Sciencemadness Discussion Board

MnO2 -> MnSO4; What is the best route?

Hilski - 2-10-2006 at 08:49

I have about 1kg of MnO2 that I would like to try and convert to MnSO4, and I was wondering what the preferred method would be.
In all my searching, really the only way I have been able to find is (from the "Toluene -> benzaldehyde" thread) to heat the MnO2 with an excess of HCl to convert it to MnCL2, and then disolve the MnCL2 in conc. H2SO4. After evaporating the liquid, one should be left with MnSO4. Is this the best/most direct way to do it?

What would you end up with if you were to heat the MnO2 with conc. H2SO4?

Thanks in advance.

-Hilski

gsd - 2-10-2006 at 09:14

1) The direct reaction of H2SO4 with MnO2 does not happen

2) MnCl2 will not become MnSO4 if you heat it with H2SO4.

There are several ways of doing it.

1) Roast MnO2 at high temp ( above 750 Deg C) to get MnO and then react with H2SO4

2) Reduce MnO2 to MnO by mixing it with a heavy oil ( furnace oil etc) and then igniting the mixture in absence of Oxygen, and then react with H2SO4

3) Use MnO2 + H2SO4 as an oxidizing agent as in oxidation of Toluene to Benzaldehyde; However this reaction is very slow and yield are not good. Oxidation of p-Hydroxy Toluene to p-Hydroxy Benzaldehyde goes very smoothly.

4) Use MnO2 + H2SO4 for oxidizing aniline to Hydroquinone however the product is contaminated with Ammonium Sulphate.

5) Direct Reaction of MnO2 with SO2 ( if you can get hold of SO2)

6) Indirect method : React MnO2 with HCl ( careful ! Cl2 gas is evolved as a by-product) to get MnCl2 solution ; convert it to MnCO3 and then react with H2SO4 to ger MnSO4

gsd

[Edited on 2-10-2006 by gsd]

12AX7 - 2-10-2006 at 09:32

MnCl2(anh.) + H2SO4 = MnSO4 + 2HCl. Heat drives to the right. Mind fumes.
More completely, MnO2 + 2HCl(aq) + H2SO4 = 2H2O + Cl2 + MnSO4.
You can even do it with salt, at the expense of sulfuric acid; MnO2 + 2NaCl + 2H2SO4 = Na2SO4 + Cl2 + MnSO4.

I don't know if MnO2 + 2NaCl = Na2O + Cl2 + MnO occurs. Probably not, at least without oxygen to produce the manganite or manganate (MnO2 + 2NaCl + O2 = Cl2 + Na2MnO4).

Manganese can be reduced with carbon mildly exothermically as 2MnO2 + C = Mn2O3 + CO and Mn2O3 + C = 2MnO + CO. The first reaction occurs easily (spontaneous ignition when the mixture is heated to dull redness, reaching orange to yellow heat in the process), but the second needs higher heat to drive.

Come to think of it, the spinel Mn3O4 may also be involved. Spinels are generally very stable crystal structures. (For instance, the iron equivalent Fe3O4 is produced from FeO by disproportionation on cooling.)

FYI, direction reaction of MnO2 + H2SO4 yields little or nothing, however impurities of Mn(III) in e.g. pottery grade MnO2 may dissolve in strong acid to produce a pink to burgundy colored solution which hydrolyzes very easily when diluted.

Anyone know if easier reducing agents than SO2 (or equivalently, sulfites) will work? Reducing sugar for example, or ascorbic or citric acid, etc.

Tim

[Edited on 10-2-2006 by 12AX7]

unionised - 2-10-2006 at 10:12

"1) The direct reaction of H2SO4 with MnO2 does not happen

2) MnCl2 will not become MnSO4 if you heat it with H2SO4."
Are you sure? I have just checked a couple of textbooks that say that MnO2 disolves in hot conc H2SO4 to produce MnSO4 and I can't see why the displacement of HCl should be any different from MnCl2 compared to NaCl.

My personal choice would be to roast the MnO2 with carbon then disolve the (mixed lower) oxide(s) in hot conc acid or to use HCl if I had a use for the Cl2.

S.C. Wack - 2-10-2006 at 12:22

MnO2 most definitely dissolves in hot conc. H2SO4 with vigorous evolution of oxygen. This is not debatable. This says nothing of the usefulness of drain cleaner and crap pottery MnO2 in producing good MnSO4, however.

[Edited on 2-10-2006 by S.C. Wack]

guy - 2-10-2006 at 12:34

hydrogen peroxide and acid will very quickly reduce MnO2 to Mn2+.

Elawr - 2-10-2006 at 19:10

What if you were to fuse the MnO2 with Na2HSO4 and powdered charcoal?

Hilski - 2-10-2006 at 19:56

Quote:

6) Indirect method : React MnO2 with HCl ( careful ! Cl2 gas is evolved as a by-product) to get MnCl2 solution ; convert it to MnCO3 and then react with H2SO4 to ger MnSO4


That's actually the reaction I intended to describe, MnO2 -> MnCO3 after converting the MnO2 to MnCl2, and then precipitating with sodium carbonate . Somehow I managed to leave that part out though.

But this one looks easy enough:
Quote:
You can even do it with salt, at the expense of sulfuric acid; MnO2 + 2NaCl + 2H2SO4 = Na2SO4 + Cl2 + MnSO4.


Yeah it uses more H2SO4, but then you don't have to do the reaction with HCl. I may give this one a shot.
Has anyone tried this reaction, or have any more details like how to separate the MnSO4 from the Na2SO4?

