Sciencemadness Discussion Board


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Omniquist - 15-2-2018 at 22:25

Quote: Originally posted by urenthesage  
Quote: Originally posted by MadHatter  
Given the recent CPSC(assholes) victory against FireFox, I'm sure that any convenient
method will be appreciated. KMnO4 sales by pyro suppliers are now limited to 1 LB a
year. There's still some OTC sources for now but we don't know how long that'll last.

So dont go to a pyro supplier. I got mine at a water depot, 5 lbs for about 45$. Farmers use it to take iron out of well water for cattle. They didnt even give me a second glance when I bought it.

I got a five pound bag from ace hardware. They don't have it in the store, at least in the state (USA, sorry everybody else not able to buy in ur country) of maryland. They put an order in for me and called about five days to a week later, it cost 39 us$ and needed my ID.
As for the calcium or sodium hypochlorite bleach I used with an old, likely dead d cell battery I had found. I didn't follow strict chemistry, but behaved like a cook, not measuring the hypochlorite bc I don't recall an exact measurement read. What ended up happening makes me concur with whoever said using hypochlorite is messy and difficult. I used about a half cup of standard bleach liquid (generic chlorox) and the crushed dark solid from the battery along with, again (and I DO know better, laziness was the lesson and why to avoid it) along with about a tablespoon of lye, which I found out may have been my problem. I boiled in measuring cup on hot plate, which took forever. Finally once it started it began to foam, bad. I took it off heat immediately and stirred like as if to make whipped cream or soufflé, and somehow succeeded in not spilling any. Thanks to whoever said that the majority of battery electrolyte is carbon bc my filter proved this. I used the top of a plastic water bottle, cut in half, as a funnel and a cotton ball stuffed in the mouth as a preliminary filter. No way was I using anything remotely dear to me for this project. What came thru was a pine green and aquatic blue homegenous transparent fluid, recognized as the manganate it was. However.... Pouring this same solution, this pretty green stuff which I meant to acidify promptly had turned into a dilute champagne/urine liquid, again transparent. Guessing it was mncl. I'll let u guys know how it goes using kno3, the battery innards but DRY, using the method employed by nurdrage, placing this mix in a container to be lit from underneath by torch. lowering ph with bicarbonate of potassium as sodium apparently eats the product before extraction. In hindsight I never have heard of sodium permanganate. Hmm.. I'll utfse. Thanks for your time and wisdom. I learn everyday from everyone on here.

CobaltChloride - 4-3-2018 at 07:04

Quote: Originally posted by DerAlte  
Forget K peroxydisulphate unless you have it on the shelf. Even then, it's marginal due to H2O2 formation.


I know this thread is pretty old, but could somebody tell me if they have tried the S2O8 2- oxidation? I can't seem to find its potential in alkaline solution. I would really like to know as Na2S2O8 is really cheap in my country, and so are 6% SO2 solution (to reduce MnO2 to soluble MnSO4) and KOH.

[Edited on 4-3-2018 by CobaltChloride]

clearly_not_atara - 9-1-2019 at 13:00

Quote: Originally posted by Romain  
Update: The crystals that formed from the evaporated anolyte are highly impure as evidenced by the presence of white crystals and brown crud among the needles of KMnO4. The crystals were dissolved in a small amount of water and the solution was heated up to boiling. It was filtered while hot on a preheated filter funnel (to remove MnO2, again). The filtrate yielded 0.37 g of small (and pure) needles of KMnO4 upon cooling.

(More KMnO4 could be extracted from the anolyte, but this run was just to prove that the method works and I didn't try to be particularly efficient in the extraction.)

The product was tested by adding a few drops of conc. H2SO4 to a crystal to form Mn2O7 and the reaction with ethanol was as expected: it ignited!

After everyone argued about it for ten years, Romain walked up, casually solved the problem on the first try, posted a picture, and nobody even noticed.

I have to say, I find the use of KCl as the potassium source to be perplexing. Chloride is obviously susceptible to oxidation by permanganate. There are a couple of other possibilities, but the most obvious IMHO is potassium acetate.

