elementcollector1 Posted February 17, 2012 Posted February 17, 2012 As an experiment, I attempted to make pure MnO2 from the black battery sludge inside a carbon-zinc battery. An 'ingredients' list gave me this for the sludge: Manganese dioxide: 35-40% Zinc: 10-25% Potassium hydroxide (35%): 5-10% Graphite (natural or synthetic): 1-5% This is from a Duracell alkaline battery, but I figure it's close to the stuff I have (might be wrong, though). So, to purify it, I dissolved the stuff in a green (impure) solution of strong HCl, which gave me an even greener (think lime or yellow-green) solution. Then I got some NaOH, dissolved in water, and added that to precipitate out the manganese. However, I have a mixture of brown (supposedly my finished product) clouds and more orangy clouds (which look like iron impurities. Upon settling out, the whole mixture is brown. Upon filtering out, the mixture is black. I tried washing with vinegar (which is supposed to get rid of iron and not manganese) but ended up with a much smaller amount of black solid than I started with, which makes me think the vinegar got quite a bit of the manganese as well. The only thing wrong with this scenario is that there are no iron impurities in the list above. So, my questions to you: -Why do there appear to be iron impurities? Is it my imagination? -Why is the manganese dioxide magically changing color from brown to black upon filtering?
John Cuthber Posted February 18, 2012 Posted February 18, 2012 Why do there appear to be iron impurities? I dissolved the stuff in a green (impure) solution of strong HCl,
elementcollector1 Posted February 18, 2012 Author Posted February 18, 2012 Why do there appear to be iron impurities? I dissolved the stuff in a green (impure) solution of strong HCl, Okay, so what is the impurity in the HCl that makes it green? It can't be iron, because it was stored in a plastic bottle.
Suxamethonium Posted February 18, 2012 Posted February 18, 2012 Um. There is one major issue here which I will get to later- but first lets address the other points. For commercial HCl to be green it is either coloured specifically with some organic dye. Or it has very significant impurities (probably of the transition metals). However Iron in concentrated hydrochloric acid is yellow due to the formation of the FeCl4-complex. It could be another transition metal though- maybe copper or nickel- although these too will be yellow in HCl of sufficient concentration. Iron in the battery depends on the battery- I happened upon some batteries where the case was iron (steel) with a zinc coating on the inside, or some with the electrode cap being iron- these could corrode into the battery over time. Was the ingredients list specific for that battery, or just generalised? Also, it may have only listed hazardous or active components in which case iron might be listed with 'other non-hazardous contents'. Now, from my experience with trying to extract transition metals from batteries. I personally got the green colour myself dissolving cadmium batteries in conc. HCl. I put the green to the presence of nickel (it went yellow when I replenished the acid), and cadmium. I concluded manganese would not impart any colour to the solution and this is the main issue in your method. The HCl is oxidised by the MnO2 into Cl2 H2O and MnCl2. Changing pH would then yeild Mn(OH)2 not Mn(OH)4 the latter of which would dehydrate to MnO(OH)2 and on heating finally to MnO2. Instead you will need to reoxidise Mn(OH)2 or change the original acid so as to maintain the +4 oxidation state (sulfuric comes to mind). Finally- acetates of manganese are soluble (In fact- most simple organic acids will have soluble metal salts- even lead!). Manganese (iii) acetate can spontaneously disproportionate into manganes (ii) and (iv) acetate- this may be an efficient way to reoxidise you manganese by utilising a reagent to oxidise (ii) to (iii) which disproportionates back to (ii) and some (iv), the (ii) can then react with the oxidising reagent again.
John Cuthber Posted February 18, 2012 Posted February 18, 2012 (edited) Okay, so what is the impurity in the HCl that makes it green? It can't be iron, because it was stored in a plastic bottle. Oh yes it can. Iron is a very common impurity in a lot of things- simply because it's rather common. Also FeCl3 is slightly volatile so it can be carried over with the HCl during the manufacture of the acid. If the HCl is concentrated enough it will be yellow but if it's relatively dilute it will be green. Edited February 18, 2012 by John Cuthber
elementcollector1 Posted February 18, 2012 Author Posted February 18, 2012 So, dissolve the precipitate in vinegar completely, add an oxidizer such as H2O2, and add NaOH to precipitate out Mn(OH)4? Manganese salts are typically pale pink. It's said that manganese hydroxide spontaneously decomposes into manganese dioxide and water upon contact with oxygen (dissolved in water). So, this explains the color change after filtering with vinegar, as some of the stuff must have already done this and made the black, pure manganese dioxide. But how does this get rid of the iron?