Thanks for all the replies



[Edited on 3-10-2006 by Hilski]

Elawr - 2-10-2006 at 21:18

Since I am at work right now, I cannot confirm this will work as a one-step procedure. However, based on other experiments I have done, I suspect that it just might. I know that NaS can be had by fusing NaHSO4 with powdered charcoal. The sulfate ion is a better oxidizer at high heat, and we're tallking dull red or hotter - 700-800 degrees C or more. No aqueous chemistry here - this is anhydrous molten salt. Sulfides in aqueous mileu will easily reduce MnO4- straight to Mn++. So the next chance I get, I gonna melt down some Na2SO4 anhydrous and add a little MnO2 and then charcoal. If I am right, I should see effervescence from the melt, and then be able to recover MnSO4 from the cooled melt. That is, unless somebody beats me to it! Should be easy enough to precipitate the Mn++ as the carbonate by adding baking soda to the solution and decanting.
]:)

[Edited on 3-10-2006 by Elawr]

[Edited on 3-10-2006 by Elawr]

evil_lurker - 2-10-2006 at 22:18

Quote:
Originally posted by S.C. Wack
MnO2 most definitely dissolves in hot conc. H2SO4 with vigorous evolution of oxygen. This is not debatable. This says nothing of the usefulness of drain cleaner and crap pottery MnO2 in producing good MnSO4, however.

[Edited on 2-10-2006 by S.C. Wack]


Haven't tried it but have read it before, so I'm guessing its true. Manganese is a weird materal due to all its various oxidation states... doesn't help that there isn't very much material on the net about it.

And yeah, I would think pottery grade MnO2 and drain cleaner would suck. I'd recommend using some Mn metal from skylighter and excess concentrated sulfuric acid in the right proportions so you don't have to seperate the MnSO4.

Heat, dissolve, dilute, charge and your ready to go.

not_important - 2-10-2006 at 22:19

If you want reasonably pure MnSO4, avoid routes that will result in Mn and Na ions in solution together.

If you need chlorine for anything, then the
MnO2 + HCl (or NaCl + H2SO4) => Cl2 + MnCl2,
MnCl2 + 2 OH(-) => Mn(OH)2 or MnCl2 + CO3(2-) => MnCO3
Mn(OH)2 + H2SO4 => MnSO4
would do well. The toluene to benzaldyhde route if you want that, the low yield is of benzaldehyde, there's no loss manganese. Otherwise MnO2 + H2SO4 + alcohol or sugar to get MnSO4

You might even try mixing ammonium sulfate in excess and MnO2, heating that until it melts, hold at around 300-350 C for awhile, then increase the hea until all the excess ammonia and sulfuric acid/SO3 are driven off. Go all the way up to 700 C, when the MnSO4 fuses.

If you make Mn(OH)2, set a little aside before dissolving the rest in H2SO4. If you go for oxidising some organic compound with MnO2 + H2SO4, take a bit of the solution and precipitate out some Mn(OH)2 or MnCO3. You want to neutralise excess H2SO4 with Mn(OH)2 or MnCO3, and have a bit of excess of the base. Boil for a few minutes, let cool and settle, filter off the remaining Mn(OH)2 or MnCO3. This removes some impurities as carbonates or hydroxides absorbed on the excess MnCO3.

Heating solid MnSO4(some water) until it dehydrates and then more strongly until it just melts (say 720-730 C), hold there for a few minutes, then cool, dissolve in water, add a little ammonium carbonate or MnCO3, boil, cool, settle, filter will remove even more d-block metals. MnSO4 is more heat stable than most of the d-block sulfates, they break down into oxides or basic sulfates, when taking the MnSO4 into solution the added or formed MnCO3 again traps those impurities and makes it easier to remove them.

The ceramic MnO2 I've purchased appears to be fairly pure, a little iron seems to be the main impurity. But the recrystallization of manganese sulfate can be difficult, there are a number of hydrates and some have lousy cold/hot solubility ratios; that's why it's best to try to get as many impurities out early as you can. and not have a lot of sodium in solution with it.

unionised - 3-10-2006 at 09:51

"hydrogen peroxide and acid will very quickly reduce MnO2 to Mn2+. "
Sure about that ?
Last time I checked, MnO2 decomposed H2O2 so fast there would be very little chance for a reduction.

guy - 3-10-2006 at 11:08

Quote:
Originally posted by unionised
"hydrogen peroxide and acid will very quickly reduce MnO2 to Mn2+. "
Sure about that ?
Last time I checked, MnO2 decomposed H2O2 so fast there would be very little chance for a reduction.


Yes I'm very sure. This is how I extracted MnO2 from batteries for my Mn experiments. You mix MnO2 and acid first then add H2O2.

MnO2 + 4H + 2e----> Mn2+ + 2H2O

H2O2 ----> 2H+ + O2 + 2e

evil_lurker - 3-10-2006 at 12:55

Are you sure about this?

Last time I checked MnO2 was simply a catalyst for the decomposition of H2O2 and was not reduced by the reaction.

Thanks

Hilski - 3-10-2006 at 15:07

This is very informative. Thanks for all the replys.

My intentions are to combine any MnSO4 that is yielded, with the proper ammounts of (NH4)2SO4 and 60% H2SO4 to make mangano-ammonium sulfate. Then manganese-ammonium alum will be made using electrolytic oxidation per US patent 808095.
I'm sure most of you have already seen the great writeup Cyclonite posted in the Toluene -> Benzaldehyde thread. If you haven't, you should definately check it out.

I will probably go the MnO2 + HCl (or NaCl + H2SO4) => Cl2 + MnCl2, MnCl2 + NaHCO3 => MnCO3 route to make the MnSO4, simply because it is more feasible for me to do so.

Does anyone have any more detailed info on the actual process for MnO2 + 2NaCl + H2SO4 => Cl2 + Na2SO4 + MnSO4 ?

Thanks again



[Edited on 3-10-2006 by Hilski]

jon - 3-10-2006 at 15:13

from what I undertstand MnSO4 is used as a fertilizer just buy some and save yourself the trouble.

12AX7 - 3-10-2006 at 17:05

Quote:
Does anyone have any more detailed info on the actual process for MnO2 + 2NaCl + H2SO4 => Cl2 + Na2SO4 + MnSO4 ?


Add heat...

Incidentially, why is this thread under organic...still?

