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Author: Subject: MnO2 -> MnSO4; What is the best route?
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[*] posted on 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.





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

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




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[*] posted on 28-12-2010 at 08:11


Quote: Originally posted by S.C. Wack  

I might try the oxalic acid method.


Explain?
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[*] posted on 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
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[*] posted on 28-12-2010 at 13:34


I see.
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[*] posted on 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.
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[*] posted on 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]
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[*] posted on 29-11-2016 at 14:29
Tartaric acid reducing manganese dioxide


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



If you're not part of the solution, you're part of the precipitate.
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[*] posted on 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?
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[*] posted on 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
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[*] posted on 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.
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[*] posted on 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?)
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