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[*] posted on 10-10-2008 at 10:06


Success: Na2MnO4
==> from NaClO3 (not even clean,since difficult to crystallize without NaCl)
==> and NaOH
==> at 400 [Cels], 10-20 minutes

10 g NaOH + 10 g MnO2 (100% MnO2, pottery-grade) + 10 g of the not very clean NaClO3 gave a _very_ green solution, as green as known from dissolving KMnO4-thermal-decomp-educts; definately has some concentration !

What didn't work: KNO3 and NaNO3 just dont give the slightest green (KNO3 was tried without any hydroxide).
Also: NaClO3 + Soda (instead of NaOH) failed, at the 400 [Cels]; didn't try at higher temp.

So:One doesn't even have to crystallize the NaClO3 too extensively, just a raw crop from the electrolysis-cell does it !

[Edited on 10-10-2008 by chief]
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[*] posted on 10-10-2008 at 10:43


How long was the mix heated and did you add the NaClO3 seperatly or just chuck all the reagents in and begin heating?
Nice work btw! Have you tried acidification yet?
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[*] posted on 10-10-2008 at 11:44


It was only in the furnace for 20 minutes, incl. heating-up-time;
the furnace was controlled to have 400 [Cels]; maybe the mixture needed up to 10 mins to reach this, and maybe part of it (bottom) got hotter by maybe 20 or 30 degrees, so local max-temp would have been 430 [Cels].

It was prepared in a old coffee-mill (10000 RPM) (which will never see any food again), everything was fine-powdered and well mixed; the MnO2 was from a pottery-supply, stated as "100 % MnO2"; maybe I'll try the battery-stuff too ... (thoroughly washed, filtrated, pre-glowed with some nitrate to remove the graphite first !)

[Edited on 10-10-2008 by chief]

Ah yes, the acidification: I preferred to ppt. out with Ba(NO3)2, to get the insoluble BaMnO4 ... to be continued ...

[Edited on 10-10-2008 by chief]

[Edited on 10-10-2008 by chief]
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[*] posted on 12-10-2008 at 07:12


... and the ppt, upon acidification, gives a violet-color; but it always is a slightly other violet (with Ba) than eg NaMnO4 OR KMnO4, somewhat reddish component when dilute.
Only: When acidification was with HNO3, at boiling-down the HNO3 gets too concentrated, and it all gets MnOx-ppt..
The other thing I tried was: NaMnO4 + MgSO4 (solution), as suggestedin a link from above (http://www.retrobibliothek.de/retrobib/seite.html?id=116000#...), gives also permanganate,but upon filtering the violet solution through the same filter as the green manganate-solution before, it turns into green again, perhaps some reaction with the MnO2 that was still in the filter (and that I left there to "better" the filtering-effect ...)

Also: Did the Na2MnO4 via chlorate, as above, several times now: It works. It doesn't harm, also, to set the furnace to 470 [Celsius]. Upon boiling down the green solution in the microwave the glass cracked quite cleanly around a ring where some already dry manganate was solidified on the wall (susceptibility of MnOx maybe high) .. maybe a way of cutting glass/bottles etc: Manganate around the bottle, somehow painted onto it ==> microwave for a few minutes ...

Now I think of how to mildly acidify it with CO2, don't want to abuse any acid for that purpose, since until now it's also an acid-free process to the manganate: Only MnO2 + NaOH + chlorate ...

Didn't find anything upon an easy-to-make CO2-generator at the forum here ...

[Edited on 12-10-2008 by chief]

[Edited on 12-10-2008 by chief]
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[*] posted on 27-10-2008 at 13:43


Quick question about potassium permanganate: If sulfur dioxide is dissolved in dilute potassium permanganate solution, the violet color disappears. Does this mean that the potassium permenganate present oxidized the sulfur dioxide into sulfur trioxide? Does the clear color mean sulfuric acid? Or did the sulfur trioxide combine with something somewhere. Like with the manganate, to permanganate. Not quite sure. HELP!?
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[*] posted on 27-10-2008 at 13:45


It means the SO2 is reducing the Permanganate... you would definitly not get SO3 just left in soloution!!!
SO3 would form H2SO4 as soon as its made practically...
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[*] posted on 27-10-2008 at 15:12


Quote:
Originally posted by Picric-A
It means the SO2 is reducing the Permanganate... you would definitly not get SO3 just left in soloution!!!
SO3 would form H2SO4 as soon as its made practically...

