Manganates... not PERmanganates (Mn V and VI Redox Chemistry)

There don't seem to be that many good pictures of manganates out there on the internet, so here's a picture of some potassium manganate solution that I made:



And here is a small amount of barium manganate suspension:

That's something that I haven't seen any pictures of before. If it doesn't turn into manganese dioxide this time after I dry it, I will make more. I currently have about 220mL of dark blue-green potassium manganate solution prepared.

Edit: The second picture's color is a little bit off because when I added the potassium manganate solution to the barium chloride solution, it became too dilute for the manganate ion to exist which caused some permanganate to form, giving the solution a purplish color. The precipitate itself is clearly dark blue now that it has settled. Next I will try using a more concentrated barium chloride solution to avoid lowering the pH too much.

Edit Again: Well, it appears that using a more concentrated solution of barium chloride still causes the solution to turn purple, just a darker shade. The blue precipitate still forms. This purple compound confuses me though, because it isn't the right shade to be potassium permanganate. It's more royal purple than dark magenta. Could it be barium permanganate?
Also, just to see what would happen, I tried using calcium chloride instead of barium chloride in another test tube, and this led to a purple precipitate of what I presume is calcium manganate. I have not seen any information about this compound anywhere.


[Edited on 9-1-2014 by zts16]

Quote: Originally posted by woelen

blue MnO4(3-). The latter is hard to obtain in a pure state, it very easily disproportionates to MnO4(4-) and MnO4(2-). The ion MnO4(4-) is brown and quickly falls apart to hydrous MnO2 and hydroxide ions when in contact with water.

Quote: Originally posted by woelen

There also is a blue ion, MnO4(3-) and this ion is unstable, except with barium ions. The solid Ba3(MnO4)2 is quite stable and insoluble.

Alright, here are the results of adding drops of potassium manganate solution to solutions of various salts:

From left to right:
KCl (as a control to ensure that chloride ions don't interfere with the results, it appears they don't as no change took place)
BaCl₂ (dark blue precipitate, lavender solution)
CaCl₂ (light purple precipitate, but color may come from the solution, magenta solution)
CuSO₄ (bluish-gray precipitate, hot pink solution)
NiCl₂ (black precipitate, clear solution)
CoSO₄ (dark grey precipitate, looks darker in photo, clear solution)

The last two appeared to have reduced the manganate to Mn(IV). I'm not sure what the precipitates are in the middle two tubes. I'd like to think that they are manganates. The color of the solution above the barium manganate still confuses me also, because it doesn't look like the color of Mn(VII) to me, unless it's a trick of the light.

Edit: I vacuum filtered the barium manganate, and it does appear that the lavender color of the solution was a trick of the light. Once the solid and liquid were separated, the solution was clearly the color of Mn(VII), and the precipitate was the dark navy blue color that it is supposed to be. I really wish that it was possible to make barium manganate more efficiently, because with the current method that I am using an annoying amount of permanganate is also produced that makes the process quite impractical.

[Edited on 9-1-2014 by zts16]

Quote: Originally posted by chornedsnorkack

So, if you produce Mn(IV) in concentrated alkaline solution, do these solutions settle to colourless and clear as MnO2 precipitates out? Say you apply something reducing to K2MnO4/KOH solution... pure KOH solution is something like 1210 g KOH (22 mol) in 1 kg water at 25 Celsius, and pure NaOH at 25 Celsius saturates at something like 1110 g NaOH (28 mol) in 1000 g water. Does your dark green solution go blue, directly green to brown, or directly green to pale green to clear and colourless with solid MnO2 in precipitate?

Quote: Originally posted by woelen

The Cl(-) ion acts as reductor and in alkaline conditions you get hydrous MnO2 or even Mn2O3 (the latter is much lighter brown than MnO2).

Crystallization of Potassium Manganate

Quote: Originally posted by zts16

I know it must be possible, as it is supposedly sold by the large chemical companies as a green crystalline solid.

Quote: Originally posted by gdflp

http://www.alfa.com/en/catalog/39506 Alfa Aesar sells potassium manganate, but it must not be that simple to make crystals of because it's pricey, $60/10g or $181/50g

Manganese redox chemistry in strongly alkaline media (hypomanganates, manganates)

Hi everyone .

Few days after I attended in discussion about hypomanganates. So I decided prepare hypomanganate again and tested its stability and some redox chemistry.