Tim

guy - 3-10-2006 at 17:40

Quote:
Originally posted by evil_lurker
Are you sure about this?

Last time I checked MnO2 was simply a catalyst for the decomposition of H2O2 and was not reduced by the reaction.


not in acid...!!!

Hilski - 3-10-2006 at 19:01

Quote:
Incidentially, why is this thread under organic...still?


Yeah, sorry about that. I realized after I started the thread that it was in the wrong forum.
Quote:
Add heat...

Yeah, but how do I get rid of the Na2SO4? MnSO4 and Na2SO4 are both soluble in water, although MnSO4 is about 3 times more so than Na2SO4.

not_important - 3-10-2006 at 22:50

Quote:
Originally posted by Hilski

Yeah, but how do I get rid of the Na2SO4? MnSO4 and Na2SO4 are both soluble in water, although MnSO4 is about 3 times more so than Na2SO4.


Fractional crystallization for the direct route. However, this can result in baldness through hair pulling; MnSO4 exists in anhydrous, mono, tri, tera, penta, hexa, and hepta hydrate forms depending on temperature during crystallization and other factors.

You could precipitate it out as the carbonate, filter and wash that, then dissolve in H2SO4.

Or you could do a little experimentation and try a direct route from heating MnO2 with excess (NH4)2SO4 and crystallizing the reaction product. The reaction is messy in the stoichiometry sense, as the 'NH4' goes off as water, nitrogen, ammonia, and sometimes a trace of NOx. You may need to add extra ammonium sulfate to the workup solution to have enough for the mixed Mn(II)NH4 salt (not an alum, that takes Mn(III)). The mixed Mn(II)NH4 sulfate can be isolate by crystallization; it is more soluble than ammonium sulfate hot, but less when cold, so concentrating until ammonium sulfate starts to crystallize, filtering hot, then cooling will let you get it out of solution. Yes, there's Mn left in solution, just let that evaporate and toss it into the next batch of MnO2 + (NH4)2SO4 before heating. Or you may be able to get away with just making sure there's enough ammonium sulfate in solution, as the procedure you mentioned says to use an excess of it, and taking the filtered solution of the cooled melt.

guy - 3-10-2006 at 22:56

Is the H2O2 method too expensive for you guys or something? Its the easiest cleanest method. I've tried it and it works perfectly and its quite rapid also.

not_important - 3-10-2006 at 23:10

Peroxide in some areas is only available as the 3% solution for medical purposes, is fairly expensive in terms of cost per mole H2O2, and adds a lot of water. Fertilizer grade ammonium sulfate is a reais or two per kg, which is over 7 moles.

For some things H2O2 is perfect, for others I use MnO2 (ceramic grade) or NaClO3 (from herbicide), and for certain fusion reactions NaNO3 (ceramic grade); there's a few things dichromate is good at but the recovery of the chromium including reduction of excess Cr(VI) is a bit of extra work.
.

Hilski - 4-10-2006 at 07:23

Quote:
Originally posted by guy
Is the H2O2 method too expensive for you guys or something? Its the easiest cleanest method. I've tried it and it works perfectly and its quite rapid also.


I have been looking for some (stronger than 3%) H202 since I saw your post, but haven't been able to find it locally yet. I definately want to give that method a try, and experiment with several of the other suggestions made by the folks in this thread. What conc of H2O2 have you used sucessfully in this reaction?
Quote:
Or you could do a little experimentation and try a direct route from heating MnO2 with excess (NH4)2SO4 and crystallizing the reaction product. The reaction is messy in the stoichiometry sense, as the 'NH4' goes off as water, nitrogen, ammonia, and sometimes a trace of NOx. You may need to add extra ammonium sulfate to the workup solution to have enough for the mixed Mn(II)NH4 salt (not an alum, that takes Mn(III)). The mixed Mn(II)NH4 sulfate can be isolate by crystallization; it is more soluble than ammonium sulfate hot, but less when cold, so concentrating until ammonium sulfate starts to crystallize, filtering hot, then cooling will let you get it out of solution. Yes, there's Mn left in solution, just let that evaporate and toss it into the next batch of MnO2 + (NH4)2SO4 before heating. Or you may be able to get away with just making sure there's enough ammonium sulfate in solution, as the procedure you mentioned says to use an excess of it, and taking the filtered solution of the cooled melt.


I might try this one as well, although I will have to get (or make) a lot more (NH4)2SO4 than I have right now in order to do it.

Man this thread has really got me interested in all the different things one can do with Mn compounds.

Thanks.

[Edited on 4-10-2006 by Hilski]

guy - 4-10-2006 at 11:46

I used 3%. But looking at the scale of your experiment, it might be ineffecient. Though I can get it for about $0.99 a bottle.

Hilski - 4-10-2006 at 23:05

Quote:
Originally posted by guy
I used 3%. But looking at the scale of your experiment, it might be ineffecient. Though I can get it for about $0.99 a bottle.


Well, I can get 3% very cheaply also, in quarts or gallons even. Am I correct in assuming that MnO2 and H2O2 are used in equimolar amounts in 1M H2SO4? And is the Mn2+ converted directly to H2SO4? That would be nice as it would save steps in my case.

Thanks

not_important - 4-10-2006 at 23:47

MnO2 + H2O2 + H2SO4 => O2 + 2H2O + MnSO4

Hilski - 5-10-2006 at 04:31

Quote:
MnO2 + H2O2 + H2SO4 => O2 + 2H2O + MnSO4


Thank you. I was hoping that was the case.

not_important - 5-10-2006 at 19:48

It's actually more complex than that, but it is a starting point.

I believe that the best way to do that would be to put the MnO2 into the excess H2SO4, warm it, and slowly add the peroxide until all the MnO2 had dissappeared, then add a little more, bring to a boil for a few minutes. Finish by adding enough 'ammonium carbonate' (or perhaps urea while keeping at boiling) to precipitate about 1 to 2% of the manganese while stirring well, let cool, and filter. The MnCO3 will capture some of the impurities.