But the Mn is still there! So either MnSO4 OR MnSO3 will be formed ? Maybe it's at least a way to oxidize SO2 to SO3 in Salt; later decompose the MnSO4 thermically, re-use the MnOx with another batch of cheap chlorate, and there is the H2SO4-engine !! Could this work that way ?

[Edited on 27-10-2008 by chief]
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[*] posted on 27-10-2008 at 16:33


hmm, i said this in another post, i thought what could happen is the SO2 would be oxidised to SO3 but reduce the permanganate to manganate. the SO3 would form H2SO4 and this would oxidise the manganate back to permanaganate... where would it end? would the SO2 reduce the KMnO4 back to MnOx/MnSO4 and K2SO4?
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[*] posted on 28-10-2008 at 00:53


With Na2S2O5, a SO2-source (cheap available), the reaction goes on like this, as I stated earlier (previous page) and observed myself:
Quote:
Originally posted by chief
The dissolution of MnOx sometimes occures, and at other times it doesn't:
KMnO4-crystals completely dissolve, nothing stays undissolved, and the severaln stages of reduction can be observed:
When shaking in a flask:
==> near the KMnO4 it's violet
==> 1/2 cm distance its brown and looks like MnO2
==> a little further and it looks whitish
==> then its dissolved.


So the Mn is reduced from the +7 to the +2 stage, giving probably either the MnSO3 OR MnSO4. Which one ?

[Edited on 28-10-2008 by chief]

[Edited on 28-10-2008 by chief]
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[*] posted on 28-10-2008 at 00:58


Quote:
Originally posted by chief
So the Mn is reduced from the +7 to the -2


Umm?

Quote:
stage, giving probably either the Mg


Umm?

Quote:
SO3 OR MgSO4. Which one ?


Manganese sulfate, if you didn't use too much excess sulfite. If you did, there may be some sulfite present. But these salts are ONLY formed if you evaporate the solution and crystallize it. Ions in solution are not compounds, they are ions in solution.

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[*] posted on 28-10-2008 at 01:18


@12AX7 MnSO3/4 (!), not MgSO3/4, of course ! Thanks for the correction, I edited it !

The reduction from +7 to +2:
KMnO4 : Mn ==> is +7 (K ==> +1 ; O4 ==> -8 ; ==> Mn ==> +7)
MnSO4: ah, yes : +2, not -2, I'm gonna edit it again.

Maybe I was with the Sulfur, which goes from neg. to pos: (H2S,S,SO2,SO3)

[Edited on 28-10-2008 by chief]
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[*] posted on 13-3-2009 at 10:38


Question: The electrolysis from manganate to permanganate:
==> Can it be done with PbO2-Anodes instead of Nickel ?
==> What about graphite ?

Besides: I reported above the manganate-formation at 800 [Cels] with Na2CO3 instead of NaOH ; I then found, that this is due to the decomposition of the Na2CO3 ... to NaO or someting . I think it's like that because the 780-800 [Cels], which are also the decomp-temp for the Na2CO3, were necessary for anything to get green: At lower temps no amount of time would do it (up to several hours; maybe a grinder/mixer can do after 48 hours ..., but I wouldn't bet on it) ; maybe just a mistake in the literature ...

I tried it with clean NaOH, and it works _well_ , _deep_ green solution ...
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[*] posted on 6-7-2009 at 12:08


Quote: Originally posted by Formatik  
It needs to be dehydrated also to form the compound. I've added some KMnO4 powder to reddish fuming HNO3 with a density of about 1.52, no Mn2O7 formed, and all of the acid fumes turned white. The same is in the well-known chlorine generator reaction of KMnO4 with conc. HCl, though the permanganate can also oxidize the conc. acid to some ClO2, which is explosive.


I've also tried it with the strongly hygroscopic conc. HClO4, where adding powdered KMnO4 didn't form Mn2O7 initially, but just a red mixture that discolors on contact with tissue. Though on standing longer the mixture does form small, shiny oily green droplets which have the reaction of Mn2O7.

