Strange manganese dioxide (?)

Here's some pix:

Left: homemade MnO2, right: the bought 'high grade MnO2'



Test tube #1 with the bought 'high grade MnO2', after boiling with 50 % H2SO4



Test tube #2 with the homemade MnO2, after boiling with 50 % H2SO4




Poor colour representation in all cases due to the flash but the green and permanganate purple can clearly be discerned.

Is your "high-grade MnO2" really that light? It looks almost orange!

The purple color can be understood. It is the color of Mn(3+) in very strongly acidic solution with a coordinating ion. This even is a test for manganese.

I have a solid with manganese(III) in it, it is a phosphato-complex, NH4 [MnP2O7], ammonium pyrophosphatomanganate(III):



If this is dissolved in strong acid, a deep purple solution is obtained. I can imagine that your purple color is due to some sulphato complex given the very high concentration of sulphuric acid.

All these manganese-things also are puzzling me, it is not what I expect. The pale green color also is not what I expect, nor the grey stuff.

If you have KMnO4, then you could try a similar experiment. First make MnO2 from the KMnO4, rinse, and dry. Then add that to the 50% H2SO4 and repeat the experiment. In this way you are 100% sure that no other metals are in play.




EDIT by woelen: Changed link, so that it works again.

[Edited on 12-6-12 by woelen]

Unfortunately my home assay of the bought 'MnO2' (based on dissolution in HCl and re-oxidation and re-precipitation as MnO2) wasn't conclusive, some precipitate must have got lost along the way. Really only a titrometric method could save me here but I don't recall any (for Mn²⁺.

I'm aware of the fact the MnO2 is always somewhat oxygen deficient.

Later on today I'll alkalise the supernatant solutions and see if Mn2O3 (Mn III) precipitates or whether manganate (Mn VI) forms.

The difference in behaviour between the two oxides in boiling 50 % sulphuric acid would seem to suggest they are distinctly different species...

[Edited on 28-10-2008 by blogfast25]

Well, I dropped about 1 ml of the supernatant fluids in a few ml of 5 M NaOH and both behaved similarly. Firstly a reddish-brown colour change is seen but without precipitation. Then after cooling back to RT and on adding a few more ml of 5 M NaOH, slowly a whitish gelatinous precipitate forms. At first this precipitate is almost transparent (like freshly precipitated waterglass). No green colour was observed at any time.

I'm to assume the Mn2O3 in these conditions is either very lightly coloured or disproportionated to Mn(OH)2 and x (?).

But it shows at least no permanganate was present...

I'll repeat these tests, this time using dilute H2SO4 (say 1 M) instead of 50 %...

[Edited on 28-10-2008 by blogfast25]

@Woelen:

Great!

Here are the strange precipitates obtained after dropping some of the purple supernatant fluids into 5 M NaOH:



(Left #1, right #2) Not really what I expected either: doesn't really look like 'Mn2O3. x H2O', does it? The precipitates have already darkened slightly since yesterday. Most of yesterday you could practically read the newspaper through them!

And I found this titrimetric determination of manganese in steel:

AS/NZS 1050.14:1994:

3 PRINCIPLE The sample [steel] is dissolved in dilute sulfuric acid, any other acid additions necessary for complete dissolution being removed by fuming after oxidation, and interfering elements removed by precipitation using a zinc oxide suspension. The resultant solution is strongly acidified to prevent interference from cobalt and the manganese is oxidized to the permanganate ion by ammonium persulfate. The permanganate ion is titrated with ammonium iron(II) sulfate and potassium dichromate, using preoxidized sodium diphenylamine sulfonate as indicator.

Unfortunately half of the paper is available for purchase only...

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I boiled both precipitates in a steambath. Also in the photo: 3 g of the unknown oxide being boiled in a conical flask with 100 ml of 1 M H₂SO₄ (ooops - must clean that cooker again!):




Left and right: #1 and #2 after a few minutes of boiling. They've lost a lot of volume (water) but remain light in colour, not very indicative of Mn2O3. x H2O.




The remainder of #2 (see post above), after simply adding DIW to the rim of the tube but without mixing. The green Mn compound can be seen clearly. To the naked eye it's much lighter...




The now diluted supernatant liquid (above) of #2 was then carefully siphoned off into a small conical flask. This presumably is the relatively pure Mn (+III) compound. It didn't suffer hydrolysis after dilution (about (2:1).




The unknown oxide, when boiled with 1 M H₂SO₄ turned black almost immediately. I simmered for about 15 min, then cooled, settled and filtered. It filters difficultly (runs through the filter, at least much of it). The filter was treated with hot 32 w% HCl: the (truly) black precipitate dissolves with evolution of chlorine. Presumably it is MnO₂.