12AX7 - 6-10-2006 at 06:35

Is FeCO3 much less soluble than MnCO3? (I get the feeling they may form a solid solution.) If so, that could take care of iron impurities well (especially important with pottery grade stock).

Tim

not_important - 6-10-2006 at 21:22

I believe it would be better to drop out Fe(OH)3 / FeO(OH). You could use fresh Mn(OH)2 to do that, adding a little aqueous ammonia as well, but the process must get the iron into the 3+ state. In that case you would want a small amount of MnO2 left, and boil for awhile.

I have these Ksp

FeCO3 2,5 x 10-11
MnCO3 6.3 x 10-11
Mn(OH)2 4,0 x 10-14
Fe(OH)2 ~1 x 10-13

Fe(OH)3 2,8 x 10-38


My pottery grade MnO2 doesn't seem to have much iron in it, maybe I should try to determine how much actually is there.

S.C. Wack - 8-10-2006 at 16:42

MnO2 from batteries may be low Fe, especially of course if from a brand that uses CMD rather than pyrolusite.

But while leafing through the 15 chemistry books that I downloaded from Google when they put their books up, I noticed an interesting preparation from pyrolusite:

http://books.google.com/books?vid=0IMFP4H-FFYEBB4t1Zak&i...

Hilski - 8-10-2006 at 20:33

That was a rather interesting read.

I decided to just dissolve the MnO2 in HCl and precipitate with Na2CO3. It was easy and worked just like I was told it would. The only real concern is how to deal with the chlorine. I just did the reaction outside, so it wasn't too big a deal, as long as I stayed some distance away from it until it was mostly finished.

After filtering the resulting solution a few times, I was left with a very pink, clear liquid. I made a saturated Na2CO3 solution and added it slowly to the pink solution. There was a lot of fizzing and foaming for a while. Once it finally stopped, the manganese carbonate started precipitating pretty quickly, and the solution turned a tan color and became a 'slush'. I havent converted it to MnSO4 yet, but judging by the MnCO3 I was able to get, it looks like I will have a decent amount of the sulfate when I am done.

guy - 8-10-2006 at 22:00

It turns brown because the air oxidizes it to MnOOH (manganese III oxyhydroxide).

not_important - 8-10-2006 at 22:11

Quote:
Originally posted by Hilski6(snip)...

After filtering the resulting solution a few times, I was left with a very pink, clear liquid.


It really is worth while investing in some good filtering gear. Do it before your government restricts access.
Quote:
I made a saturated Na2CO3 solution and added it slowly to the pink solution. There was a lot of fizzing and foaming for a while. Once it finally stopped, the manganese carbonate started precipitating pretty quickly, and the solution turned a tan color and became a 'slush'. I havent converted it to MnSO4 yet, but judging by the MnCO3 I was able to get, it looks like I will have a decent amount of the sulfate when I am done.


Excess acid, it sounds like. If you are after MnCl2 for further processing, you could do the following )next time) :

After the reaction between MnO2 and Hcl seems to have stopped, warm the solution add a bit more MnO2 and stir; repeat until no further reaction seems to happen. Note that 'warm' means warm, not put it on a red hot heater and try to melt the beaker/pan. After than add still a little more MnO2, and bring the solution to a gentle boil, then cool and filter. This is especialy true when using MnO2 from batteries, as that generally has carbon in it which makes it difficult to tell when all the MnO2 is gone. When there is an excess of MnO2 then you've used most of the HCl up.

The MnCO3 is white, but turns tan from oxidation. Not likely to be a problem in your case.

When using battery MnO2, it might be a good idea to wash the MnCO3 with aqueous ammonia to remove any zinc that may have been in the MnO2.

Hilski - 9-10-2006 at 18:02

Thank you guys for all the tips. All this info will definately come in handy for future projects.

MnSO4 ---> Mn2O3

roamingnome - 21-4-2007 at 13:10

variation on a theme here...

should i just buy the Mn2O3 probably, but i have MnSO4 stock to work with

the question is can the sulfate be subsituted for the nitrate in the catayst prep?

This catalyst was prepared by precipitation as follows. A solution of Mn((NO3)2)4H2O (20.8 g) in 100 ml of water was slowly admixed with a solution of KOH (11.2 g) in 100 ml of water in a well stirred beaker. The slurry was continuously stirred and heated to 80.degree. C. for one hour. The product was filtered off, washed with 1 liter of hot water and dried at 110.degree. C. overnight. This material was calcined in air at 650.degree. C. for 4 hours. The material was analysed by XRD and found to be pure Mn2O3.

Journal of Organic Chemistry, vol. 50, No. 17, 1985, pp. 3143-3148
Process for the catalytic preparation of -butyrolactone having the general formula
The process as claimed in claim 5, wherein the carboxylic acid is acetic acid and the olefin is ethylene.



it appears that from
http://www.drycleancoalition.org/download/tn_MnO4_injections...

2KMnO4 + 4H2O2 ---> 2KOH + Mn2O3 + 3H2O + 4O2



[Edited on 21-4-2007 by roamingnome]

not_important - 21-4-2007 at 17:08

Quote:
Originally posted by roamingnome

the question is can the sulfate be subsituted for the nitrate in the catayst prep?


The main reason to use nitrates, besides their often greater solubility, is that nitrate is less likely to carry over on the ppt, and that which does decomposes cleanly while sulfate doesn't and may result in traces of sulfate in the ignited catalyst. How important this is varies all over the place, it's sort of try it yourself and see. The industrial books regarding catalyst prefer nitrates in most cases, but in some case you can add sulfate - V2O5 for making SO3 is an example.

roamingnome - 24-4-2007 at 11:02

thanks, manganese is continually amazing me with all its valance states....

from the latest bout of patent scouring, they say that the valance state of the metal is one of the important parts first off.... so 3+ here i come

and from further reading Mn(III) acetate can be used to treat the acetic acid as well forming a radical... that acetate is crafty stuff, ill tell you that....
---------------------------------

Regeneration, pretreatment and precipitation of oxides of manganese
http://www.freepatentsonline.com/20040018936.html

Sedit - 11-7-2009 at 09:01

I just wanted to mention something here. It seems pretty accepted that MnO2 will not react with H2SO4 to produce MnSO4 yet I have used a direct method with battery MnO2 for a little while now and just repeated it after someone reported low yeilds to me.