Though as far as oxidation of HCl by KMnO4 goes, this is said by at least several references in Brethericks to have given a sharp explosion when conc. HCl was added to solid KMnO4, the explanation was that there is a remote possibility of the permanganate oxidizing the Cl2 to "chlorine oxide". I've added 0.4g conc. HCl to 1g KMnO4 dust in a clean, dry test tube, after some gas evolution a lit long match was inserted two times but nothing happened. Addition of liquid Cl2 to (moist or dry) KMnO4 would surely prove the hypothesis wether Cl2 is actually oxidized.

[Edited on 6-7-2009 by Formatik]
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[*] posted on 1-9-2009 at 07:39


I should have looked over the forums more thoroughly; I didn't know there was a dedicated thread on permanganates. I've wasted a lot of time recently on various attempts at manganates and permanganates. Complicating matters was that KMnO4 or K2MnO4 were out of consideration because I wanted the substance as an aid in detection of potassium as KClO4 with the microscope. I've been trying to keep notes of my progress using Google Docs; here's the document pertaining to all the manganate/permanganate attempts:

http://docs.google.com/Doc?docid=0AWsze3MPrfu6ZGh0eDUzZ3dfMj...

Quick summary: I succeeded in preparing some impure BaMnO4, but failed at converting it to Ba(MnO4)2 with CO2. Fusions of MnO2 with NaOH and added NaClO3, followed by extraction of the melt with dilute NaOH, gave Na2MnO4 solutions but I mostly succeeded only in destroying them when trying to convert them to some kind of soluble, crystallizable salt. The Na2MnO4 solution from the last attempt was divided, one half acidified to yield permanganate, and both solutions promptly frozen until I can figure out how to process them further.

I did succeed in preparing one well-crystallized salt, mixed crystals of Na2SO4.10H2O and Na2MnO4.10H2O in undetermined proportion, unstable in open air but stable in a humidor. It's enough for my purposes--all I need is a small, unmeasured amount of a manganate or permanganate for treating drops on a microscope slide--but not much use for large-scale preparation.

A practical note that doesn't come across well in the notes (I need to improve my note-taking): I cracked a lot of cheap porcelain crucibles doing this. Thermal stress was no doubt sometimes to blame: my sources of heat are poor, comprising only an alcohol lamp (wretched), a couple of different propane torches (one the standard pencil torch, another with a fairly broad flame), and an electric range. But even using the electric range and gradual heating to the maximum value, the porcelain cracked, sometimes extensively. I know that alkaline melts attack porcelain but I'd underestimated how much it would weaken it.

I'll have to read up on this thread. I probably could have avoided some rookie mistakes had I known it existed.




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[*] posted on 1-9-2009 at 07:54


Quote: Originally posted by monoceros4  
(cut)
A practical note that doesn't come across well in the notes (I need to improve my note-taking): I cracked a lot of cheap porcelain crucibles doing this. Thermal stress was no doubt sometimes to blame: my sources of heat are poor, comprising only an alcohol lamp (wretched), a couple of different propane torches (one the standard pencil torch, another with a fairly broad flame), and an electric range. But even using the electric range and gradual heating to the maximum value, the porcelain cracked, sometimes extensively. I know that alkaline melts attack porcelain but I'd underestimated how much it would weaken it.(cut)
More resistant materials than porcelain (or borosilicate glass) should be used as crucibles for reactions involving fused or hot concentrated alkalis. The reaction mixture would become contaminated with silica and alumina. The first step upwards would be fused zirconia, or failing that, graphite or nickel; or better still, platinum if it could be afforded.
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[*] posted on 1-9-2009 at 08:11


Quote: Originally posted by JohnWW  
More resistant materials than porcelain (or borosilicate glass) should be used as crucibles for reactions involving fused or hot concentrated alkalis. The reaction mixture would become contaminated with silica and alumina. The first step upwards would be fused zirconia, or failing that, graphite or nickel; or better still, platinum if it could be afforded.

I suppose I could try to find a used platinum crucible but then I'd have to go without paying rent for at least a couple months. Vitreous carbon is somewhat cheaper. I have a small nickel crucible but was hoping to manipulate larger amounts of material. Would magnesia hold up? It's reasonably priced.

