Sciencemadness Discussion Board

Peroxymanganates

Taoiseach - 9-12-2008 at 08:01

Ammonium peroxymanganate is mentioned in the Hazardous Chemicals Handbook.

Also, peroxymanganates with 2 or 4 peroxo ligands are mentioned here:

http://www3.interscience.wiley.com/journal/109798451/abstrac...

It is said that the sodium compound can be crystalized in a pure form.
Unfortunately I cannot read this article :(

These compounds sound very interesting, as permanganates alone are very powerful oxidizers. Replacing the O(2-) with O(1-) should make them even more energetic.

woelen - 9-12-2008 at 12:16

Too bad that this article cannot be read without an account. I did not know about the existence of peroxo-complexes of manganese. These compounds must be extremely unstable. I know from experience that permanganate reacts with hydrogen peroxide at any pH. At low pH, it is converted to Mn(2+), at medium to high pH, it is converted to MnO2.

Formatik - 9-12-2008 at 13:42

There is also this one here:

http://www3.interscience.wiley.com/journal/109798436/abstrac...

Basically, reaction of alkali manganates (VI) or Mn(II)-salt with H2O2 in a cooled alkaline solution forms the easily soluble alkali peroxymanganates. They describe the preparation and properties of a "very unstable", crystalline potassium peroxymangante. They determine a formula as: K2H2[MnO(O2)3].

[Edited on 9-12-2008 by Formatik]

JohnWW - 9-12-2008 at 14:23

Quote:
Originally posted by woelen
Too bad that this article cannot be read without an account. I did not know about the existence of peroxo-complexes of manganese. These compounds must be extremely unstable. I know from experience that permanganate reacts with hydrogen peroxide at any pH. At low pH, it is converted to Mn(2+), at medium to high pH, it is converted to MnO2.

Someone please download and post these articles! The reaction of Mn(VII) as MnO4- with H2O2, resulting in double decomposition with evolution of O2 and Mn(II) or (IV), is likely to be highly temperature-dependent, in view of its exothermicity, so peroxymanganates, whether Mn(VII), (VI), or (V), may be possible at low temperatures, particularly in supercooled solutions.

sparkgap - 9-12-2008 at 19:35

Now all you need is somebody with good chemical German... ;)

sparky (~_~)

woelen - 9-12-2008 at 23:51

Thanks very much. I have no problem reading this. This evening I'll read this more thoroughly and if wanted, I can summarize it.

Taoiseach - 10-12-2008 at 00:21

According to the paper (if I understand correctly), the well-known catalytic decomposition of H2O2 with MnO2 goes through an intermediate peroxy compound which is instantly decomposed by water at room temperature. MnO2 is said to dissolve(!) in H2O2 at -20°C forming a peroxy complex.

Formatik - 10-12-2008 at 08:52

Quote:
Originally posted by sparkgap
Now all you need is somebody with good chemical German... ;)

sparky (~_~)


They first showed from an earlier cited procedure addition of H2O2 onto K2MnO4 in a cooled, strongly alkaline solution formed the salt of a peroxymanganic(IV) acid. Treating such a solution with cooled 20% pure KOH solution and finally with cooled 3% KOH containing methanol, could pure, very decomposable dipotassium-oxotriperoxo-manganate(IV) K2H2[MnO(O2)3] . aq be obtained.

Here is a preparation procedure they give for a potassium tetraperoxomanganate(IV): 50 mL of pure perhydrol (25 mL H2O, 25 mL H2O2, so it looks like 50% by vol) has under strong cooling (-10 deg.) gradually 20-50 g KOH added and solubilized. Now another solution cooled at -16° which consists of 5 g moist K2MnO4, 20 mL H2O + 0.5 g KOH is slowly dripped in.

Using 50 g KOH: there was an immediate precipitation from the deep dark-red brown solution, with 20 g first after 20 minutes. This is filtered under cooling. And finally washed 4-6 times with excess of acetone at -30°, then it was stored in acetone at -70°.