Properties of pure permanganic acid

chornedsnorkack - 20-2-2019 at 02:49

It can be and has been made. Stable and easily isolated.
Published in Jacs, in 1969. Sadly, only first page is visible. Has anyone reached the next pages?
To summarize the procedure:
  1. Produced pure dilute permanganic acid from stoichiometric amounts of dilute Ba(MnO4)2 and dilute H2SO4, and centrifugating the BaSO4 precipitate - about 0,3 M (3,6 %) permanganic acid
  2. Froze the acid. Did it by putting it into -75 Celsius acetone/dry ice bath and rotating, about 240 ml acid in 500 ml glass flask
  3. Vacuum evaporation. Their vacuum pump was capable of 0,001 Torr, and they sent the vapours through two traps of -75 Celsius acetone/dry ice and -193 Celsius liquid nitrogen. At first, pure ice was sublimed
  4. At some point, pink appeared in the trap - then the -75 Celsius trap was changed to collect separate fraction
  5. Finally, volatile fractions had left. The residue was found to be quite hygroscopic, solid, soluble - and pure HMnO4

    So, what is on the next page/s?
    Did they find out what happens when frozen solid HMnO4 is warmed above +1 Celsius? At what temperature does it melt? Congruently or incongruently?
    Has anyone repeated their synthesis?


HeYBrO - 27-3-2020 at 01:03

[Edited on 27-3-2020 by HeYBrO]

[Edited on 27-3-2020 by HeYBrO]

Attachment: pg2forguy.pdf (279kB)
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Bedlasky - 14-12-2020 at 21:13

I found on youtube this really interesting reaction - chemiluminecense from reaction between KMnO4 and Na[BH4]:

SuperOxide - 7-6-2021 at 11:05

Quick question (that I didn't think was deserving of its own thread, since I found this).

I purchased some Pot Perm Plus some time ago as a technical grade potassium permanganate source, and for just playing around with basic reactions/demonstrations, it's pure enough for that. The MSDS says its 60-100% KMnO4, but you can tell by the photos that there's clearly something else:

I figured maybe I could dissolve the KMnO4 in a solvent, filter it to get any insoluble crap out, then crystalize it. But from what I've gathered, KMnO4 isn't too soluble in acetone[src]. So perhaps just water? Though I know that heating it will cause it to decompose into MnO2.

So I suppose just distilled water and let it evaporate slowly at room temperature? Is that the best way?

P.S. I did look in the latest Purification of Laboratory Chemicals book for for potassium permanganate, but wasn't able to find anything on it.

[Edited on 7-6-2021 by SuperOxide]

Morgan - 7-6-2021 at 12:39

Quote: Originally posted by SuperOxide  
Quick question (that I didn't think was deserving of its own thread, since I found this).

I purchased some Pot Perm Plus some time ago as a technical grade potassium permanganate source, and for just playing around with basic reactions/demonstrations, it's pure enough for that. The MSDS says its 60-100% KMnO4, but you can tell by the photos that there's clearly something else:

I figured maybe I could dissolve the KMnO4 in a solvent, filter it to get any insoluble crap out, then crystalize it. But from what I've gathered, KMnO4 isn't too soluble in acetone[src]. So perhaps just water? Though I know that heating it will cause it to decompose into MnO2.

So I suppose just distilled water and let it evaporate slowly at room temperature? Is that the best way?

P.S. I did look in the latest Purification of Laboratory Chemicals book for for potassium permanganate, but wasn't able to find anything on it.

[Edited on 7-6-2021 by SuperOxide]

Isn't it maddening for a simple product to proclaim it could be 60-100% KMnO4 if you were trying to calibrate some repurposed use. Imagine if food was sold that way. Well I guess with potato chips it is as they sometimes list 3 different oils, possibly combined, "and/or" fried in.

[Edited on 7-6-2021 by Morgan]

SuperOxide - 7-6-2021 at 12:56

Quote: Originally posted by Morgan  
Isn't it maddening for a simple product to proclaim it could be 60-100% KMNO4 if you were trying to calibrate some repurposed use. Imagine if food was sold that way. Well I guess with potato chips it is as they sometimes list 3 different oils, possibly combined, "and/or" fried in.