I didn't worry about silica or alumina contamination, partly because it's highly unlikely that any significant Si or Al would find its way into a manganate or permanganate crystallization, partly because nobody expects permanganates to be all that pure anyway.

If I mess with this again I hope to avoid methods involving fusion. Even when it works out it's so damn tedious.

[Edited on 1-9-2009 by monoceros4]




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[*] posted on 2-9-2009 at 22:17


I have made some progress in an unexpected direction and prepared BaMnO4 from MnO2 without needing to melt or electrolyze anything.

User Paddywhacker in another thread (that I shouldn't have created) suggested to me that a good procedure for making barium manganate, BaMnO4, would be useful because of the compound's utility as a mild oxidant. I was interested because BaMnO4 is a route to barium permanganate, thence to other permanganates. I tried a fusion method the first time and got poor results.

I think I've done better this time and with a very simple method requiring no special equipment. Finding references to "hydrothermal" methods of preparing potassium manganate, i.e. operating in aqueous solution at high temperature and pressure, I reasoned that BaMnO4 ought to be even easier to prepare that way because the insolubility of the compound would help the reaction along. To test this I heated 2.0 g of MnO2, 9.0 g of impure, homecooked Ba(OH)2.8H2O, and 3.0 g of KNO3 (I wasn't sure of the necessity of this--maybe air is enough?) moistened with a little water in a loosely closed container in an ordinary pressure cooker and lots of water at 15 psi overpressure. The result was a greyish-green, somewhat heterogeneous mass that, after exhaustive washing with cold water, drying, and grinding, yielded 6.1 g of olive-green powder. The weight is not far off theoretical but the substance is definitely not pure. While I saw only a few black particles, indicating that most of the MnO2 had reacted, there was more whitish matter that was probably BaCO3 from the impure Ba(OH)2. (This was later confirmed: a pinch of the solid fizzed briefly in weak acid.)

I'm in the process of trying to assay the product. My attempt to determine "active oxygen" by boiling a small sample with strong HCl and collecting the Cl2 in dilute NaOH failed; my hastily improvised apparatus was too leaky and a fair amount of chlorine escaped, judging from the smell; also the NaOH I used was not strong enough to deal with all the HCl vapor also coming off. The HCl solution was free from sediment and bright yellow, indicating Fe contamination no doubt from the crude MnO2. Now the plan is to determine Ba, after removal of Mn as MnO2 and Fe as Fe(OH)3, by precipitation of barium chromate and iodometric estimation of Cr(VI) in the BaCrO4. I haven't worked out what method might be best for determining Mn. I'm leaning towards removing the Fe by shaking with BaCO3, removing the Ba by precipitation of BaSO4, then determining Mn as MnNH4PO4.H2O. I wouldn't consider a gravimetric method but I did once succeed in determining Mg by precipitating MgNH4PO4.6H2O--my one stab at a gravimetric procedure--so I might be able to carry it off. If I had any sodium bismuthate I'd consider oxidation to permanganate; I have ammonium persulfate but Vogel states that it gives poor results in quantitative work.

One last note of possible interest: in trying to precipitate MnO2 from a sample, I decided to use NaOCl, which was a mistake. As noted above, NaOCl is capable of at least partial oxidation of Mn to permanganate, and that is what happened. Even prolonged boiling did not destroy all of it probably because there was always plenty of excess NaOCl around. Adding NH4Cl destroyed the excess hypochlorite and decolored the solution.

In any case, I'm fairly optimistic of the results--at least in getting as far as barium manganate. I see no reason why this preparation could not be scaled up considerably, either, which is definitely not true of fusion methods without going to a lot of trouble. The next trick will be to attempt preparing barium permanganate by gassing a suspension of BaMnO4 with carbon dioxide.

It might also be possible to dispense with the barium hydroxide and use a mixture of barium carbonate and sodium hydroxide. That would save me the effort of preparing Ba(OH)2.8H2O first; it's not terribly difficult but it is a tedious process.




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[*] posted on 3-9-2009 at 03:03


By the atached diagram, it doesn't look as if air would do the job under aqueous conditions.

For alkaline fusions, iron or nickel is acceptable. There will be some contamination, especially if chlorates or perchlorates are added as oxidiser - halogens are bad news even under alkaline conditions.