The filtrate is now decanting but I can already see the supernatant liquid to be clear and colourless: dilute MnSO₄?

Any ideas on what tests to perform on the green compound?

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The clear, now decanted filtrate precipitated a creamy, white precipitate when saturated Na₂CO₃ was added (after cessation of bubbles). Adding bleach to the precipitate and it turns brown/black. Presumably the clear filtrate contained MnSO₄ and the precipitate was MnCO₃, then oxidised to MnO₂ with hypochlorite.

So it's safe to say the unknown oxide reacts with dilute sulfuric acid according:

y Mn oxide + H₂SO₄ ---> x MnSO₄ + z MnO₂ (+ water?)

That would be consistent with Mn₂O₃ but apparently also with Mn₃O₄...

And another (but similar) method for titrimetric determination of managnese (free article)

It would be interesting to use this method to determine the proportions of Mn (II), Mn (IV) and Mn (III) obtained in various conditions of dissolution.

The method essentially calls for oxidising the Mn to permanganate with amm. persulfate, then adding a standardised excess of Fe²⁺ and back titrating the excess with standardised KMnO₄:

Mn (II) + persulfate ---> Mn (VII)
Mn (VII) + x Fe (II) ---> Mn (II) + 5 Fe (III) + (x-5) Fe (II)

(x-5)/5 Mn (VII) + (x-5)Fe (II) ---> (x-5)/5 Mn (II) + (x-5) Fe (III)



[Edited on 29-10-2008 by blogfast25]

Some more experiments...

1) On the white precipitate obtained from the purple solution + alkali:

Firstly, the precipitates (3rd pic, previous post) that had lost so much volume had reconstituted themselves overnight:




The swollen precipitates dissolve in cold, strong HCl but not eagerly and a clear solution was obtained. I didn't notice any chlorine... Also remarkable is the fact that when I neutralise this obtained clear solution and then add sat. sodium carbonate solution, no precipitate is formed. Only by adding more NaOH (5 M) does a gelatinous precipitate eventually show up. This is a clear indication that the solution does not contain much, if any, Mn²⁺.

Another portion of the precipitate was subjected to bleach (hypochlorite, at pH 13 - 14) and a steam bath: hardly any oxidation at all took place: there is no colour change to black, so typical when bleach is added to freshly precipitated Mn(OH)₂.


2) On the green precipitate:

After having siphoned off the purple liquid from both test tubes, I added 32 w% HCl to the green precipitates: both reacted very quickly, turning black, with generation of chlorine and a lot of heat, the test tubes became too hot to hold:




The black stuff (combined from both tubes) with added strong HCl and some heat quickly clears to MnCl₂ and chlorine (lots of it):




3) On the purple solution, suspected to contain Mn (+III):

Adding oxidisers:



Left: the purple solution as such.

Middle: after adding bleach. It slightly darkens. No precipitate. We're at pH ≈ 0 though...

Right: after adding weak H₂O₂. Goes clear.

Adding reducers:



Left: after adding strong HCl. It darkens considerably, no chorine though. MnCl₃? Which is allegedly dark coloured but should lose colour through decomposition to MnCl₂ on heating (like MnCl₄ does).

Edit: after heating the solution does clear up and chlorine is formed during that. That would strongly point to MnCl₃... Even a saturated NaCl solution (about 6 M, 32 w% HCl is about 10 M) added to the purple solution caused it to darken slightly, which disappeared on heating and the faint smell of chlorine could be observed. This would point to MnCl₃ being analogous to MnCl₄ in being more covalent than heteropolar and unstable at higher temps.

Middle: after adding bisulfite. Clears up completely. Presumably reduces Mn(III) to Mn(II).

Right: after adding strong Fe²⁺. Clears up. Presumably reduces Mn(III) to Mn(II). The residual green is due to excess Fe²⁺, IMHO.

Some more to follow.

I read that Mn₂(SO₄₃ forms alums with Rb and Cs, albeit not very stable ones. I wonder if it would be possible to pull a potassium alum from the purple solution... It'd be nice to have a solid Mn(III) bearing compound...



[Edited on 30-10-2008 by blogfast25]

Wow, you have done a lot of experimenting. Let's try to organize things a little...
- concentrated acid vs. dilute acid
- presence of chloride ion vs. presence of sulfate only
- use of reductor
- use of oxidizer
- ...
By organizing all things, a pattern may appear.

Some remarks which I can give already: the strong heat you observe by mixing the purple liquid with conc. HCl most likely is due to heat of hydration of conc. H2SO4. The purple liquid contains a lot of H2SO4 and hardly any water.