When H2SO4 is heated to is boiling point it decomposes into H2SO4 = H2O + SO2 + 1/2 O2 IIRC.
SO2 can react with MnO2 to yeild MnSO4.

I started doing this after attempting the BnO synthesis that was put forth by Neograviton which involves boiling MnO2 in Sulfuric acid until white fumes are let off and leaving it for a little bit to react. I noticed that the Brown/black leftover mixture would precipitate as a white MnSO4 after boiled in H2O.

Whatever the mechanics of it maybe when you boil H2SO4 with MnO2 till its so concentrated that it fumes and this is left to react for a few minutes then a large amount of MnSO4 can be extracted right away without the need to convert to the chloride then to the carbonate or whatever process one uses. I feel that the concentrate H2SO4 must reach its boiling point before the reaction will proceed though.

chloric1 - 11-7-2009 at 09:45

If the 3% peroxide route seems wastefull of money/resources, don't overlook oxalic acid. Oxalic acid is available as rust remover or wood bleach. Simply mix your MnO2 with the appropiate amount of dilute sulfuric acid and heat and add oxalic acid. The oxalic acid should be converted to carbon dioxide but don't overlook the possibility of carbon monoxide generation. Hence this needs to be done outside/fume hood.
No sodium
No SO2
No problem

P.S. It might be a good idea to make a separate portion of Manganese carbonate and add to the final solution to remove any iron by sedimentation.

Saerynide - 11-7-2009 at 21:09

Interesting reads, but I think if you want to make so much, the most economical way is to just buy fresh MgSO4. If you live in the US, go to the healthcare section in Target and you'll find huge cartons (about 1 kg i think) of pure MgSO4, and its only a couple dollars.



[Edited on 7/12/2009 by Saerynide]

497 - 11-7-2009 at 21:33

Quote:
If you live in the US, go to the healthcare section in Target and you'll find huge cartons (about 1 kg i think) of pure MgSO4, and its only a couple dollars.


Need a little more coffee this morning? :P
In case you hadn't noticed, its Mn not Mg they're talking about.. not quite the same thing..

Saerynide - 12-7-2009 at 19:32

Oops my bad - dunno what I was thinking. Sorry :D

bilcksneatff - 15-7-2009 at 07:18

I was searching on this topic this morning, and found something. I'm not sure how much you guys trust Wikipedia, but this is what I found at http://en.wikipedia.org/wiki/Manganese_dioxide .

"One of the two chemical methods starts from natural manganese dioxide and converts it with dinitrogen tetroxide (N2O4) and water to manganese(II) nitrate solution, which is purified and after evaporation of the water a cristaline solid forms. At temperatures of 400°C the reverse reaction releases the N2O4 and manganese dioxide is formed.[4]

MnO2 + N2O4 → Mn(NO3)2
Mn(NO3)2 → MnO2 + N2O4"

The Manganese(II) nitrate could be converted to sulfate by precipitating MgCO3 with soda and then reacting that with sulfuric acid.

[Edited on 15-7-2009 by bilcksneatff]

[Edited on 15-7-2009 by bilcksneatff]

driedfloral - 7-2-2010 at 13:12

You may be able to find 27% H2O2 in the pool supply stores or in hardware store in pool section, I was able to get 1 gal for I believe $20 I know it wasn't over $30

chief - 7-2-2010 at 13:56

Above "gsd" wrote:
================
Direct Reaction of MnO2 with SO2 ( if you can get hold of SO2)
================

How fast/good does this reaction proceed ?
==> It could be a route towards H2SO4 from plaster (CaSO4; SO2 somehow obtainable by roasting), the SO2 --> SO3 -step would be no problem any more ...
==> MnSO4 would be electrolyzed for the acid, and pure Mn-Metal (which is possible too) ...

===============

franklyn - 28-4-2010 at 18:54

Direct combination of MnO2 and SO2 gives excellent yield and purity.
Sulfur burns readily if lit to give SO2. Available as 90 % + 10 % clay
Espoma brand , from your lawn and garden center. Purer more cheaply
available online _ http://www.dudadiesel.com/all_chemicals.php
These papers refer to gasing solid MnO2 with SO2 with only 33 - 40 %
yield of MNSO4. No problem just dissolve in water and again gas the dried
MnO2 filtrate _ http://pubs.acs.org/doi/abs/10.1021/es60120a014
also _ http://pubs.acs.org/doi/abs/10.1021/es60148a001
A better way is to gas a water slurry as shown in the second half here _
http://www.youtube.com/watch?v=2gXByJkg0iY
You could just buy MnSO4
http://www.starnursery.com/fertilizers/fertilizer-supplement...
If it needs to be purified _ http://www.youtube.com/watch?v=BLJgBSrhZI8
Beware and be alert and suspicious of Ebay purchases - a very good point ! -
http://www.youtube.com/watch?v=a5XZCZy3CvY

.

The WiZard is In - 29-4-2010 at 09:12

Quote: Originally posted by gsd  
1) The direct reaction of H2SO4 with MnO2 does not happen


I haven't kitchen tested this -

Pradyot Patnaik
Handbook of Inorganic Chemicals
McGraw-Hill 2003

Manganese (II) sulfate is prepared by prolonged heating with
any manganese salt with concentrated sulfuric acid. The
compound produced commercially from pyrolusite (MnO2)
or rhodochrosite (MnCO3). Either mineral is dissolved in
sulfuric acid and the solution evaporated.

NB This produces the tetrahydrate. Gentle heating produces
the monohydrate.

Manganese (II) sulfate also may be produced by the action
of sulfur dioxide with manganese dioxide. [Find an old
refrigerator.]

