In regards to chemically producing it, with epsom salt, you can react the epsom salt, with ammonium hydroxide (ammonia + water solution), producing ammonium sulfate and magnesium hydroxide. (Yes, I know, same subtance as would be procured in your electrolysis, but you would obtain a better yield), which when heated, produces magnesium oxide... and then.... I have absolutely no idea, what to do. Best of luck!
max.yevs Posted March 20, 2009 Author Posted March 20, 2009 thanks, although i will probably have to try the reaction anyways out of my own curiosity the trouble is that magnesium chloride and sodium sulfate are both soluble in water. What that means is that there is no reaction. all that happens is both your reactants will dissolve leaving you with four different ions floating about in solution. If one of the products were insoluble, a precipitate would be seen and u could isolate it. yes, i realize that there would not be any actual reaction, but when you get Na+, SO4-2, Mg+2, and Cl-, trying to electrolyse it should be the same as electrolysing MgCl2, since the sodium and sulfate ions are harder to reduce/oxidize. and thanks theo, ive considered trying to get magnesium hydroxide, magnesium oxide, and magnesium carbonate, but trying to decompose any would give magnesium oxide... which will probably be impossible to decompose, except maybe electrolysing but i guess the only approach is to try it... just as soon as i get some epsom salt
Theophrastus Posted March 20, 2009 Posted March 20, 2009 In regards to what has been said, firstly, one has to become careful of the electrodes you use, is this is how you're planning to attain magnesium, as while magnesium will precipitate into the solution, most metal anodes, will then react with the sulfate, for example, if you are using copper electrodes, the reaction would form magnesium and copper sulfate. The magnesium, now magnesium hydroxide, having been dissolved in water, would then displace the copper, and you would be left with copper hydroxide (which actually has quite a number of useful applications!) and magnesium sulfate. To summarise, using any metal, that is less reactive than magnesium, would simply lead to the reverse of the initial reaction, going nowhere. Due to this, I recommend you use carbon rod electrodes. For this, graphite will do. However, I have a feeling that the magnesium, would then become magnesium hydroxide, reacting with the water. But hey, you can't knock it, 'til you've tried it, right? Best of luck! ,theophrastus ps: let us know the outcome. It may be quite interesting! (magnesium powder, could be quite useful)
max.yevs Posted March 21, 2009 Author Posted March 21, 2009 (edited) bought some epsom salt today... and set up the experiment, its still running for the electrodes, i used two small steel plates... i'm actually very impressed with how well steel works for this... when i use copper electrodes, they quickly get corroded and stop working... i think its the carbon content...or better yet, the fact that iron doesn't passivate as much as copper... I'm using pretty weak power, for now, 5volts/.55amperes.... anyways, one of the steel plates is actively bubbling... the other i cant see bubbling at all (although it might be bubbling very little and the gas dissolving)... the one that's not bubbling is slowly turning a graphite like grey color- with no residue... the one that's bubbling (i think its bubbling hydrogen) has some white or silvery white deposits.... which means either silvery white magnesium or white and insoluble magnesium hydroxide... its probably Mg(OH)2..., if so, ill try using more power... i.e. a small lead-acid battery or a bunch of 9v alkaline in series... im sure that with enough power you can plate magnesium onto the cathode... To summarise, using any metal, that is less reactive than magnesium, would simply lead to the reverse of the initial reaction, going nowhere. Yes, of course, but what if i take out the electrodes, one of which is hopefully plated with magnesium, before turning off the power? the higher electropositivity of magnesium would make it more willing to stay at the cathode, assuming enough power is provided... Merged post follows: Consecutive posts mergedyeah it was definitely insoluble magnesium hydroxide produced at the bubbling electrode, which was the cathode...and the bubbles were hydrogen, i assume... i'm surprised i did not get any chlorine or oxygen bubbles... so i guess everyone was right... however i havent given up yet... i did find a tiny slice of magnesium on the side of the glass i was doing it in... (i think its magnesium because it burned very well, but idk how it got there) Mg(OH)2 is electrolyzable, and i think with enough power I will be able to plate an anode with magnesium, and maybe, just maybe, prove wrong all those other magnesium threads... if somehow i do get it to work, ill post some photos of the setup and results... maybe over the weekend Merged post follows: Consecutive posts mergedbut it seems to me that since the electrodes are basically losing and gaining electrons, the cathode is like an oxidizer and the anode is a reducer... so if enough electricity is supplied, the cathode's oxidizing potential should be higher than the oxidizing potential of the OH radical... so, there is some truth to trying to electroplate a piece of metal with magnesium, which in essence means electrolysing a magnesium compound to make magnesium Edited March 21, 2009 by max.yevs