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I also have done some experimenting by now, this weekend I'll have more time.

Experiment 1: Add black MnO2 (laboratory grade material, not from some eBay seller or some pottery shop, but a real lab chem) to a 1 : 1 mix of conc. H2SO4 (also reagent grade) and distilled water. Heat this stuff.

The result of this experiment is no reaction at all. The black solid does not dissolve. I kept on boiling until white smoke was produced from the sulphuric acid. The solid does not dissolve.


Experiment 2: Use the same setup, but now add a small amount of solid Na2SO3, such that some of the MnO2 could be reduced. When this is done, then some SO2 is produced (pungent smell!!), but the liquid also turns bright purple. When the liquid is allowed to stand for a while, a nice bright purple liquid above a black solid appears. So, the purple material really is due to some lower manganese compound (most likely +3 oxidation state).

When water is added to the purple liquid, then it becomes turbid. After one day of standing, a black precipitate is formed, with a nice brown/red liquid above it. This color is due to Mn(3+) in aqueous solution at high concentration of water. This color is familiar to me.




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In another experiment I took MnSO4 and dissolved this in 3% H2O2 and added this to excess 5% NH3. A dark brown flocculent precipitate is formed. I'll let this settle to the bottom and clean this, such that next weekend I have some material to play with. This brown flocculent precipitate most likely is hydrous MnOx, with x < 2.

A few more experiments.

I made a larger amount of the homemade MnO2/boiling 50 % H2SO4 mixture (3 g oxide + 50 ml acid). The pic below gives a better view of the unknown green compound (but photographed it looks more blueish than green):




I then tried to wash it to get rid of the H2SO4 but on adding water the compound hydrolyses to a brown compound. Lighter than freshly precipitated MnO2, more like Fe2O3. Could be Mn₂O₃ or maybe Mn₃O₄:



So I've not been able to confirm/infirm the reaction of the green compound with HCl.


Taking a small amount of this suspension and adding 32 % HCl and it dissolves instantly to this w/o chlorine development - MnCl3?:




Heated it clears up with generation of chlorine:




What's disappointing is that the brown compound does not react at all with 1 M H2SO4, not even on heating. I tried twice with different ratios. I will try again with a slightly more concentrated H2SO4, perhaps 25 w%, instead of 1 M. If the stuff doesn't react with dilute H2SO4 then it can't be Mn2O3 or Mn3O4. For now I'm assuming it is (either) but that I'm not testing properly.


One other test I carried out was to add NaClO₃ (chlorate, sat.) to the purple liquid. No colour change was noted. It makes me think that the darkening I observed yesterday when adding bleach may simply have been due to the chloride content of the bleach: Mn (+III) species + 3 Cl^- ---> MnCl₃.

Summarising so far we've got (in all likelihood):

Homemade MnO₂ or unknown Mn oxide + 50 % H₂SO₄ + heat---> soluble purple Mn (+III) species + unknown, insoluble green Mn-based species

Unknown, insoluble green Mn-based species (in 50 w % H₂SO₄ + water ---> Mn₂O₃ or Mn₃O₄



[Edited on 31-10-2008 by blogfast25]

@Xenoid:

You took the words right out of my mouth!

I've been pondering based on these results so far about what the 'sulfato Mn(+III) complex' could be and the lack of any cation (apart from H₃O⁺ makes me think that the violet colour is indeed simply Mn₂(SO₄₃.

I've inadvertently obtained it by heating homemade MnO2 with 50 w% sulfuric acid, Woelen basically by dissolving MnSO4 in strong acid, followed by mild oxidation.

There remain nonetheless some questions. The nature of the blue-green compound not being the only one.

<i>"It dissolves in water to a violet solution (Mn3+) which deposits hydrated manganese dioxide on standing (dark-brown coloured)."</i>

Clearly not when the concentration of H₃O⁺ and/or SO₄^- is great enough: my purple coloured solutions doesn't budge even on considerable dilution...

I carried out one more experiment. On standing overnight, the tube (second photo, previous post of mine) had cleared: all precipitate had collected at the bottom. After decanting off the clear supernatant liquid, I added a few centimeters of approx. 25 w% H2SO4. Even before heating, a purple solution started to form at the bottom of the tube. On heating in direct flame the tube unfortunately 'burped' and most of the content ended up splattered over my lab. The centimeter or so that was left was a clear purple solution. I need to corroborate (with steam bath heating this time) this but it seems to confirm Mellor: freshly precipitated Mn oxides dissolve readily in H2SO4 (of the right concentration?), forming Mn₂(SO₄₃.