For sure. I think they maybe do it because that gives them some room to vary the amounts of components in the product based on where they're sold/made at. If a certain chemical is cheaper in one city than another, then they may or may not use it for that area. But that's just a theory.

clearly_not_atara - 7-6-2021 at 16:39

Stop trying to dry permanganate. KMnO4 has a very low solubility at 0 C. The amount that remains in solution is almost always going to be less than what you would lose to decomposition on drying. If you really want to get more out, add excess potassium.

Antiswat - 14-2-2022 at 06:13

ok so i decided to delve into this very demonic preparation
NaOH melted with MnO2, supposedly a very slow reaction- this is when it just barely started to melt, i dont think you need 10-24 hours to form the sodium manganate- it is however blue, which explains its 2 steps away from permanganate

anyhow moving on, heres 3 solutions, from right to left: solution leached/decanted, acidified with HAc (this is a mistake!!), middle, blue, diluted leached solution without added acid, maybe some NaOH to keep solution basic and finally on the left, the acidified solution diluted down well
acidification does indeed turn it purple. infact this happens before it even hits pH 7, the manganate ion requires high basicity, otherwise it turns into permanganate, irreversibly
photo of sodium manganate, with adjusted color, very nice color.

as said, HAc acidification was a mistake. ITS ORGANIC, duh! look at that disgusting brown bottom layer, this was decanted off earlier, the MnO2 has formed since then
this is where it gets interesting, after concentrating down the solutions, decanting, and putting in fridge overnight, decanting the dark mystery solution i ran into this. a small amount of .. most likely potassium permanganate, with a load of MnO2- i didnt get to isolate it.

so, melted 25g MnO2 with 25g NaOH again (excess NaOH is better, it solidifies and the whole thing cant melt up if you use 1:1)
precipitated the MnO4 using barium chloride and barium nitrate- barium manganate is insoluble, and purple

i managed to produce just about 16 grammes of barium manganate with a faint purple/pink tint to it, it was dried at just about 100*C, maybe a bit less and i have a bunch more in solution, when its very thick solution it appears almost blue, but upon dilution with water you see that its really purple- the precipitate in photo was very volumnious for some reason- so its very obviously purple, that one precipitation yielded about 10 grammes of dry matter

so, i went ahead and added some phosphoric acid to it, as barium permanganate is soluble in water, this gave me a brown-red solution
red would mean that it lost some of its oxidation powers- oh no, how? appears phosphoric acid ruined the whole thing
H2SO4 would give me BaSO4- and would basically just stop the whole reaction
HCl is just a no-go with manganese chemistry for me. one could buy some CO2 water and use that, but i'd rather not, its what- 5% by weight, super unstable? would be better to find an alternative

i have an amount of times also dumped some KNO3 into the melt, but the result doesnt seem much different
i believe one could simply just boil the manganate dry, acidify for whenever you wanna use it and use it like that, it will contain x amount of carbonate/hydroxide however, but mangananate it is for sure. the barium manganate is said to work, just like that on its own for some organic reactions, a bit weaker than permanganate supposedly- this can also be acidified

main reason i made this was to use for emergency water disinfection, and since you dump manganate into water, when very dilute it does indeed turn purple- into permanganate, that should do. upon standing around it does turn red, probably because of organics in the water. as for water disinfection copper overnight will do too, for water disinfection the permanganate has to be just barely visibly tinted pink and i believe 12 hours standing time is recommended

MAYBE ill try to melt it down again, lets call it.. 20g NaOH 10g MnO2 10g KNO3, decant off, simmer down at low temperature, add in enough phosphoric acid to turn it acidic, add in 10g KCl. hm? extremely frustrating projection for a decantation fanatic like me, youre just completely blindfolded. a challenge, and eventually a major relief. i already got what i wanted, a bunch of color chemistry, done very easy, water disinfection. but now i want to crop out those diabolic KMnO4 crystals.

on a sidenote ammonium permanganate which i made long ago from KMnO4 + ammonium salt you shouldnt mention full name of yielded me long darkpurple spikey crystals which, when lit in one end would act like a small rocket, flying about in the air like a little, intoxicated insect, emitting fumes of MnO2, energetic potential for sure, risky single-compound rocket propellant certainly.

edit: i might add, adding phosphoric acid to the BaMnO4 gave me a clear solution that was pink. im not sure how, barium phosphate is supposed to be insoluble.