Theophrastus Posted March 21, 2009 Posted March 21, 2009 totally unrelated, but I just realised the answer to my previous problem in refining the magnesium oxide, to pure magnesium. If you are to heat the magnesium oxide with charcoal, the following reaction would ensue: C + 2MgO > 2Mg + CO2 I really can't believe I didn't think of this earlier, having recently read some of isaac azimov's nonfiction books, regarding the history of chemistry. This was the primary method used by alchemists and metallurgists, to extract pure metals, from their respective oxide ores. It mentions it, in the context of copper, iron and tin, however, I can't see why it couldn't work for this. Best of luck in your electrolysis though. Tell us whether it was pure magnesium or magnesium hydroxide that formed. If the first, than your goal has been accomplished, if not, you can simply dehydrate the hydroxide to oxide, and then, do the above. Best wishes, ,Theophrastus
John Cuthber Posted March 21, 2009 Posted March 21, 2009 One of the interesting things about magnesium is that it burns in CO2. 2Mg + CO2 -> 2MgO +C
Theophrastus Posted March 21, 2009 Posted March 21, 2009 Ah, alas, a method gone wrong. I suppose upon buring, it would oxidize to magnesium oxide, reversing the reaction that originally took place. Shame... (Though thanks for the clarification)
max.yevs Posted March 21, 2009 Author Posted March 21, 2009 I've given up on this magnesium extraction... I hooked up several 9v batteries in series, to get a pretty decent voltage... enough to create very bright sparks... voltage in my case was 20 times more important than amperes... Didn't work... some specks of magnesium may have formed on the cathode, but not any reasonable amount... I'm sure with enough voltage, it could be done... I even know how much is needed... the amount of voltage produced from reaction of magnesium and free hydroxide... (which would take quite a few alkaline zinc-manganese dioxide reactions) I'm sure I will try this again, when I happen to have a good enough power source lying around, but probably not anytime soon... P.S. the electrolysis quickly created a very dense blue suspension... i suspect some kind of iron salt, since i was using steel electrodes. Anybody know what it is? P.P.S. I'm very impressed with using steel as electrodes. P.P.P.S. Please don't comment on how I'm doing this in my kitchen.
max.yevs Posted March 22, 2009 Author Posted March 22, 2009 i took another picture of the suspension, after i poured it in an improvised tray... anyone know what it is? im guessing some kind of iron salt, but i don't think its iron sulfate, i made iron sulfate before and it looks different... it starting to rust a bit... pretty quick actually
max.yevs Posted March 22, 2009 Author Posted March 22, 2009 overnight, the bluish green compound slowly turns a brilliant orange... perhaps iron oxide, but its a brighter orange than any rust i've ever seen. i've given up on the magnesium a long time ago, i'm just interested as to what this complex is.
max.yevs Posted March 22, 2009 Author Posted March 22, 2009 a drop of hydrochloric acid makes the solution a clear bright green... (bottom of cup) looks very much like iron(II) chloride... so... what is an iron compound that: - Makes a dark bluish green suspension - Turns into a brilliant orange overnight - And the anion must be weaker then chloride so that it forms FeCl2 with HCl
max.yevs Posted March 23, 2009 Author Posted March 23, 2009 (edited) ah, ok i know what happened. Either the Cl2 or SO4 attacked the steel anode, creating some iron sulfate or chloride. Then the magnesium or sodium hydroxide that formed at the cathode reacted with the iron sulfate/ chloride. So, im guessing the first picture is some kind of 'green rust' which apparently has the formula Fe(III)x Fe(II)y (OH)3x+2y-z (A-)z; where A- is Cl- or 0.5 SO4-2 ... anyways, something complicated. Green rusts oxidize very easily to Fe2O3, so the second picture was just a high grade iron (III) oxide. And i guess iron oxide reacts with hcl to form iron chloride and water. Edited March 23, 2009 by max.yevs
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