Sciencemadness Discussion Board


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Xenoid - 22-8-2007 at 23:27

I tried something like this, in a half hearted way, a few days ago.

Sorry I didn't take a photo of it in operation but it was really just a try out to see what would happen and that the arrangement would work.

I used the set-up shown below, a stainless steel bowl was the cathode and the small porous pot (plant pot) was the anode partition. I foolishly used a gouging rod (carbon) anode. I knew it would dissolve but I thought it would last longer than it did (~18 hours).
I filled both sections of the cell with 2M NaOH solution (I used NaOH because I have plenty of it, but only a small quantity of KOH). I put a few heaped teaspoons of MnO2 in the central anode pot, along with a magnetic stirring bar, the slurry stirred up nicely! The cell was attached to a constant voltage lab. power supply, and initially ~30 volts was required to get any current flowing (because the porous pot was dry). As the electrolyte soaked through, I was able to turn the current up to 2 amps and for this current, the voltage fell to about 9 volts where it stabilised. After about an hour I extracted an aliquot of a few mls (shown in the test tube on the left after settling). It appeared slightly greener than plain water. I then decided to increase the strength of the NaOH in the central pot, and added more solid NaOH which brought the strength up to about 3Molar. The other two test tubes show aliquots after about 4 hours and after 18 hours. The cell came to premature end after 18 hours when the carbon rod was totally eroded away. I am assuming that the greenish colour of the aliquots was manganate, there did not seem to be any permanganate production. The electrolyte in the cathode partition remained clear during the whole time.

NOTE; Also in the image are two composite epoxy-MnO2 anodes. The one on the left was cast the one on the right was molded (rolled), they have steel rods down the centre. I used a very minimum amount of epoxy, slightly diluted with thinners, but even so they both have a resistance of 2-3 Megohms.... :(
So don't bother going down this line of research!

When the cell was in operation, the hydrogen from the open cathode chamber was carrying quite a lot of NaOH mist. This tends "to catch in the back of the throat". My first modification will be to use the lid (shown at the bottom of the image) which came with the stainless steel container, and cut a hole in the centre of it to make a tight fit with the plant pot. I can then lead the fumes away using some PVC hose.

I am also in the fortunate position of having a little Pt foil to use for anodes, so I will try this next time;

NOTE; Porous "plant pots" are quite soft, after only 18 hours, the magnetic stirrer bar had ground a depression about one third the way through the base. It will be nescessary to put a piece of hard plastic or something similar in the bottom of the plant pot.

Hope this helps!

Regards, Xenoid

Permanganate-Cell.jpg - 17kB

DerAlte - 23-8-2007 at 08:52

Xenoid, very interesting! Your set up is just about what I was thinking of, bar the carbon anode. I was thinking of some sort of compressed MnO2 as they make in batteries. In fact I once tried the cathode part of an new alkaline cell, after scooping out the zinc and binder and carefully extracting it from the case. Since KOH is the electrolyte, the only problem should be the 10% graphite added for conduction. In an undivided cell I did get a red coloration before the MnO2 fell to pieces, which was quickly.

Epoxy does tend to insulate any filler particles. Silver epoxy is used to seal RF leakage in electronic equipment and is conductive at DC but I’m sure the grains have to contact for it to work. Another material is a silver filled silicone (?) rubber .gasket used the same way. Whether any such would ever work with MnO2 I have no idea.

The results seem disappointing. Could you try a nickel anode in the same set up?

Thanks for the patents, ciscosdad and not_important. Haven’t had time to peruse them yet.

Der Alte.

Filemon - 23-8-2007 at 13:11

Originally posted by DerAlte

The starting materials are KClO, MnO2, KClO and possibly KOH. We have the following reactions:

2MnO2 + 3KClO + K2CO3 --> 2KMnO4 + 3KCl + CO2

The reaction with nitrate the following one is?

MnO2 + 2KNO3 + 2KOH => K2MnO4 + 2KNO2 + H2O

DerAlte - 23-8-2007 at 13:49

To balance it,

MnO2 + KNO3 + 2KOH => K2MnO4 + KNO2 + H2O

In the fusion method, yes. It won't work in solution because the nitrate --> nitrite redox potential is very weak in alkaline solution. However, this applies only to a water solution - one can't predict a solid fusion reaction from this. Since the H2O is able to escape, the reaction should go to the right, as your equation shows, by Le Chatelier's principle.