So I prepared these three solution: 41% NaOH, 25% Na2S2O3.5H2O in 20% NaOH and approximately 0,0125M KMnO4.

All three solution were cooled to cca 2°C in fridge. After that I added few drops of sodium thiosulfate and six drops of potassium permanganate in to 41% NaOH solution. After that dark violet solution immediately turned colour to dark green by presence of manganate (picture one). I let it stand for hour in the fridge. After cca 30-40 minutes solution turned colour in to turquoise by presence of hypomanganate (picture two), but I let it stand for another 20 minutes just for sure that reaction was complete. Solution is now stand in the fridge in plastic beaker with lid and it still have beautiful turquoise colour (picture three).

I also made some solution of manganate for some redox reaction.

I found in one document effects of manganates and hypomanganates in strongly alkaline media on organic compounds, so I tested some of them. According to the document - ethanol should reduce manganate in to MnO2 and acetone in to MnO2 through unstable hypomanganate. Ethanol and aceton should reduce hypomanganate in to MnO2.

I tried reactions betwen strongly alkaline solution of manganate and ethanol/acetone. The reaction with ethanol took some time (few hours). Ethanol slowly reduced managanate in to MnO2. Acetone after few minutes reduced manganate in to hypomanganate. But reduction in to MnO2 wasn't occurre. The course of the reaction show pictures 4-8. In the pictures 4-6 is in left test tube manganate/acetone and in right test tube manganate/ethanol. In the picutres 7-8 is test tubes reverse (I transfered it in to plastic test tubes because i crushed one glass test tube).

In the pictures 9-11 is the same reaction but in hot water bath. Reaction was much quicker, but result was the same.

I tried reduction of hypomanganate by ethanol/acetone, but nothing happened. After few hours in the test tube with ethanol/hypomanganate disproportionation in to manganate and manganese dioxide has occurred. I tried reduction of hypomanganate by ethanol in hot water bath and it was slowly reduced in to MnO2 (pictures 12-15). Reaction lasted few minutes.

Last reaction I tried was reducing of hypomanganate by 3% H2O2 in to MnO2. This reaction was very quick and after one minute was complete (picutres 16-17).

Here are photos of hypomanganate solution for five days storage.

I haven't any barium salt so I can't try this . In my country isn't very e-shops which sell chemicals to people without business (for example I don't know where I buy manganese sulfate or chrome alum which are common chemicals).

Quote: Originally posted by Texium (zts16)

I was thinking that allowing a basic solution of potassium manganate to evaporate would leave me with nice green crystals sitting in an ultra-concentrated NaOH solution that doesn't evaporate because of its deliquescence, but instead it looks like it oxidized to permanganate despite being in strongly basic solution.

I made sodium hypomanganate for the first time a few days ago, used a totally different method based on the one used when it was first discovered. I'd tried various aqueous routes but they never seemed to work, although I haven't yet tried the one from Bedlasky shown here.

I melted about 10 grams of sodium nitrite by heating in a beaker using a Bunsen and then added about 250mg of manganese dioxide.This was then stirred with constant heating for about 10 minutes and much of the manganese dioxide dissolved to give a dark grey-brown solution, but a lot remained suspended. Next I added about 2.5g of sodium hydroxide pellets and heated the mixture gently to keep it molten with occasional stirring until a dark blue-green solution formed. This took about 15 minutes. A little of the manganese dioxide still remained suspended, but not much. The mixture was left to cool and formed a blue-grey solid that got steadily more blue-green and then bright blue once it was completely cool. This took about an hour as it kept producing heat as it cooled. Photos are attached.

If you do this then be careful when you add the sodium hydroxide pellets as it spits like mad (go slowly). The mixture has also eaten the end of a glass rod and has corroded the surface of a pyrex beaker. Sometimes the beaker will crack during cooling, so plan for this. I had great fun dripping a little of the mixture onto paper towel and watching it ignite on contact with a single drop, so again take care.

Sometimes the product is somewhat greener than in the photos, generally when I heat it more aggressively or use too little sodium hydroxide, so be prepared to experiment a bit as the recipe above could be refined further for better results. The product decomposes to a dark grey-brown mess after a day or two.

Quote: Originally posted by Bedlasky

I haven't any barium salt so I can't try this . In my country isn't very e-shops which sell chemicals to people without business (for example I don't know where I buy manganese sulfate or chrome alum which are common chemicals).