The question of what my strange looking Mn oxide really is still remains open too.

And now I'm seriously tempted to try and produce an Mn₂(SO₄₃ based alum, with K, Rb or Cs...

$$$$$$$$$$$

Is there no end to this weirdness??? In an attempt to corroborate the dissolution in 25 % H2SO4 of the brown hydrolysed product of the insoluble green compound, I transferred some of the green slurry (top pic, 7.46 post) into an Erlenmeyer and added a small amount of water to provoke hydrolysis. Instead the green compound dissolved - completely:



The colour is slightly more reddish that the previous purple solutions but for now I'm attributing this to concentration: in other words, it's still Mn2(SO4)3 but at a higher concentration, giving a slightly different hue. Only spectroscopy could tell that with certainly, of course.

My first thought was that it might be concentrated MnSO4 (and not Mn2(SO4)3) but adding 32 w% HCl makes it go quite dark and heating the dark solution makes it clear up to yellow, with evolution of chlorine. Mn(II) doesn't do that.

It would seem to indicate that concentration of the H2SO4 is really important: the liquid part of the green slurry was 50 w% H2SO4, diluting it a bit we're probably in the 25 % range, enough to dissolve it to Mn2(SO4)3...

I'm still nonetheless puzzled as to how a predominantly Mn (IV) compound (homemade MnO2) can be reduced to a relatively stable Mn (III) compound, seemingly in the absence of any reducing agent...

It should also be possible to determine the Mn³⁺ concentration by adding a standardised excess of Fe²⁺ and titrating the excess with standardised KMnO<sub>4</sub...

[Edited on 1-11-2008 by blogfast25]

I also have done some more experimenting. I found the following:

In conc. 96% H2SO4, the purple compound cannot exist. I added some of my solid NH4[MnP2O7] to conc. H2SO4. The purple solid quickly turns brown and when the acid is heated, I obtain a reddish brown precipitate in a colorless liquid. To this, I added water, dropwise. At a certain point, the very fine precipitate dissolves and a deep purple and clear liquid is obtained. When more water is added, then the color of this solution shifts from purple (like permanganate) to reddish purple and from there to red (something like shown in the picture above of the erlenmeyer) and on addition of even more water, the solution becomes turbid and a dark brown precipitate is formed.

So, in fully concentrated H2SO4, manganese(III) seems to exist in the form of a reddish/brown solid. Is this Mn2(SO4)3? On moderate dilution of the acid, a deep purple solution is obtained. This is not plain Mn(3+) in solution, but must be some sulfato-complex of Mn(3+), a possible option could be [Mn(SO4)2](-). On even stronger dilution, the color shifts to red/brown and then aqueous Mn(3+) is obtained (as shown in my testtube picture in a previous post in this thread). When dilution is even stronger, the Mn(3+) hydrolyses to Mn2O3.

I also added some of the bright purple NH4[MnP2O7] to conc. HCl (30%). When this is done, then the liquid becomes VERY dark, almost black. But in thin layers, such as when the liquid sticks to the glass, it appears to be olive-green. On heating, the liquid becomes lighter and chlorine is released. So, the very dark brown/green compound must be a manganese(III) complex with chloride. or simply MnCl3.

Up to now, however, I have not been able to obtain the green solid material. I get an understanding of all of this chemistry, but the green material still is lacking from my experimental results. Right now, some solid MnO2 (reagent grade) is brewing in H2SO4 of medium concentration, we'll see tomorrow what this will do. I do not have such light brown manganese oxide, so that part of the experiments cannot be reproduced by me

[Edited on 1-11-08 by woelen]

These colour changes on dilution are puzzling. I'll try and reproduce that with my own purple solution. My green compound seems more in line with Xenoid's Mellor quote of a green Mn2(SO4)3, it could explain why the stuff just dissolved and not hydrolysed. But it did that only once (out of three times).

You think the purple might be Mn³⁺ + 3 Mn(SO₄₂^-, huh?

Any idea how the Mn (IV) ends being reduced to Mn(II) in H2SO4 w/o the absence of an oxidiser?

Acc. 1911 Encyclopedia:

Obviously O [-II] ---> 1/2 O [0] + 2 e^-

could provide the missing electrons. I do not recall seeing any bubbles when heating with 50 % H2SO4, though...


And strange how you don't seem to get the green compound...

If I can repeat the green-to-purple trick with a known quantity of MnO₂, known volume of H2SO4 and a controlled volume of diluent water and everything dissolves I could have a good idea of the concentrations and the corresponding colours... And if I could get the unknown oxide to dissolve completely in the same way, it would open up a simply, titrimetric method for determining the Mn content of the oxide...