[Edited on 14-2-2022 by Antiswat]

Jome - 11-4-2022 at 09:12

A few days ago I did a quick experiment of trying to make KMnO4 through air oxidation, melting KOH:MnO2 in a 3,5g:2,5g ratio using a propane burner.

It quickly melts and bubbles (probably from water in the KOH), then the crucible was elevated around 10mm to give a lower temperature, and it was held there for a few minutes.

The hard "cake" was then hacked to pieces (difficult), and added to 50 or so mL of water. It looked disappointingly gray. It might've been "cold green" since that is the least visible color to my spectrum-challenged eyes, but to me it just looked gray. I decided to filter the solution through some cotton wool to be rid of remaining MnO2, the supernatant turned out to be weakly purple. A meal (1.5 h) later I returned, and to my surprise the supernatant had turned a stronger purple, and the part of the solution that had remained from the soon-to-be-filtered beaker (the decanted-from solids) were black. I decided to try these by pipetting a small amount into 100ml and this solution turned out to have even more permanganate, by the looks of it.

Last but not least (for that experiment), 24h later the remaining small pieces of the molten cake that were stuck to the crucible had turned soft, and when 100-200mg of this was added to water it turned so strongly purple that it looked black!


So, I decided to re-do the experiment paying a bit more attention to conditions, and then of course titrate it against a known reductant.

2.0g MnO2, 2.7g KOH was heated strongly until they melted and bubbled, they were then heated less strongly for 20 minutes. The cake was left for 23 hours and, having turned into a goop, was scraped off and added to 250ml of water. To my surprise, the first few pieces of the cake turned to water green, then quickly black. I guess this was either manganate or a combo of manganate and hydrated MnO2.

My reductant measuring standard is a 10.0g oven dried sodium oxalate with water added to a total weight of 1000g.

Some of this was diluted around fifty* times, and then for each 20 or so grams of the resulting solution, 12ish grams of a 10% sulfuric acid solution was used to acidify in a small beaker, the whole thing was then heated until "vapor visible/ finger pain, which I think is is 60-70 degrees of fluid temperature.

Another beaker had some of the permanganate solution, and it was weighted before and after enough of it was added to the first beaker to make the purple persist.

It turns out my yield based on MnO2 was 12.5%, with a cumulative error margin probably around 5 or so %, so 11.9-13.1... Permanganate really is pretty damn purple. It looked so promising... Perhaps I dead-burned it, perhaps higher partial pressure of O2 is needed for any reasonable yield. Do note that my MnO2 was ceramic grade, around 70% when using it for MnSO4 by nurdrages method, and contains iron, a small amount of chlorides (enough to smoke in the mentioned process) and perhaps 10% inert, clay-looking material. The MnO2 part is black and contains very fine particles, probably electrolytically produced MnO2 ("EMD"). The KOH was lab grade, but it has been sitting in its container for 15 years.

*yes, these rough sounding numbers were carefully written down and used in the calculations.

[Edited on 11-4-2022 by Jome]

Jome - 17-4-2022 at 23:44

Yesterday I performed two experiments, both using 2.0g of MnO2, to test various theories I had.

2.7g KOH/2.0g MnO2 burned until the bubbling stops (2-3 minutes), to test the "dead burn" theory.

4.7g KNO3/2.0g MnO2, because it would be fantastic not to have to use a difficult to acquire lab chemical (KOH) to produce another difficult to acquire lab chemical (KMnO4), instead possibly using the more available and relatively easy to make KNO3. I vaguely remember having heard some transition metal oxides catalyzing the decomposition of nitrates, and the equivalent process works well with Cr2O3.