Quote: Originally posted by Fery

this shops sells to chem fans too, no need to have business ID Ba(NO3)2 https://www.kovyachemie.cz/dusicnan-barnaty/ Mn https://www.kovyachemie.cz/kovy-spony-a-kousky/mangan-elektr... they sell other chemicals too like BaCO3 and even BaO2 I'm buying from this shop for 15 years already, very good prices, quick delivery - usually the second day after order, at most 2 days (if you order during night :-D)

Thank you for this shop!

Hi.

Here are photos from experiment.



Reduction of manganate. From left: glucose, fructose, sucrose, tartaric acid, citric acid, ascorbic acid.

Test tube with fructose reduction is orange due to decomposition products of fructose in strongly alkaline solution.



Reduction of manganate. From left: nothing (just manganate solution for comparison), dithionite, oxalic acid, sorbitol



Reduction of hypomanganate. From left: glucose, fructose, sucrose, tartaric acid, citric acid, ascorbic acid.

Test tube with fructose reduction is red due to decomposition products of fructose in strongly alkaline solution.



Reduction of hypomanganate. From left: nothing (just hypomanganate solution for comparison), dithionite, oxalic acid, sorbitol

[Edited on 11-2-2020 by Bedlasky]

After discovering this thread a few months ago and successfully preparing my first hypomanganate sample with the thiosulfate procedure by Bedlasky, I went on to do some additional hypomanganate experimentation, using different reducing agents and reaction conditions. I wanted to see the colour a very concentrated hypomanganate solution would have. Trying to maximize the reduction efficiency of manganate, hoping to minimize any contamination with residual manganate and bring out the bluest blue of hypomanganate.

My starting material was a rather concentrated solution of manganate (roughly 0.1M), made by adding KMnO4 to a boiling 10M NaOH solution:



A small amount of this solution was diluted in 2M NaOH to have a sample for comparison later.

When the solution completely cooled down, it was put in the freezer, to get it very cold. After about 2 hours of sitting in the freezer, it was taken out. It had a syrupy/pasty consistency, as some of the hydroxide precipitated out. But this wasn't a bad thing, as I wanted to limit any decomposition as much as possible and I would assume a solid NaOH lattice would help to further stabilize hypomanganate at such a high concentration.
A refrigerated solution of 0.25M Na2SO3 (Na sulfite) in 5M NaOH was added in roughly a 1.5 molar excess of Na2SO3 in regards to manganate. This was thoroughly stirred to ensure proper mixing of Na2SO3 with the manganate paste. It was placed back in the freezer. After 1 hour it was taken out, with the following result:



Comparing the colour of the thin layer of liquid on glass with the starting solution of manganate, it is very clearly more blue.

An amount of this new material (approximately the same as with manganate from before) was diluted in refrigerated 10 M NaOH. The resulting diluted solutions of manganate and hypomanganate were compared:



To be frank, my camera doesn't pick up intense greens to well, so the above image doesn't accurately reproduce the colours. The manganate was much deeper and saturated in reality, and the hypomanganate was also slightly deeper in colour. I transfered both solutions into test tubes, to examine the colour when it's made less intense, which resulted in a far more accurate representation of both colours. I judged both the photo and the actual view of the test tubes, and this time the difference was much less significant, so you can take my word that this is photo is the closest in terms of colour reproduction out of all the ones I have taken.



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I also tried what has been suggested by others - adding an alkalized solution of barium salt to the hypomanganate. I dissolved Ba(NO3)2 in 5 M NaOH. I eyeballed it without paying attention to molarity. I then diluted some of the thick hypomanganate "syrup" in 10 M NaOH. I added some of the Ba solution to the diluted hypomanganate and left it in the refrigerator for 4 hours to see if I get a precipitate. Here was the result:



Important was making the Ba nitrate solution fairly alkaline itself. I tried with making a saturated solution in pure water and adding a really small amount of this saturated solution to the hypomangate, but the hydroxide concentration becomes insufficient and the Ba hypomanganate precipitate disproportionates, rapidly turning a murky brownish green, unlike Ba manganate which becomes sufficiently stable once it is precipitated.
Also, it should be considered that other solids are in the precipitate, such as BaSO3, BaSO4, BaCO3, etc. So perhaps this isn't the colour of "pure" Ba3(MnO4)2, but since it is a lot lighter in colour then what you get when precipitating manganate with Ba, I'll claim this as a success in obtaining at least some barium hypomanganate.