And adding extra H2SO4 to my reddish-purple solution could also push towards Mn(SO₄₂^-.

[Edited on 1-11-2008 by blogfast25]

If only I knew of such a specific reaction!

But guys, we're starting to sing from the same hymn sheet!

Having pondered my observations some more, I too am now convinced that the green compound is Mn (+III) bearing and probably simply manganic sulphate.

One thing makes me almost 100 % certain of this:

If you go back to my post of 7.20, second picture, both me and Woelen misinterpreted that. I thought MnO2 was being formed there, which then later dissolved in the extra added HCl, giving MnCl2 and Cl2. Woelen thought the heat I observed was simply due to dilution (hydration heat) of the 50 w% H2SO4. With hindsight, we were both wrong .

What happened in all likelihood is that the solid Mn (III) species reacted with the strong HCl: Mn (III) + 3 HCl + water ---> MnCl₃ + 3 H₃O⁺.

The dark liquid that I (rather prejudicially) mistook for a slurry of MnO2 was in fact quite concentrated MnCl₃ and its formation caused the heat to be generated.

I know this almost for certain because of the very unusually fast rate this liquid cleared up on heating: the decomposition of MnCl₃ is much faster than the dissolution of freshly precipitated MnO2 and subsequent decomposition of the formed MnCl₄. All the MnCl₃-bearing solutions I've heated up show that characteristic: they decompose remarkably fast to MnCl2 and Cl2.

Ironically, what started out as a mystery regarding the reddish Mn oxide, <i>may just</i> (note cautious emphasis) have yielded the possibility of titrimetrically assaying such oxides. If I can find the right conditions (acidity and sulphate concentration) in which the green Mn (III) compound dissolves without hydrolysis to a dilut<i>ish</i> Mn³⁺ solution, then the road is open to adding a standardised excess of Fe²⁺:

Mn³⁺ + Fe²⁺ ---> Mn²⁺ + Fe³⁺

followed by back titration of the excess Fe²⁺ with permanganate and calculation of the Mn content from sample weight, added ferrous mmol and permanganate volume. The method would of course be sensitive to the presence of chlorides.

This would really only be useful if the 'mystery oxide' behaves much in the same way but preliminary results seem to indicate that. Alternatively, if it doesn't, then the mystery oxide would have to be converted first to MnO2 via HCl and re-oxidation to Mn (IV).

Thanks for your cooperation so far!

+++++++++

And it may not even be too hard to find the window of conditions in which Mn³⁺ can operate relatively stably. One more very small experiment was carried out. Another 3 g of homemade MnO₂ was boiled up with 50 W% sulphuric acid for about 15 min. The green compound appeared. It was allowed to cool down to near-RT. And this time I saw bubbles (quite small) long before the mixture came to the boil..

A small amount of approx. 1 M H2SO4 was then poured into a conical flask and to it about the same amount of green slurry (carefully stirred for homogeneity) was then added to it. The green compound dissolved effortlessly, forming a dark, beautiful wine-red solution, here seen in a test tube This must be quite concentrated Mn2(SO4)3. The photo doesn't do it justice:




My working hypothesis for now is that any reasonable amount of the manganic sulphate is probably soluble in 1 M H2SO4, without hydrolysis, reduction or oxidation.

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Here's a few dilutions using the concentrated Mn (III) solution above:



Left: 1 ml diluted about 10 times with 1 M H2SO4. Light amber/reddish yellow.

Middle: 2 ml, diluted with approx. equal amount of 1 M H2SO4. An attractive wine-red.

Right, an original MnO2 boiled up with 50 w% H2SO4, with green Mn (III) compound and violet supernatant liquid, presumably containing an Mn (III) sulphato complex, as suggested by Woelen.

No hydrolysis noted so far.

(((((((((((((

Adding deeply cooled 50 w% H2SO4 to a small amount of the liquid from the middle tube causes the colour of the liquid to shift towards the violet obtained directly from reacting solid MnO2 and 50 w% H2SO4:



Left: the middle tube from above. Wine-red.

Right: a small amount of this solution with ice cooled 50 w% H2SO4 added to it. Shifts to more purpl<i>ish</i> - violet, definitely much more like the typical purple solutions obtained by attacking MnO2 with strong H2SO4.

The mangani sulphato complex can thus also be obtained by pushing:

Mn2(SO4)3 + SO₄^2- <---> mangani sulphato complex

to the right by increasing the SO₄^2- concentration, not only by attacking solid MnO2 with strong H2SO4.


[Edited on 2-11-2008 by blogfast25]