Suxamethonium Posted February 19, 2012 Posted February 19, 2012 So it is proposed that: Mn(OH)2 + 1/2O2 --> MnO2 + H2O This would probably work- but I doubt it would be spontaneous. You would surely need some energy if not just to drive off the water. Manganese oxides structures tend to be fairly variable at the best of times, you're probably more likely to get something like (and I just pulled this formula out of my head to illustrate the point): Mn(OH)2 + O2 --> MnO(OH)2 or even worse a co-complex like MnO.MnO(OH)2 ... Going straight from (ii) to (iv) is also less likely to be spontaneous as say (ii) to (iii) to give the very common brown Mn2O3 (also forms messy complexes like MnO.Mn2O3). You're best bet is probably to find something unusual that Mn forms a soluble salt with (possibly ligands would be a good start), or go the opposite way and see if ligands like EDTA are non-effective for Mn ions (in which case EDTA could stop Fe precipitating with the Mn). Then worry about oxidation states. As for the H2O2 and NaOH. Hydrogen peroxide is not a good idea. Mn salts catalyse its decomposition, and salts (iv) and higher are often reduced to (ii) or (iii) salts. Also, you wouldn't add the NaOH until the end of the reaction or you would precipite everything and get no reaction at all.
elementcollector1 Posted February 20, 2012 Author Posted February 20, 2012 How about this (NurdRage method): -Dissolve crude mix in HCl. Filter to get impure, iron contaminated solution. -Take 1/3 of solution and add NaOH to precipitate out metal hydroxides. -Add precipitate back into solution after washing. -Let stand overnight. Reactions: MnO2, Fe3+? + HCl -> MnCl2 + Cl2 + FeCl3 + H2O MnCl2, FeCl3 + NaOH -> Fe(OH)3 + Mn(OH)2 MnCl2, FeCl3 + Mn(OH)2, Fe(OH)3 -> MnCl2 + Fe(OH)3 is the link. This works for MnCl2 as well. So, now that I can get MnCl2 in a pure aqueous solution, what can precipitate out MnO2? Could H2O2 work (oxidize II to IV? Unlikely, apparently)? Or should I use potassium peroxymonosulfate (seen here )?
Suxamethonium Posted February 20, 2012 Posted February 20, 2012 I feel that last equation is wrong, but the process/products is correct. Atmospheric oxygen is doing the oxidation (I think of Fe(OH)2 to Fe(OH)3 which is even less soluble)- this is why it's 'purifying' the Mn(OH)2- which somehow dissolves and is left in solution with the chloride ions. It looks like it could be a complex combination of reactions, which my best guess is something along the lines of: Fe(OH)2 (s) (miniscule amount)> Fe(OH)2 (aq) Fe(OH)2 (aq)+ 0.5O2 +H2O (or + H+) > Fe(OH)3 + -OH (or + H2O) Fe(OH)3(aq) (large percentage)> Fe(OH)3(s) Now there is 'more room' for more Mn(OH)2 and Fe(OH)2 to dissolve in miniscule amounts once again (even though they are 'insoluble' a miniscule ammount dissolves enough to get a slow reaction- these are both more soluble then the Fe(OH)3 which is why the reaction is driven forward- by Le Chateliers Principle). Overall, all of the Fe2+ becomes Fe3+. And as such all the iron recipitates out as iron (iii) hydroxide in preference to the Mn2+. If the Fe3+ somehow sequested all the HO- present in solution then only Mn2+ and Cl- would remain in the solution. You also have a significant amount of Cl- present and the video isn't exactly clear, but I imagine it is slightly acidic initially- otherwise the alkalinity created by the reaction would cause more precipitation. In such a case maybe the OH ions in the Mn(OH)2 can be displaced by the Cl- ions? Like I said, I'm unsure, but it is definately not a single reaction occurring. By comparison the second video is straight forward. I would use the KHSO5salt if you can get it, if not use a similar salt like NaHSO5 or NH4HSO5 or even peroxysulfuric acid. If still no luck try the dication salt mixed with sulfuric acid or as a last resort try an acidic solution of S2O82- salt (much less likely to work). I still feel hydrogen peroxide will not work effectively- although it seems Mn(iv) wont be reduced by the peroxide like Mn(vii) it still directly catalyses its decomposition. The reason is because even without the loss of Mn(iv) product, as more of it is prodcued the rate of hydrogen peroxide breakdown will increase thus decreasing the rate of the usefull reaction significantly (and increasing the amount of hydrogen peroxide needed, and heat generated).