I intent to try precipitating a larger amount and perhaps collecting it to dry out, but seeing how easily it disproportionated once the hydroxide concentration dropped, I doubt that this is possible at all.

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Here is one more photo of just the hypomanganate, from an earlier run where thiosulfate was used as the reducing agent, according to Bedlasky's method, except again I've made a concentrated hydroxide "syrup" to create an enviroment that minimizes any chance of hypomanganate decomposition. The material, diluted in 10 M NaOH, exactly matches the "sky-blue" colour description found in some literature. Very pretty colour.



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I mainly did this small project out of abundance of spare time, but also for disambiguation purposes, so if someone wants to also prepare hypomanganate in the future, they will know what to look for to see if the method worked. Manganate is often described as a bluish-green, so it might be tricky to tell the difference between manganate and hypomanagate (especially in very small concentrations) if all you know is that hypomanagate is supposed to be blue in colour. I hope the above photos give a better idea of the difference between the respective colours of MnO4^2- and MnO4^3-. And also, there's not enough photos that show the very rare hypomanganate, so I'm glad I could provide additional ones here.

Many people have tried to produce hypomanganate without success, but it's perfectly possible to do it - the key is keeping low (preferrably ice cold) temperatures, as hypomanganate very easily disproportionates with increased temperature. And copious amounts of NaOH or KOH at every stage of the experiment, which should be taken into proper safety considerations. Also it is important to use a slight excess of reducing agent, but not too much, because hypomanganate will be reduced further if too much reducing agent (such as sulfite) is used. Thiosulfate seems to be mild enough to not overreduce it, however.

I would also like to try the method of dissolving MnO2 in a molten mixture of NaNO2 and NaOH and see if I can extract the hypomanganate by dissolving it in cold 10 M NaOH and see if the resulting colour of solution and Ba precipitate is any different. I just need to find a suitable vessel to contain the molten NaNO2/NaOH mixture.

[Edited on 10-12-2020 by EthidiumBromide]

Quote: Originally posted by EthidiumBromide

Manganate is often described as a bluish-green, so it might be tricky to tell the difference between manganate and hypomanagate (especially in very small concentrations) if all you know is that hypomanagate is supposed to be blue in colour.

Quote: Originally posted by EthidiumBromide

Also it is important to use a slight excess of reducing agent, but not too much, because hypomanganate will be reduced further if too much reducing agent (such as sulfite) is used. Thiosulfate seems to be mild enough to not overreduce it, however.

Quote: Originally posted by EthidiumBromide

I also tried what has been suggested by others - adding an alkalized solution of barium salt to the hypomanganate. I dissolved Ba(NO3)2 in 5 M NaOH. I eyeballed it without paying attention to molarity. I then diluted some of the thick hypomanganate "syrup" in 10 M NaOH. I added some of the Ba solution to the diluted hypomanganate and left it in the refrigerator for 4 hours to see if I get a precipitate. Here was the result: Important was making the Ba nitrate solution fairly alkaline itself. I tried with making a saturated solution in pure water and adding a really small amount of this saturated solution to the hypomangate, but the hydroxide concentration becomes insufficient and the Ba hypomanganate precipitate disproportionates, rapidly turning a murky brownish green, unlike Ba manganate which becomes sufficiently stable once it is precipitated. Also, it should be considered that other solids are in the precipitate, such as BaSO3, BaSO4, BaCO3, etc. So perhaps this isn't the colour of "pure" Ba3(MnO4)2, but since it is a lot lighter in colour then what you get when precipitating manganate with Ba, I'll claim this as a success in obtaining at least some barium hypomanganate. I intent to try precipitating a larger amount and perhaps collecting it to dry out, but seeing how easily it disproportionated once the hydroxide concentration dropped, I doubt that this is possible at all.

I tried to make barium manganate more or less according to the procedure in the attached document, following the below extract but with reduced quantities. But I had to use more water for the BaCl2 to get it to dissolve; not sure if the below is a typo or whether it was supposed to remain in suspension:

Attachment: Convenient Routes to Ba Permanganate.pdf (80kB)
This file has been downloaded 337 times

"118.5 g of potassium permanganate was dissolved in 1.5 L of water,
then the solution was mixed with 183 g of BaCl2.2H2O (in 45 ml
H2O), 30 g of NaOH (in 130 ml H2O) and 15 g of KI (in 60 ml
H2O). The mixture was boiled for 10 min, decanted, then the precipitate
obtained was washed 10 times with distilled water and
dried at 110 °C. Yield is 161 g. Analysis of the product was performed
by the method of Schlesinger [9]. Composition of the dried
product was 66.3 % of BaMnO4, 13.9 % of MnO2 and 19.8 % of
water."