For the preparation of sodium peroxymanganate(IV): this is easier to do because it is more difficultly soluble than the potassium salt. So it is easily obtained from relatively lower concentrations of NaOH and H2O2, so that at an about same time occurring crystallization out of the cooled H2O2 containing NaOH precipitating Na2O2.8 H2O can be avoided.

They used cooled H2O2-containing NaOH solution, then a filtered solution containing 3 g freshly prepard, moist Na2MnO4 in 15 mL H2O + 0.5 g NaOH, which was added dropwise. The amounts and temperature are in table 5. From the deep dark-red solution, very quickly a rich amount of an almost black crystallization precipitates. Which under the microscope consistently are small, multiple aggregated, raddled, plates that have a green shimmer. Then the precipitated compound was filtered under cooling, washed richly with acetone at -3°, then washed with acetone at -15 to -20°. Then stored under acetone at -70° just like the potassium compound.

I didn't read the paper in full, but they are also covering mixed crystallizations also containing K3H[MnO(O2)3].aq.
Quote:
Originally posted by Taoiseach
According to the paper (if I understand correctly), the well-known catalytic decomposition of H2O2 with MnO2 goes through an intermediate peroxy compound which is instantly decomposed by water at room temperature. MnO2 is said to dissolve(!) in H2O2 at -20°C forming a peroxy complex.


They are saying according to Bredig and Marck, freshly precipitated MnO2 solubilizes in a very small amount in 30% H2O2 at -20 deg. colorless, forming a manganese peroxyhydrate, the solution decomposes at RT forming MnO2 and lively giving off O2.

[Edited on 10-12-2008 by Formatik]

kmno4 - 14-12-2008 at 12:25

With all respect for authors of these old publications but peroxymanganates (or peroxomanganates) seem to exist only in these two german publications.
ACS, RSC, Science Direct and Wiley gives only two articles when I ask about "peroxy(o)manganate" - just these ones mentioned in this thread.
It seems that all those experiments and crystal structures of prepared salts should be reivestigated or determined.
Peroxymanganates can be only wishful thinking......

Per - 14-12-2008 at 14:13

The German papers says that the Peroxymanganates could be prepared under cooling to -15°C and strong alkaline conditions, the crystals of the manganese peroxides must be washed with - 25°C cold acetone, in complete everything else than easy to make because of the ease of decomposition of the stuff.

Also they say that the isolated Peroxomanganates decomposed very rapidly at room temperature giving oxygen, so I would say that these are interesting compounds unfortunately with no practical use for the home scientist and maybe for everybody else.

UnintentionalChaos - 14-12-2008 at 16:08

Perhaps these compounds are similar to, but even less stable than their period table neighbors, the Tetraperoxochromates. Excess of base and ice cold temperatures sure sound like similar conditions. And if you did it at room temp, you'd get...oxygen and Cr(III).

It might be worth a shot. At least the products aren't carcinogenic in this case, and they may turn out to be more stable than the literature states, as woelen showed us with the potassium tetraperoxochromate.

[Edited on 12-14-08 by UnintentionalChaos]

woelen - 15-12-2008 at 00:01

I tried adding H2O2 3% to an icecold solution of K2MnO4 in icecold conc. NaOH. KMnO4 is not possible at these high concentrations of NaOH, it decomposes, giving oxygen and the green manganate.

As soon as the H2O2 reaches the liquid, it fizzles and the green color changes to a turbid brown. The solutions I used have temperatures of appr. -10 C (maybe a little colder, I cooled them in our freezer).

Nothing interesting happened though. So, if these peroxymanganates exist, then they only exist at really low temperatures, which I cannot easily achieve. My book "Chemistry of the elements", which is from 1998 or so, does not mention the existence of peroxychromates.

Someone with access to CO2-snow should try this experiment and see if the red solutions can be made. I only had brown precipitates of MnO2.