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[*] posted on 3-9-2009 at 07:23


Muthmann (Ber. 26 (1893) 1016) has a procedure for Ba(MnO4)2 starting from Ba(OH)2, KMnO4 and Ba(NO3)2. The product so made is said to be free of potassium. The same procedure appears in Vanino's book, Handbuch der präparativen Chemie.
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[*] posted on 3-9-2009 at 07:39


Quote: Originally posted by not_important  
By the atached diagram, it doesn't look as if air would do the job under aqueous conditions.

Aqueous standard conditions, to be sure. But the hydrothermal reactions are under conditions decidedly nonstandard. Have a look at this abstract:
Quote:
Chemische Technik (Leipzig, Germany) 17(8):493-494 (1965)
The effect of reaction parameters on the formation of potassium manganate(VI) from manganese dioxide and potassium hydroxide
Teske, Klaus; Lehmann, Hans Albert
In the system MnO2-KOH-O2, K3MnO4 alone is only formed at a K/Mn ratio of >3 and temps, of >500°. At lower temps. and K/Mn ratios, some K2MnO4 always formed. K3MnO4 was oxidized at 100-300° for 4 hrs. in O and a PH2O of 23 mm. (PH2O = pressure of H2O vapor). From ∼ zero at 100° the conversion to K2MnO4 was practically complete at 150-200°, and almost completely reversed at 300°. In air, the range of max. conversion was narrowed (175-200°), but increased (175-300°) by raising the PH2O to 107 mm. With increasing PH2O at 250° for 4 hrs., the conversion was almost complete at PH2O of 82 mm., and raised only slightly at PH2O of 150 mm. Addn. of an equimolar amt. of KOH lowered the initiation of the thermal decompn. of K2MnO4 from 500 to 200°, and caused almost complete decompn. at 250°. The addn. of more KOH caused only a slight further lowering of the decompn. temp. In the presence of moist O and KOH, K2MnO4 decompd. at 200-75°. The results may be partially explained by the equil.: 2K3MnO4 + H2O + 1/2O2 2K2MnO4 + 2KOH.
That will give you an idea of the conditions needed to effect oxidation to manganate. The phase diagram above is based off standard electrode potentials: 25 deg. C, 1 atm, 1M solutions. None of these is the case with a reaction at high temperature, increased pressure, and highly concentrated solutions used. Nor is the paper above the only reference to air-oxidation of MnO2 to manganate at high temperatures. A Japanese patent (JP 55085425) describes production of K2MnO4 at 275 deg. C and 3 atm with air the only oxidant.

Bear in mind also that preparing BaMnO4 instead of K2MnO4 means that, even if the equilibrium concentration of manganate is very small, the formation of barium manganate need not be hindered because the substance is highly insoluble.




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[*] posted on 5-9-2009 at 02:43


Very interesting stuff monoceros!

Quote:
there was more whitish matter that was probably BaCO3 from the impure Ba(OH)2. (This was later confirmed: a pinch of the solid fizzed briefly in weak acid.)


You had MnO2, KNO3 and Ba(OH)2 in a tightly closed vessel. I wonder where the CO2 would come from to form the carbonate?

That fizzling must be due to BaCO3 that was already present as impurity in your Ba(OH)2. Also keep in mind that KNO3 should be reduced to KNO2 and nitrite is decomposed by HCl. A solid containing a small amount of NO2- gives quite some fizzling with HCl but no visible NOx whereas a commercial sample of pure nitrite produces copious red fumes. I noticed this during my nitrate->nitrite reduction experiments.

If you find preparation of pure Ba(OH)2 is too hasslesome you could try MnO2+Ba(NO3)2+NaOH.

Anyways keep up the good work, you might be close to finding a working permanganate synthesis.

-----------------------------------------------------------------------------

Here's my short summary - the gist of this lengthy thread:

1. Direct oxidation with hypochlorite: It works but yield is low and permanganate has to be extracted with dry acetone (plenty of chloride & chlorate has to be removed)

2. Direct electrolytic oxidation of pyrolusite/psilomelane (pottery grade MnO2): Doesn't seem to work at all

3. Electrolytic oxidation of manganate made by fusion of MnO2 with KNO3: Works but rather low yield

4. Fusion of MnO2 with KClO3 and subsequent disproportionation of manganate: So far this is the only proven procedure which gives good yield at acceptable effort & expense: http://www.versuchschemie.de/htopic,10934,kaliumpermanganat....
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[*] posted on 5-9-2009 at 05:37


Quote: Originally posted by chief  
(thoroughly washed, filtrated, pre-glowed with some nitrate to remove the graphite first !)

