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Author: Subject: MnO2 -> MnSO4; What is the best route?
not_important
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[*] posted on 4-10-2006 at 23:47


MnO2 + H2O2 + H2SO4 => O2 + 2H2O + MnSO4
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[*] posted on 5-10-2006 at 04:31


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


Thank you. I was hoping that was the case.
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[*] posted on 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.
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[*] posted on 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




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[*] posted on 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.
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[*] posted on 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...
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[*] posted on 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.
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[*] posted on 8-10-2006 at 22:00


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



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[*] posted on 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.
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[*] posted on 9-10-2006 at 18:02


Thank you guys for all the tips. All this info will definately come in handy for future projects.
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[*] posted on 21-4-2007 at 13:10
MnSO4 ---> Mn2O3


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]
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[*] posted on 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.
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[*] posted on 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
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[*] posted on 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.





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[*] posted on 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.




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[*] posted on 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]




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[*] posted on 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..
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[*] posted on 12-7-2009 at 19:32


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



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[*] posted on 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]
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[*] posted on 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
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[*] posted on 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) ...

===============
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[*] posted on 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

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[*] posted on 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.]
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[*] posted on 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.
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[*] posted on 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…
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