Doktor Klawonn - 20-11-2010 at 06:46

Quote: Originally posted by The WiZard is In  
...Manganese (II) sulfate also may be produced by the action
of sulfur dioxide with manganese dioxide. [Find an old
refrigerator.]


You can use an acidic solution of sulfite instead of SO2. It works quite well. I used it to make manganese carbonate MnCO3 form spent battery crud. See this video.

Dr.K.

blogfast25 - 20-11-2010 at 09:28

Search and yee shall find. Here’s thread of mainly woelen and me doing work on MnO2 + H2SO4:

http://www.sciencemadness.org/talk/viewthread.php?tid=11309#...

MnO2 and Mn2O3 both react with conc. H2SO4 (even with 50 % H2SO4 – draincleaner) to form a wonderful green/red-burgundy sulphato complex of Mn2(SO4)3 (Mn [+III]), which dissolves easily in dilute H2SO4 as burgundy Mn2(SO4)3 (it’s far more stable than you might think).

When a chloride is added to a Mn3+ solution, very temporarily MnCl3 is formed which, very similarly to MnCl4, breaks down to MnCl2 and Cl2 by oxidising a third of the chloride.

So yes, it’s possible to avoid the nasty chlorine factory from MnO2 + HCl by treating the oxide (or trioxide [+III]) with fairly concentrated H2SO4 but adding chloride to that reaction product releases a third of the chlorine as elemental chlorine: Mn3+ + 3 Cl-  MnCl2 + ½ Cl2. There’s probably a chlorine-free route to reduce the Mn3+ to Mn2+ but I can’t recall it right now…

Sedit - 20-11-2010 at 10:07

Quote: Originally posted by The WiZard is In  

I haven't kitchen tested this -

Pradyot Patnaik
Handbook of Inorganic Chemicals
McGraw-Hill 2003

Manganese (II) sulfate is prepared by prolonged heating with
any manganese salt with concentrated sulfuric acid. The
compound produced commercially from pyrolusite (MnO2)
or rhodochrosite (MnCO3). Either mineral is dissolved in
sulfuric acid and the solution evaporated.

NB This produces the tetrahydrate. Gentle heating produces
the monohydrate.

Manganese (II) sulfate also may be produced by the action
of sulfur dioxide with manganese dioxide. [Find an old
refrigerator.]


This is exactly what I was stating above even though the general consensous is that the reaction does not happen im sure it does and I have even recrystalized the MnSO4 to asure myself this was indeed what I had. I found out just by testing Neogravitons Benzaldahyde route which involved refluxing MnO2 and H2SO4 with Toluene. In the end MnSO4 precipitates out of the solution as the Toluene is added. Further test showed it to be effective route but I question the yeilds. Since my Mn came from batteries im not sure how much was carbon and how much was Mn compound so it makes determination of yeilds impossible. I plan on repeating it very soon so I suppose it wouldn't be to hard to give rough estimates on it.

I really believe that the formation is the result of concentrated Sulfuric decomposing into SO2 forcing the reaction but I have no real way of proving it other then the fact that little to no reaction seems to happen until the sulfuric was concentrated enough to start producing fumes. These SO2 fumes are what I feel are causing it to react.

chen - 7-12-2010 at 23:34

Hey, I tried a reaction similar to the one that guy talked about, that is, I mixed a small amount of MnO2 with dilute sulfuric acid and added an equal amount diluted (3%) hydrogen peroxide and came up with an orange solution. This is manganese sulfate right, with some iron contaminants? The source of the MnO2 was a pottery store.

blogfast25 - 9-12-2010 at 09:44

Quote: Originally posted by chen  
Hey, I tried a reaction similar to the one that guy talked about, that is, I mixed a small amount of MnO2 with dilute sulfuric acid and added an equal amount diluted (3%) hydrogen peroxide and came up with an orange solution. This is manganese sulfate right, with some iron contaminants? The source of the MnO2 was a pottery store.


Orange? Smacks of a peroxo complex...

ScienceSquirrel - 9-12-2010 at 10:04

Try adding some sodium hydroxide solution to a small sample.
Whitish precipitate, no iron; red brown precipitate, lots of iron.

not_important - 9-12-2010 at 10:22

Note that the method of prolonged heating of H2SO4 and MnO2 often was done with a long ramp-up to dull red heat, then cooling and extracting with cold water. This both drives off excess H2SO and convert iron compounds to Fe2O3 which remains as part of the insoluble leftovers (SiO2 is another portion).

blogfast25 - 9-12-2010 at 13:32

Below, left, a tube with a weeks old solution of Mn2(SO4)3, obtained dissolving MnO2 into 50 % H2SO4 (nothing else). Below, right, an empty tube:



Below: after transferring half of the left hand tube into the right and adding weak peroxide solution to the right hand tube’s half:



It appears that in these conditions, Mn (III) is oxidised to Mn (IV). In the presence of Cl-, covalent MnCl4 would form which would shed half of its chlorine almost immediately via: MnCl4 --- > MnCl2 + Cl2.

No orange on this occasion…

blogfast25 - 10-12-2010 at 07:54

Well, I have to self-correct because I was wrong in the above post. Thinking about it last night I recall there being some very small bubbles of clear gas being evolved when adding the peroxide to the Mn3+ solution. That would indicate O (-I) --- > O (0) and Mn (III) --- > Mn (II).
I added some 3 % peroxide to the left hand tube (Mn3+) of the last photo. Again it clears up and fine bubbles can be seen with the naked eye.
The right hand side tube is the right hand side tube from the last photo but with strong HCl added: no MnCl4 formed as I erroneously predicted… No oxidation of Cl- occured either, normally expected in the presence of Mn (IV)…



Below I added 5 M NaOH to the left hand tube, during neutralisation (of the excess H2SO4) brown-black MnO2 precipitates (the two layers are due to insufficient mixing), presumably due to excess H2O2.