Der Alte

MnO2 Electrodes etc

ciscosdad - 23-8-2007 at 15:45

I have seen a reference that specifies Portland Cement as a suitable binder for MnO2 electrodes. It was in a patent I read recently, but memory fails me. It will probably be more porous than epoxy. A metal rod down the centre would be a good idea.
However, Stainless steel and Iron will do for the electrodes for this, so we can keep the exotic electrodes for something that needs them.
Ref solibility of KMnO4:This may or may not be a problem depending on where it deposites after formation. If it crystallises in solution and becomes a part of the slurry, then it may not be an issue. However real life is unlikely to be that kind to us, so it will probably cover electrodes, stirrer, thermometer or whatever.
Perhaps the approach will be to use NaOH, then after the run is finished, add sufficient KOH to form the expected amount of KMnO4. Filter and the solution is now prinicpally NaOH plus residual MnO2, NaMnO4 and small amounts of KMnO4. (I expect that the MnO4 formed will hydrolyse quite quickly, so we will probably have to accept the presence of more or less MnO2 in the product).
This soln will now be ready to take more MnO2 and returned to the cell for a rerun. This will work in litre quantities,and from the patent will produce around 1 Mole of KMnO4 per day. Scaleup to ~ 5 litre quantities should not be difficult. If this works, the method will make KMnO4 as accessible as Chlorates.
The only hangup I see at the moment is a reliable way of filtering the solution. Does anyone have any suggestions? Everything I can think of that is readily available will be attacked.:(

DerAlte - 23-8-2007 at 17:21

It sounds a nasty brew, but in my (small) experience as an amateur, I think the erosion of glass by alkali hydroxides is overstressed. The permangantes do eat up ordinary filter paper. I'd use a glass frit filter if I had one, but a glass plug in an ordinary filter fummel does OK. A Buchner funnel and a frit or tightly woven gass mats would also do better, under suction. The MnO2 usually clogs glass wool in a funnel. A funnel of HDPE seems to be fairly immune to attack by KMnO4 - I have not observed any problems using them.

Haven't read your patents yet, folks - I've been a bit busy but will do so tonight.


Der Alte.

DerAlte - 24-8-2007 at 12:39

I have read the patent. I have the following comments:

(1) No diaphragm is mentioned. No details of cell structure were given. I assume that anode and cathode were well spaced to avoid dissolved MnO4-, MnO4 - - or MnO4- - - ions intermixing too seriously with cathode products likely to reduce them .back to MnO2.

(2) I found the restriction of the KOH concentration to 15-25% interesting.. I IM solution of KOH has roughly 5.6%, ignoring density, so these solutions are of the order of 2.5 to 4.5 M. Such a solution has a pH of 14+ but less than 15. If the pH of a manganate solution is dropped to say to below 10-11 (I’m not sure exactly to what) then it goes to permanganate 3MnO4-- + 4H+ <-- --> 2MnO4- + MnO2 + 2H2O. Increasing the pH by adding alkali reverses this (a nice Hey Presto! demo for grandchildren!). Manganate should be stable in the 15-25% KOH solution but not permanganate.

(3) In the fusion with KNO3, the authors assumed Hypo-manganate was produced. This is a new one to me! MnO4--- is definitely very unstable in aqueous solution. It oxidizes to manganate as soon as you produce it except in very strong alkaline solution. See Brauer for a way of producing it (It works – I’ve tried it, but things have to be done just right.).

The redox potential of MnO4 - - against MnO4- - - is only about 0.27 V while the MnO4- --> MnO4 - - transition gives about 0.57V. But these mean nothing except a general trend for any non-aqueous condition, such as in the dry state, fused or in some other non-polar solvent.

I always assumed the hypomanganate was very unstable above 0C and am surprised to learn of its existence in a fused mixture at 300+C. But then again the industrial process as reported in many chemical books says that manganate exists at 500C, a low red heat. CRC says it decomposes at 190C. and permanganate at 200-250C is often quoted. I believe the latter, and you can show it by heating KMnO4 in an oven till it crackles. These compounds may be stable in the presence of OH_ ipns, even dry or fused.

I can believe the use of a Ni anode, even perhaps Stainless Steel. But iron? It oxidizes rather easily - in moist air, e.g. It reacts with dilute acids to give H2. Nickel is far less reactive, displacing H from acids very slightly. But iron can be passivated by strong nitric acid, so maybe it passivates under these conditions.

I was aslo fascinated by the realtively low current densities - 500A/m^2 is only 50ma/cm^2. And aslo the reuirement for a high temperature. I have to re-read the thing again to get the full gist...

Well, all patents are written by very specialized lawyers to be as obscure and misleading as possible while claiming that they also cover all possible variants of the method/invention (and rightly so). So one has to read between the lines a bit…

Der Alte

[Edited on 24-8-2007 by DerAlte]

Xenoid - 24-8-2007 at 15:01

Hi, DerAlte

Well I have set up a simple cell to try out the Japanese method.

