max.yevs

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

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.

ah, that's interesting... i do apologize, my experience with chemicals does not exceed what you could buy at the local supermarket, so most of the things i know are looked up... but i do like the idea of bubbling SO2 into H2O2 instead of SO3 into H2O... Merged post follows: Consecutive posts mergedI was down at a hardware store today, and i happened to find a big case of sodium metabisulfite, used for treating rust stains (i guess Fe2O3 + Na2S2O5 > Fe + Na2S2O8 or something similar?)... i read somewhere that sodium bisulfite releases S2O when mixed with water, but i hope you have the equipment to catch it!

i tries elctrolysing salt water with ac straight from the plug... its very energetic and makes tiny explosions at both electodes (because hydrogen and oxygen are being produced one after the other really quickly, and with enough electricity to spark them)... but i do suppose some gas escapes

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... 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

thanks, although i will probably have to try the reaction anyways out of my own curiosity 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

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...

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.

i believe mercury is a liquid because one of its subshells is completely filled, which makes it not very reactive with neighboring mercury atoms (its right at the end of the transition metals)... on the other hand titanium and tungsten have those shells just half filled which makes the atoms more chained with each other and results in higher boiling and melting temp.

well iron sulfide, i haven't had the pleasure of encountering it, but i thought it ignited at room temperature...? isn't it pyrophoric?

why are all the metal oxides insoluble in water... and is that the reason why metals generally form hydroxides in water instead of oxides... i.e. 2Li + 2H2O > 2LiOH + H2 not 2Li + 2H2O > 2LiO2 + 2H2 Merged post follows: Consecutive posts mergedah ok i know why hydroxides are more favorable then oxides.. why should a metal, i.e. lithium, have to share its oxygen with another lithium, when it could share it with a hydrogen, which doesnt reduce oxygen's oxidizing power as much, allowing oxygen to be a stronger oxidizer to lithium... forget i posted this post

well you could always get HCl by reacting NH4Cl with a stronger acid, i.e. H2SO4, but then you might as well use table salt instead of ammonium chloride... you can get NH3 by reacting NH4Cl with almost any base, i.e. react it with baking soda to make ammonium carbonate, which decomposes with only a little heat...

actually bleach with hcl works pretty well, free chlorine is much better than chloric acid... but don't do it inside, quite a bit of chlorine gets generated

hydrogen bonding in water, pehaps? that is much stronger than london dispersion... if not, probably some other polar attraction... oh sorry misunderstood the question

oh, right... thats how NI3 is made... so not only would it not work, but unless you keep the chlorine and ammonia seperate, youll get dangerous chloramines Merged post follows: Consecutive posts mergedin fact thats the main way to produce NCl3 nowadays http://www.chm.bris.ac.uk/motm/ni3/ni3c.htm i should probably learn a bit more chemistry before suggesting such ideas

interesting problem electrolysing probably won't work, it will probably get you 2NH4Cl > 2NH3 + H2 + Cl2

yeah, thanks, so far i've just been trying to determine whats more reactive just by using a thermochemical calculator , chemix... its just i had kind of assumed that reactivity is a direct reflectment on electronegativity... which works most of the time but not always

ok yeah you're right it would not have to be spontaneous, but in this case, i think it would be since ammonium nitrate and hcl ionize in solution anyways... and yeah sorry when i was posting this i didnt realize id be writing way back on the second page

some plasters are made with calcium carbonate, although probably not of the purity you need... of course you can always react (almost) any calcium compound with sodium carbonate or bicarbonate... i.e. plaster of paris has 99% calcium sulfate... and then filter out CaCO3 since it has very low solubility... of course idk i would think that you can just replace Na2CO3 for CaCO3 of course there is another way... keep boiling tap water in a kettle or pot, adding more when it almost boils away... you'll find some white deposits, calcium carbonate.

By any chance, are you thinking of a BZ reaction? oscillates in color back and forth... http://en.wikipedia.org/wiki/Belousov-Zhabotinsky_reaction It's really quite remarkable, colors oscillate for quite a while Merged post follows: Consecutive posts mergedactually in that video, they only oscillate once, but just type bz reaction you should get quite a bit. i believe they all have something to do with bromine and different colored cations

for a while, i had assumed that a substance's reactivity depends on its electrongativity... but i'm starting to find some faults with this... for example, what is more reactive, copper or hydrogen? copper has lower electronegativity but is lower in the reactivity series. And if hydrogen is more reactive, shouldn't copper chloride or copper sulfate be a stronger lewis acid then hcl or h2so4, respectively? and what's more reactive, chlorine or oxygen? oxygen has higher electronegativity, but say preheated aluminum foil will burn in chlorine but usually not in oxygen... and for example sulfur, iodine, and carbon all have similar electronegativity, but I2 will react with aluminum scraps, while if anything, carbon and sulfur burn. (of course for the chlorine and iodine, how close they are to the stability of a full shell might explain their reactivity)

yes thats why aluminum is so reactive yet corrosion proof... but it somehow works in thermite compositions, i think ist either because its powdered or al gets vaporized... But its alright i'll just have to try it again sometime

it actually kind of surprises me you can buy uranium online if not from ebay then somewhere like here http://www.unitednuclear.com/chem.htm

try to add an oxidizer to the acid, like nitric suggested, as it helps pull the hydrogens off tha acid to react faster.... (vinegar plus hydrogen peroxide can dissolve copper)... if you have liquid bromine available, that might be worth a try Also i've found that 3% hydrogen peroxide by itself makes short work of iron, even more so than chloric acid...

well it would definately be harder to oxidize argon then krypton or xenon, but theres at least one compound, argon fluorohydride... but these compounds are not too useful, they often decompose well below room temperature...