The first preliminary test (the goop scraped off and dissolved in 250g water) turns the water green, it then stays green as opposed to the previous experiments with the same mix. Diluting this solution 40x looks gray to me, but cold greens are not my friend, and upon adding dilute H2SO4 it turns purple. but simple visual inspection (comparing to photos of known-concentration MnO4- solutions in the same type of beaker) shows that it is at most 1/3 or so as concentrated as the equivalent solution produced during my previous 20-min burn. I didn't titrate the solution since it was so obviously weaker than last time. "Dead burn theory", that the compound (manganate) forms quickly and is lost by excessive heating has thereby been disproved, at least under "my" conditions. Next I'll try a 30 min burn.

The hard cake remains a hard cake. Since potassium oxides are hygroscopic, this is a bad sign. The reaction did produce bubbling for a relatively long time (15 minutes), but in hindsight not particularly much. Small scale test of pulverized cake to water turns it gray, adding dilute H2SO4 turns it.... gray, possibly a small hint of purple but so weak that is barely visible. In MnO4- terms this means nothing formed. A hint of NOx could be smelled upon adding the dilute acid, but no visible effervescence, this means that the other possible product of this, KNO2, didn't form to any particular degree either. The cake consists of KNO3:MnO2, just molten together rather than a powder.
It is not impossible that this process could work, but it seems it would then require ridiculously high heat for a long time.

Jome - 25-4-2022 at 22:43

I did another experiment, a 30 minute burn. Titrated the same way, got 11-12%, which is basically the same result as the 20-minute burn, possibly marginally worse. I wonder what the reaction mechanics are, for some reason, the literature says, sodium manganate doesn't form from fusing NaOH with MnO2. Why would the MnO2 care what the counter ion is? Either the equivalent compound Na2MnO4 is thermally unstable, or potassium does something special like when it forms K2O2 much more readily than sodium forms Na2O2.

I'm aware that others have got the manganese-metal electrochemical route to work*, but Mn-metal isn't OTC enough to suffice IMO.

*has "anyone" produced any actual crystals, titrated them to see if they're really KMnO4 rather than MnO2?


I see two ways forwards: trying to isolate K2MnO4 and electrolyze that electrochemically, or raising the partial pressure of O2 above the melt by using a stream of oxygen. I'll post here if I do more experiments.

Alkoholvergiftung - 26-4-2022 at 08:05

to Antiswat.In the early days of industrialisation they melted NaOH and MnO2 at 400 C together and than the added to the solution MgSo4 solution or MgCl2,CaCl2.
3Na2MnO4+ 2MgSO4 +2H2O= Na2Mn2O8 +Na2SO4+2Mg(OH)+MnO2 (i know old formula permanganetes are written NaMnO4 they wrote it so that time.)

The final product wasnt clean but you are on the right way it seems. :)
methode from Tessie du Mothay.

Electrolytical the used Manganesecarbide anodes.

Chemical way to produce permanganates is only when you use Ozone Gas on the manganetes it goes 1to1 into permanganetes

[Edited on 26-4-2022 by Alkoholvergiftung]

[Edited on 26-4-2022 by Alkoholvergiftung]

Antiswat - 9-7-2022 at 04:59

i bought some carbonated water to try for this, but i also acquired a CO2 extinguisher, as CO2 appears to be superior for this reaction to complete
i believe adequate filtration is a crucial step in homelab to have KMnO4 produced, but also KOH to start from

many organic reactions calls for manganate, or basified permanganate, so maybe making KMnO4 is really just a challenge?

i found something on KMnO4 from a textbook and ill translate it:
"Potassium manganate, K2MnO4, is a greencolored salt thats formed when MnO2 is melted with KOH and a bit of salpeter, KNO3. If chlorine or CO2 is bubbled into the solution potassium permanganate, KMnO4 is acquired."

so- K2MnO4 into KMnO4 requires chlorine gas or CO2.
but at this point you want to concentrate the more stable manganate solution, after copious filtrations.. and then pump in CO2, maybe adding in solid CO2
solid CO2 can be extracted from CO2 extinguishers by wrapping the nozzle in a buch of cloth in a bucket and firing it rapidly

i got a vacuum filtration system since i worked on this last so i might give it another go.