elementcollector1 Posted February 20, 2012 Author Posted February 20, 2012 I feel that last equation is wrong, but the process/products is correct. Atmospheric oxygen is doing the oxidation (I think of Fe(OH)2 to Fe(OH)3 which is even less soluble)- this is why it's 'purifying' the Mn(OH)2- which somehow dissolves and is left in solution with the chloride ions. It looks like it could be a complex combination of reactions, which my best guess is something along the lines of: Fe(OH)2 (s) (miniscule amount)> Fe(OH)2 (aq) Fe(OH)2 (aq)+ 0.5O2 +H2O (or + H+) > Fe(OH)3 + -OH (or + H2O) Fe(OH)3(aq) (large percentage)> Fe(OH)3(s) Now there is 'more room' for more Mn(OH)2 and Fe(OH)2 to dissolve in miniscule amounts once again (even though they are 'insoluble' a miniscule ammount dissolves enough to get a slow reaction- these are both more soluble then the Fe(OH)3 which is why the reaction is driven forward- by Le Chateliers Principle). Overall, all of the Fe2+ becomes Fe3+. And as such all the iron recipitates out as iron (iii) hydroxide in preference to the Mn2+. If the Fe3+ somehow sequested all the HO- present in solution then only Mn2+ and Cl- would remain in the solution. You also have a significant amount of Cl- present and the video isn't exactly clear, but I imagine it is slightly acidic initially- otherwise the alkalinity created by the reaction would cause more precipitation. In such a case maybe the OH ions in the Mn(OH)2 can be displaced by the Cl- ions? Like I said, I'm unsure, but it is definately not a single reaction occurring. By comparison the second video is straight forward. I would use the KHSO5salt if you can get it, if not use a similar salt like NaHSO5 or NH4HSO5 or even peroxysulfuric acid. If still no luck try the dication salt mixed with sulfuric acid or as a last resort try an acidic solution of S2O82- salt (much less likely to work). I still feel hydrogen peroxide will not work effectively- although it seems Mn(iv) wont be reduced by the peroxide like Mn(vii) it still directly catalyses its decomposition. The reason is because even without the loss of Mn(iv) product, as more of it is prodcued the rate of hydrogen peroxide breakdown will increase thus decreasing the rate of the usefull reaction significantly (and increasing the amount of hydrogen peroxide needed, and heat generated). Potassium peroxymonosulfate can be bought at pool stores, so that might be a good source. The reaction does work, though? I did guesswork on the equations, and left out the O2 and H2O in the last one (guess I shouldn't have).
Suxamethonium Posted February 25, 2012 Posted February 25, 2012 Yeah, that guy is pretty reliable as far as I have known- the reaction should work.
elementcollector1 Posted February 25, 2012 Author Posted February 25, 2012 Yay! As soon as I get back from California, I'll check the color of the solution. I was in a bit of a rush, so I simply put the precipitate (still in the filter) into the beaker and left it there for a few days. This won't affect anything, right? Ions still pass through and react.
elementcollector1 Posted March 1, 2012 Author Posted March 1, 2012 Just tried the bleach-NaOH method (src. Pyroguide) and it appeared to work. I now have a very dark brown, chunky precipitate and a chocolate-colored solution (probably stuff that escaped through my mediocre filters). My starting product was contaminated with Al, but this doesn't seem to be a problem as Al(OH)3 is soluble in acids and alkalis (just not water). How do I confirm that this is in fact MnO2?
Suxamethonium Posted March 12, 2012 Posted March 12, 2012 hmmm confirmation methods. Best way is probably google lol. I don't know any specific reactions to try.
elementcollector1 Posted March 12, 2012 Author Posted March 12, 2012 hmmm confirmation methods. Best way is probably google lol. I don't know any specific reactions to try. Doesn't it decompose H2O2? Then again, a lot of things seem to be able to do that. I have some pure, clear HCl now (I restocked), so I might see if I get a clear or pink solution and chlorine production upon dissolving this stuff.
John Cuthber Posted March 12, 2012 Posted March 12, 2012 Probably the most characteristic reaction of manganese is that you can oxidise it under strongly alkaline conditions to green Mn(VI) and then if you add acid it disproportionates to give purple Mn(VII).
elementcollector1 Posted March 14, 2012 Author Posted March 14, 2012 Probably the most characteristic reaction of manganese is that you can oxidise it under strongly alkaline conditions to green Mn(VI) and then if you add acid it disproportionates to give purple Mn(VII). On a new run, I now have a solution which is not contaminated by iron, but something else. Judging from precipitations of carbonate and hydroxide, both being off-white, it's something similar to manganese. Any thoughts on figuring this out?
Suxamethonium Posted March 17, 2012 Posted March 17, 2012 Um, do you have access to Ion-exchange chromatography at all? If so, you should be able to identify any cations present. Otherwise, I don't really know- a large number of compounds have white or off white precipitates (I included white, because often impure precipitates of "white" compounds appear off-white).
elementcollector1 Posted March 17, 2012 Author Posted March 17, 2012 Um, do you have access to Ion-exchange chromatography at all? If so, you should be able to identify any cations present. Otherwise, I don't really know- a large number of compounds have white or off white precipitates (I included white, because often impure precipitates of "white" compounds appear off-white). Ahaha, no. I do, however, have access to a fresh supply of battery stock.
Recommended Posts
Create an account or sign in to comment
You need to be a member in order to leave a comment
Create an account
Sign up for a new account in our community. It's easy!
Register a new accountSign in
Already have an account? Sign in here.
Sign In Now