I have been searching to find out what happens when you add H2SO4 to MnO2.
It seems as if the answer to this is only available on sciencemadness and no other site has it.
ANSWER : nothing happens if you use dilute H2SO4. I tried it and yes, nothing happens.
You need a reducing agent added like oxalic acid or H2O2 and you get Manganese(II) Sulfate and water and O2 (in the case of H2O2 added).
So, Manganese(IV) Sulfate doesn't form? Doesn't exist?bfesser - 15-12-2013 at 07:07
: nothing happens if you use dilute H2SO4. I tried it and yes, nothing happens.
About this; I've always wondered if MnO2 could be purified by adding dilute H2SO4 to it (to turn the impurities into soluble compounds).
Also couldn't Manganese (IV) sulfate be made by the action of H2SO4 on MnCl4?
[Edited on 15-12-2013 by bismuthate]woelen - 15-12-2013 at 09:44
All these references are about very dilute solutions in quite strongly concentrated H2SO4 (0.005 M in manganese(IV), while being 9M in H2SO4). I
indeed think that Mn(SO4)2 as a pure compound does not exist. In concentrated acid, soluble manganese(IV) species certainly exist, but I strongly
doubt one can speak of free aquated Mn(4+) ions. chornedsnorkack - 15-12-2013 at 10:38
If you increase your concentration all the way from 9M (about 60%?) H2SO4 to 100 % SO3 (remember, liquid above 17 degrees until it condenses into beta
or alpha), would the solubility of manganese increase all the way or would it go through maximum?
Can you get solid precipitates other than MnO2? And what would they be? Like
Mn(SO4)2?
Mn(HSO4)4?
Mn(S2O7)2?
Mn(HS2O7)4?
MnOSO4?
MnO(HSO4)2?
Something else?
[Edited on 15-12-2013 by chornedsnorkack]
[Edited on 15-12-2013 by chornedsnorkack]chornedsnorkack - 15-12-2013 at 14:10
If you increase your concentration all the way from 9M (about 60%?) H2SO4 to 100 % SO3 (remember, liquid above 17 degrees until it condenses into beta
or alpha), would the solubility of manganese increase all the way or would it go through maximum?
It should drop.
On closer examination, SO3 is planar trigonal. Therefore no dipole moment. And modest polarizability. The dielectric constant is quoted as 3,6 or so.
For comparison, sulphuric acid is in 80-100 range, like water.
If you are handling dilute solutions of sulphuric acid in sulphur trioxide, they should have high pH and low conductivity because the low dielectric
constant should be hostile to charge separation and electrolytic dissociation. Also sulphur trioxide should be a poor solvent for ionic salts.
Sulphuric acids should dissolve as molecules.
So if sulphuric acid is diluted with sulphur trioxide at high temperatures, over 35 degrees, then the solvent should gradually lose dielectric
constant and precipitate salts. If sulphuric acid is diluted below 35 degrees, then solid disulphuric acid should precipitate, binding sulphuric acid
and forcing any solutes to saturation.
What solid salts of manganese should precipitate from oleum as it is diluted with sulphur trioxide?vmelkon - 19-12-2013 at 18:23
I have this written in my old notes. It comes from a chemistry dictionary.
Mn(SO4)2, black crystals (MnSO4 in H2SO4 + KMnO4) readily hydrolysed.Metacelsus - 15-6-2014 at 14:10
I just did this 30 minutes ago (adding 98% sulfuric acid to hot manganese dioxide from batteries). I was hoping that it would form manganese (ii)
sulfate with liberation of oxygen. I was expecting a slow reaction, even with the heating. However, I got a very vigorous reaction that was barely
contained within my flask. It also created a fine mist of sulfuric acid.
Details: 250 g manganese dioxide (2.87 mol).
I was going to add 3 50 ml portions of sulfuric acid, but stopped after the first.
I estimate the manganese dioxide was at around 120 C when I added the acid.
Conclusions: The reaction of hot manganese dioxide with concentrated sulfuric acid releases oxygen and proceeds rapidly. The presumed product is
manganese (ii) sulfate.
I will add the rest of the acid very slowly and at a lower temperature once my lab is clear of sulfuric acid mist.Oscilllator - 15-6-2014 at 15:51
Cheddite it is interesting that that happened, since in my experience sulfuric acid does not react with MnO2 at room temperature. To create MnSO4 I
first reduced the Mn(IV) to Mn(II) with oxalic acid, then reacted the product with sulfuric acid to get manganese(II) sulfate. I guess the sulfuric
acid acted as a reducing agent here somehow. Does anyone know of other instances where that can happen?bismuthate - 15-6-2014 at 16:09
Cheddite Cheese, did you use pure MnO2 or the battery variety? Metacelsus - 15-6-2014 at 19:32
I used a purified version of the battery variety (washed to remove all compounds soluble in water and/or acetone).
The reaction definitely does not happen at room temperature.
I have reacted the rest of the manganese dioxide, and will soon filter off the insoluble impurities (mostly graphite). The solution is the pink color
of Mn(ii).