NaOH + MnO2, dissolve in boiling water, vacuum filter, concentrate down and add CO2 and cool it down slowly? decant liquid and check for crystals

Keras - 11-10-2022 at 11:18

I’m not sure what the problem is here?
I tried this reaction myself yesterday. Stainless steel cup loaded with some (didn't weigh) flakes of KOH. Waited until it fused, added MnO₂ extracted from D (formerly LR20) new alkaline batteries. Because it set into a goo, I added more KOH, until it remained liquid, and I periodically stirred with a spatula. I didn't use any solid oxidiser to avoid contaminating the product.

That was left to cool, I used a hammer on the back of the cup to force it out, dissolved the pieces in hot water. Filtering. Got unreacted MnO₂, don’t really care, and a deep green, clear solution (btw, this stained my glass frit. No matter how much HCl I added, there were MnO₂ particles left. Wish I had glass wool filter left. It maybe useful to add Celite).

Next step is to neutralise the excess potash. This must not be done with HCl, because the permanganate will oxidise the Cl⁻ anion and revert to a lower state of oxidation. Sulphuric acid is used.

Took a sample at some point, dissolved it in more water. With a bit more of sulphuric acid it turned lovely purple (see picture).

Did another filtering to remove the MnO₂ formed by the disproportionation, then boiled the solution until I got purple crystals filling around to 30/40 mL of my beaker (see other picture). Had no time to go further. Obviously, there is potassium sulphate there, but also a fair amount of KMnO₄. It needs recrystallisation. I hope I can do that next week-end.

[Edited on 11-10-2022 by Keras]

IMG_1130.JPG - 1.4MBIMG_1131.JPG - 1.5MB


RU_KLO - 19-12-2022 at 07:14

the main Idea was:

2 MnSO4 + 5 PbO2 + 3 H2SO4 → 2 HMnO4 + 5 PbSO4 + 2 H2O

HMnO4 + KOH -> KMnO4 + H2O

It did not work.

probably because:

PbO2 was not pure (was from a lead-acid battery)
PbO2 is almost insoluble

I got a very light violet color (more like pink), but probably is a mixture from light black (PbO2 solution) + pink (MnSO4 solution).

[Edited on 19-12-2022 by RU_KLO]

Bedlasky - 27-12-2022 at 14:21

Mn2+ is pink (not violet) only at very high concentrations. So you surely got some permanganate. Did you heat the solution?

[Edited on 28-12-2022 by Bedlasky]

RU_KLO - 28-12-2022 at 04:49

No, did not heat it.

here are some (bad) pics - taken from a video.

the first pic: trying to make a solution of Lead Dioxide (checking solubility). 0.1gr in aprox 10ml 1M NaOH, 1M H2SO4, H2O) ambient temperature.
Not fully disolved. Did not weight not disolved PbO2, but in all there were some.
It seems more a suspension than a solution.

then aprox 10ml 2M MnSO4 solution was added.
the "pink-violet" solution was seen in H2SO4, a precipitate was seen on NaOH solution (probably Mn(OH)2)

from what I saw, if it was permanganate, it was very diluted.... So did not proceed further.

My thinking was: black color suspention + pink color (Mn(OH)2) = Pink/violet/brown color (or the color you see in te picture) and not violet from permanganate.

(Did not add KOH, because of precipitation in NaOH solution. - so I expected the same...)

Is there a way to increase solubility of PbO2?

(sorry I do not know how to add pics, because "Add Image" button did not work as expected

[Edited on 28-12-2022 by RU_KLO]

[Edited on 28-12-2022 by RU_KLO]

[Edited on 28-12-2022 by RU_KLO]

Bedlasky - 28-12-2022 at 17:38

PbO2 is insoluble. Just use suspension of it. Heat it - everytime I did oxidation of Mn(II) to permanganate, I must heat the reaction (I tried hypochlorite and persulfate - both need heating, I also have bismuthate so I will try that in the future). And PbO2 work as an oxidizer only in strongly acidic environment. So addition of NaOH stop the reaction. I personally would use stronger H2SO4 than 1M - maybe something around 20%? 1M acid should do the work in theory, but sometimes reactions need more acidic conditions.
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