So here Mn3+ oxidised hydrogen peroxide in acid conditions to oxygen and water and peroxide in alkaline conditions oxidised the Mn2+ back to MnO2 (Mn [IV])…

A recent paper I found on the net suggested revisiting Mn3+ solutions as oxidising titrant solutions: it is much more robust than generally assumed. But solid Mn3+ compounds seem very rare: I recall reading about an unstable Rb Mn (III) alum but not much more…

[Edited on 10-12-2010 by blogfast25]

chen - 17-12-2010 at 16:28

Okay, I brought some pictures this time.


Several days after combining dilute sulfuric acid, dilute peroxide and manganese dioxide, a precipitate formed:





The second photograph is blurry but captures the color of the precipitate okay. The precipitate looked a more pinkish/beige irl.

I don't know a whole lot about manganese precipitates but after snooping a bit I found this site: http://www.chemguide.co.uk/inorganic/transition/manganese.ht...
which describes it as Mn(H2O)4(OH)2. There was also a darker precipitate which might be Mn(III) oxide or unreacted MnO2.

I'm sure you all know that Hydrogen peroxide mixed with Manganese Dioxide catalyzes the decomposition of Hydrogen Peroxide. I looked in an old science journal from the 1940's and it listed a possible mechanism for the catalytic action:

MnO2 + H202 + 2H+ -> Mn(2+) + 2H2O + O2
Mn(2+) + 2H2O2 <-> Mn(OH)2 + 2H+
Mn(OH)2 + H2O2 -> MnO2 + 2H2O

I got the idea that adding sulfuric acid would make the equilibrium in step 2 favor formation of Mn(2+) but had no idea what it would form after that. I have no clue how the hydroxide complex formed (if it even is a hydroxide complex), especially in such an acidic solution.


I will test the remaining clear solution for iron. I think the first thing I am gonna do is redo this experiment and take pics of the solution immediately after mixing and every day afterwards. I also have a scale now so I can test the mystery ppt for weight.

[Edited on 18-12-2010 by chen]

[Edited on 18-12-2010 by chen]

gsd - 17-12-2010 at 19:00

When I made MnSO4 by this method, I followed following procedure:

Reagents : MnO2 - low grade ore containg about 42% MnO2 and 40% CaCO3
H2SO4 - 95%
H2O2 - 53 %

In 1 lit beaker provided with O/H stirrer, a slurry of 100 gm ore and 200 ml water was made and 97 gm H2SO4 added in portions. The amount of acid stoichiometrically corresponds to CaCO3 & MnO2 present in the ore. The temperature rose to about 75 Deg C.
29 gm of H2O2 was diluted with water ( 1 :1 ) and added dropwise thru' a separating funnel over a period of about 30 min. The reaction was quite vigorous and temp of about 65 to 70 sustained throughout the addition. The stirring continued for another hour and the reaction mix allowed to settle. A clear but slightly reddish pink colour solution was obtained.

The point to be noted here is that this reaction occurs only on acidic conditions. Otherwise MnO2 just catalyses the decomposition of H2O2 without getting it self reduced to MnO.

Gsd

chen - 18-12-2010 at 15:22

I added some hydroxide to the supernate solution and a pale brown complex precipitated first; a dark brown complex and a whitish complex precipitated after the pale brown complex. I believe there was a layer of clear liquid underneath the light brown and dark brown precipitates, suggesting that the brown precipitates might be less dense than a liquid phase. At the very bottom of the film canister was the white complex. It's hard to say since I did the whole thing in a film canister.

I also took a cotton swab of the precipitate of the original solution and exposed it to air. It did not change color and stayed the original beigish pink.


Mass balancing etc. later. Will use 1 oz. clear glass spice jars instead of film canisters.

S.C. Wack - 28-12-2010 at 04:02

Quote: Originally posted by S.C. Wack  
MnO2 most definitely dissolves in hot conc. H2SO4 with vigorous evolution of oxygen. This is not debatable. This says nothing of the usefulness of drain cleaner and crap pottery MnO2 in producing good MnSO4, however.


MnO2 from Rayovac alkaline batteries readily reacted with excess colorless technical 96% sulfuric acid seconds after applying minimal heat from a blowtorch to a test tube containing them. It set to a black paste after a minute. Extraction 3x with water dissolved the majority of the paste. The majority of what remained didn't dissolve in HCl, likely because it's carbon.

The clear filtered extract has a dark violetness. Heating at 250F causes it to look exactly like MnSO4 solution, over a precipitate of grey-black MnO2. On oven evaporation until the color is gone from the acid, the filtered MnSO4 in H2SO4 precipitated a buff colored crystalline crust. The color is not as dark as in chen's picture. The acid is quite strong at this point and the dry, crystalline Mn cpd. retains considerable acid after washing with ethanol. It dehydrated to white, is plenty soluble in water, and in solution it looks like MnSO4. But the pH is 2.5, and I don't really care about recrystallizing for MnSO4, so the Mn was precipitated with baking soda as a white paste. This gave a tan powder on drying at 200F, weighing the same as the original battery sample.

I might try the oxalic acid method.

blogfast25 - 28-12-2010 at 08:11

Quote: Originally posted by S.C. Wack  

I might try the oxalic acid method.


Explain?

S.C. Wack - 28-12-2010 at 13:23

Like H2O2, it's another old redox method from the analysis world; usually described as or based on "the method of Fresenius and Will" References for that sort of MnO2 analysis to 1900: http://books.google.com/books?id=oJYZAAAAMAAJ&pg=PA114

blogfast25 - 28-12-2010 at 13:34

I see.

NurdRage - 28-12-2010 at 14:00

@blogfast

Go to 1:35 of this video: http://www.youtube.com/watch?v=2gXByJkg0iY

There i detail the oxalic acid approach for making manganese sulfate. Its very easy if you can get your hands on oxalic acid and sulfuric acid.

blogfast25 - 29-12-2010 at 07:16

Yep, that makes sense: I've used standardised oxalic acid to standardise KMnO4. It's a neat trick, I've gotta say...