That last item seems to me an unnecessary precaution. Why waste the KNO3? The graphite won't interfere anything and it will be left behind anyway when the melt is extracted.
Quote: Originally posted by chief  
Ah yes, the acidification: I preferred to ppt. out with Ba(NO3)2, to get the insoluble BaMnO4

For me that's going in the wrong direction. BaMnO4 is not the ultimate goal, KMnO4 is, and so a method for making BaMnO4 from KMnO4 is not useful to me.




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[*] posted on 7-9-2009 at 09:15


monoceros4 - nice to see someone else trying to get permanganates. I found it especially interesting that you need NaMnO4 and not the usual lab KMnO4. Barium manganate, being almost insoluble, is probably the only stable manganate - oxygen from the air seems to convert all others to permanagante and MnO2 except in strong alkaline solution.

You are certainly innovative! I like your test for potassium and I don't know why more amateurs do not use the microscope for a quick analysis of reaction products. Often the crystal forms can tell you what you have and very roughly, how much. I use this all the time - it only needs a drip and a bit of patience, plus very rudimentary knowlege of crystallography.

The solubility of NaMnO4 makes it very difficult to crystallize. I believe Na2MnO4 is near impossible. Your idea of a mixed sulphate/manganate is neat. Never realized the salts were isomorphous. This mirrors the formation of an alum from Mn(III) sulphate which can be isolated, although Mn(III) salts are very unstable.

Quote:
That last item seems to me an unnecessary precaution. Why waste the KNO3? The graphite won't interfere anything and it will be left behind anyway when the melt is extracted.


I think chief is right. Try taking some KMnO4 and mixing it with graphite. Apply a match or flame. Even in solution graphite is slowly oxidized by permanganate if you boil it.

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[*] posted on 3-2-2010 at 12:18


Quote: Originally posted by Formatik  
... Though as far as oxidation of HCl by KMnO4 goes, this is said by at least several references in Brethericks to have given a sharp explosion when conc. HCl was added to solid KMnO4, the explanation was that there is a remote possibility of the permanganate oxidizing the Cl2 to "chlorine oxide". ...


Finally understand now. For the preparation of Cl2 using KMnO4 and HCl, the KMnO4 should have so much water over it, just so that it's covered. Because action of hydrochloric acid onto the dry crystals easily forms small amounts of gaseous MnO3Cl, which can't be entierly removed by washing with water. In light, the permanganyl chloride also decomposes into Cl2O and MnO2. This all being found out by Kümmel, Wobig (Z. Elektrochem. 15 [1909] 252).
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Paddywhacker
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[*] posted on 16-5-2010 at 21:43


An equimolar amount of NaOH and KOH forms an eutectic that melts well below 200 degrees. So I tried that with MnO2, and tried to oxidize it to manganate with equimolar KNO3 and NaNO3 ... equimolar to keep the Na/K ratio constant and the melt molten.

The MnO2 reacted only slowly, and the nitrate seemed to have no effect. The melt turned green on the surface from reacting with oxygen, but the bulk stayed black/grey.

The photo is after adding the nitrate. You can see the dark bulk of the melt with blue manganate on the surface.
The melt bubbled slowly like thick boiling mud. Most of the MnO2 did not go into solution after 40 minutes of heating.

My reason for doing this was to make the manganate, then stir in zinc oxide to form a zincate/manganate melt. Then to cool, dissolve up in water, filter, and precipitate zinc manganate by lowering the pH with CO2. ZnMnO4 is a mild oxidizer like BaMnO4 with a very easy workup ... you just filter it off.

Edit --- another attempt to upload the photo.
The quantities were NaOH/KOH/MnO2/NaNO3/KNO3 44/61/47.2/23/27.5


[Edited on 17-5-2010 by Paddywhacker]

Sdc10045.jpg - 119kB

[Edited on 17-5-2010 by Paddywhacker]
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