- 23.9g KMnO4 in 300g water in beaker, heat and stir
- 36.3g BaCl.2H2O in 100g water in separate beaker, heat to near boiling to fully dissolve
- 6.2g NaOH in 25g water in separate beaker
- 3g KI in 12g water in separate beaker
- KMnO4 solution on stirring hotplate. Add BaCl2 solution.
- Add NaOH solution.
- Add KI solution.
- Dark solution. Heat to boil and boil (took some time) and boil for 10minutes.
Note: The attached mentioned 10 minutes boiling, but for larger quantities. I kept the same boiling time, but did wonder whether I should have shortened the time.
- Deep purple coloured solution, dark ppt. Diverge from procedure and try to vacuum filter: big mess and my vacuum pump also decided to stop working at that point!
- Scratch all of the dark ppt (which stained the beaker dark green) into a 2 litre beaker.
- Add 1.5 litre water. Stir 10 minutes. Leave to settle for one hour.
- Very dark ppt with purple supernatant solution - similar to a very dilute KMnO4 solution color.
- Repeat the above 2 steps a total of 6 times. Thus, in total "washed" with 9 litre of water.
- Not much change in the color of the ppt or the supernatant solution, after all the washing. Not sure if some of the barium manganate decomposes to permanganate with excess water thus the constant purple color of the supernatant liquid?
- On steam bath. Dry some 6 hours till weight of the product remained constant at 26.2g Very dark black-blue product, see photo.





The attached reference says it cannot be avoided to have around 14% of MnO2 in the final product, following this method. Whether I have more than that I do not know. Wikipedia says the color or barium manganate can be "light blue to dark blue and black" which means the color alone does not necessarily say a lot.

I did some very simple tests with acids, mainly in the process of cleaning everything:
- With concentrated or slightly dilute HCl it gives a brown ppt, when heated with either concentrated or diluted HCl the brown ppt dissolves and the solution eventually becomes clear. Plus a faint chlorine smell.
- With dilute H2SO4 there are some bubbles and a brown ppt, at room temp.
- With concentrated HNO3 there was an interesting observation: a while 'smoke' comes off which has no smell.

Very nice result Lion850! I always think that barium manganate is green .

If you want to do some analysis of your product, try chelatomeric determination of Ba (ammoniacal buffer pH = 10, eriochrome black T as indicator). Firstly dissolve barium manganate in sulfuric acid, filter, washed well filter paper, neutralize, add buffer and indicator and titrate with 0,01M EDTA.

I did today some spectroscopic study of MnO₄^n- anions in UV-VIS region using cca 0,8mM solution of permanganate, manganate and hypomanganate.



All anions have common peak with maximum at 285 nm.

Manganate and hypomanganate have similar peaks in the range 540-800 nm. Manganate peak is shifted more to the left with 608 nm absorbition maximum. Hypomanganate have maximum at 675 nm.

Permanganate have typicall "finger tips" peak in the range 450-600 nm.

Manganate have typicall peak in the range 410-480 nm with maximum at 434 nm.

Hypomanganate have typicall sharp and tall peak in the range 310-320 nm with maximum at 316 nm.

There are also some small peaks for manganate with maximums at 260 and 340 nm and for hypomanganate with maximum at 321 nm.

Nice Absorbční Křivka
Very interesting to see how part of the spectrum changes, whereas the peak around 285 nm stays virtually the same.
How did you record these? Do you have a spectrometer yourself?

I had another go at potassium manganate. By boiling potassium permanganate in potassium hydroxide. I got the green product as a ppt / suspension, as can be seen in the below photo.


But I got stuck trying to filter it off. Plan was to filter it and then press dry the obtained green crystals between filter papers, and bottle it still slightly wet with KOH. Problem, which I should have foreseen!, is that it reacts with the filter paper and decomposes. So I stored the wet product, best I can do until I can filter it.


Any suggestion for filtering this product will be appreciated.