[Edited on 29-12-2010 by blogfast25]

Tartaric acid reducing manganese dioxide

symboom - 29-11-2016 at 14:29

Is it possible to reduce MnO2 with tartaric acid
Just as manganese oxalate then adding sulfuric acid

Copper tartrate is soluble so I figure that manganese tartrate might be

[Edited on 29-11-2016 by symboom]

Random - 2-12-2016 at 12:37

MnO2 mixed with non-reactive acid for example vinegar along with bisulfite will be reduced and the impurities will be left. After that it can be precipitated with Na2CO3 and the resulting carbonate will gladly dissolve in sulfuric acid.

clearly_not_atara - 2-12-2016 at 13:09

I know this sounds utterly silly, but why not just try:

MnO2 + SO2 >> MnSO4

SO2 is a much less annoying gas than HCl or Cl2 or any such thing. Aqueous SO2 is not hard to handle at all.

EDIT: Orgsyn likes using SO2 in H2SO4, which at the very least is nicer than HCl in H2SO4 or any such thing:

http://www.orgsyn.org/demo.aspx?prep=cv2p0315

[Edited on 2-12-2016 by clearly_not_atara]

Dmishin - 3-12-2016 at 03:16

I several times used citric acid as reducing agent, with good results.

My procedure was the following:

- Wash with water and dry the black gunk form the used batteries.
- Put dried gunk to the reaction vessel, add 36% H2SO4 to it in stoichiometric proportion, assuming it is 80-90% MnO2.
- Add citric acid, not more than 2-3 grams, and stir. At first, there is no reaction, but after 15-20 seconds it starts hissing and bubbling and temperature rises, reaction gradually becoming more and more intensive. Add following portions of citric acid, when reaction slows down.

I think that in this reaction citric acid is oxidized completely to CO2 and H2O. The reaction is very exothermic, so cold water bath might be required. Adding a lot of citric acid at once can cause runaway reaction.

nezza - 4-12-2016 at 03:50

In acid conditions Mn2+ is the stable species. Adding Hydrogen peroxide and dilute Sulphuric acid to Manganese dioxide will indeed give Manganese(II) sulphate in solution. The solution needs to kept acidic to avoid oxidation to Mn3+ and Mn(IV). The addition of NaOH to a solution of MnSO4 gives a whitish precipitate of Mn(OH)2 which rapidly oxidises in air to brown Mn(OH)3. The presence of peroxide will immediately oxidise it to Mn(OH)3 and MnO2 and decompose any remaining peroxide. Basically Mn(II) requires an acidic and reducing environment to be stable.

Random - 11-12-2016 at 09:12

Quote: Originally posted by Dmishin  
I several times used citric acid as reducing agent, with good results.

My procedure was the following:

- Wash with water and dry the black gunk form the used batteries.
- Put dried gunk to the reaction vessel, add 36% H2SO4 to it in stoichiometric proportion, assuming it is 80-90% MnO2.
- Add citric acid, not more than 2-3 grams, and stir. At first, there is no reaction, but after 15-20 seconds it starts hissing and bubbling and temperature rises, reaction gradually becoming more and more intensive. Add following portions of citric acid, when reaction slows down.

I think that in this reaction citric acid is oxidized completely to CO2 and H2O. The reaction is very exothermic, so cold water bath might be required. Adding a lot of citric acid at once can cause runaway reaction.


What is the reaction behind this?

symboom - 11-12-2016 at 14:47

Quote: Originally posted by Dmishin  
I several times used citric acid as reducing agent, with good results.

My procedure was the following:

- Wash with water and dry the black gunk form the used batteries.
- Put dried gunk to the reaction vessel, add 36% H2SO4 to it in stoichiometric proportion, assuming it is 80-90% MnO2.
- Add citric acid, not more than 2-3 grams, and stir. At first, there is no reaction, but after 15-20 seconds it starts hissing and bubbling and temperature rises, reaction gradually becoming more and more intensive. Add following portions of citric acid, when reaction slows down.

I think that in this reaction citric acid is oxidized completely to CO2 and H2O. The reaction is very exothermic, so cold water bath might be required. Adding a lot of citric acid at once can cause runaway reaction.



I wonder if this procedure of using citric acid instead of oxalic acid could be replicated with potassium bitartrate or ascorbic acid

Dmishin - 24-12-2016 at 16:28

Quote: Originally posted by Random  
Quote: Originally posted by Dmishin  
I several times used citric acid as reducing agent, with good results.
...


What is the reaction behind this?


Produced gas has no smell and extinguishes flame, thus it is CO2 for sure.
There are also no strong smells; particularly, no smell of acetic acid or acetone.
From this, I suppose that citric acid is oxidized completely, to water and CO2:

C6H8O7 + 9 MnO2 + 9 H2SO4 = 9 MnSO4 + 13 H2O + 6 CO2

This also matches well with the amount of citric acid required.

exodia - 24-12-2016 at 21:51

So seeing all this different reactions, how do you think iron contaminants will behave with them?

1. with the solution of MnO2 + H2SO4 and Toluene, will the iron sulphate precipitate upon toluene addition as well?

2. in the procedure given by gsd, adding H2O2 dropwise, will the iron contaminants stay precipitated and the MnSO4 will stay in solution?

3. and with the citric acid addition, will there be any precipitate, iron maybe? (will it stay in solution mixed with the manganese ions?)

Random - 29-9-2024 at 13:51

Quote: Originally posted by exodia  
So seeing all this different reactions, how do you think iron contaminants will behave with them?

1. with the solution of MnO2 + H2SO4 and Toluene, will the iron sulphate precipitate upon toluene addition as well?

2. in the procedure given by gsd, adding H2O2 dropwise, will the iron contaminants stay precipitated and the MnSO4 will stay in solution?

3. and with the citric acid addition, will there be any precipitate, iron maybe? (will it stay in solution mixed with the manganese ions?)


You see. Look how unexplored Chemistry is.