It's based on their 1 litre beaker method, but is only 1/5th the quantity.
I'm using a 250ml beaker with 200ml of 3.5 molar KOH solution and I added about 18g of MnO2. The anode is stainless steel sheet, wrapped around the inside of the beaker. The anode area is about 108 cm^2. The cathode is a 10mm steel rod in the centre of the beaker. The operating temperature is 70 oC, and stirrer setting is 750 rpm.
The cell is running at ~2 volts / 2 amps.

After running for 10 hours, there is no sign of any permanganate production, the settled solutions taken as aliquots are quite clear.

Note: The bare copper wire around the top of the cell is not an integral part of its construction. It is holding the black plastic cathode support in place! I found with magnetic stirring that the steel cathode was bouncing around along with the stirrer bar (flea).

Note: 3.5 molar KOH + 70 oC = NASTY

My old eyes are tired from the small print, and my brain is sore from the obfuscation in this patent. But I get the impression that this will only work with the precipitated, active form of MnO2. Is this obvious, am I missing something, or does everyone know this already.

Ist embodiment: MnO2 (tetravalent Mn) produced as a by-product of other Mn processes (ie. active form of MnO2)

2nd embodiment: Pentavalent Mn produced by fusing Mn ore with alkali hydroxide and alkali nitrate.

I assume that my powdered pottery grade MnO2 is actually just pyrolusite or psilomelane, and will be unreactive in this reaction without further chemical treatment. That is, fusing with alkali hydroxide and nitrate.

So really we are no further down the track, the industrial process is fusion of MnO2 ore, followed by electrolysis in KOH. This was already known! This patent process is essentially the same as the current industrial process.

I have retrieved several old lantern batteries from the compost heap, where I discarded them after removing the carbon rods and zinc metal. I will have a try with the MnO2 recovered from them, it may be more reactive.

Regards, Xenoid

Cell3.jpg - 22kB

DerAlte - 24-8-2007 at 19:37

Great try, Xenoid! Your turn around time is spectacular and your set up looks highly professional. I can't move that fast, I don't call myself Der Alte for nothing. I don't even have any KOH at present. You've tried it and I sit here dreaming!

As indicated in my last post, I'm always a bit sceptical about patents. (Having worked on one or two with a patent lawyer, I know why!)

Your anode current density I estimate at 363 ma/cm^2, a bit higher than the patent. Not that I believe this is going to make much difference.

That black liquid looks perfect. You could have fooled me that it was stiff with KMnO4! No doubt about the stirring. You have a thermometer and are at temp. A negative result says only one thing - either we're missing something or the patent is sheer Japanese bull shit. Somehow I think there is something there we have missed, and the patent information is probably both optimistic (you wouldn't expect it to be pessimistic!) and purposely misleading. I am thinking hard but nothing occurs.

As for the impurity of pottery MnO2, I beleive it is usually native ore but I could be wrong. I use battery MnO2 purified as shown earlier in the thread. And I don't like the idea of an iron anode, as I said before. Any obvious corrosion? I rather see the SS as the anode or better nickel.

In scientific experimentation a negative result is nether good not bad, Xenoid. Both provide information. You, Ballermatz , Cicsosdad at al. have livened up a moribund thread muchly. And we haven't finihed yet. A clever chemist can do (almost) anything! Even amateurs (professional advice welcomed, of course),

Regards to all,

Der Alte.

not_important - 24-8-2007 at 20:22

If you're using battery MnO2, there's a good chance that it still has some carbon in it that will have to be oxidised before any yield of MnO4(-1) can be had. Likewise much iron will compete, certainly be oxidised to Fe(III) if not higher.

Doing a small run with hydrated MnO2 would be informative. An alternative is the MnO(OH) and mixed (III)/(IV) oxides formed by air oxidation of alkaline suspensions of Mn(OH)2. This has been used in the old method of making chlorine from HCl and MnO2, treating the MnCl2 solution left with alkali and air to produce higher oxidation state hydrated oxides that can be reused to make more chlorine.

People playing with managanese containing stuff that kicks up a fine spray, or involved in fusion of managanese compounds, might wantto read this

Potters Manganese Toxicity

Xenoid - 24-8-2007 at 22:19

Hi DerAlte

The current is 2000mA/108 cm^2 = 18.5 mA/cm^2. It's actually on the low side, the range given was 5 - 50 mA/cm^2

Note: The anode is SS on the outside of the beaker. The cathode is iron (steel) in the centre, haven't noticed any erosion yet!

I am now running the cell with MnO2 from Zinc/Carbon batteries. It was boiled 3 times with water and dried and reground.

After 4 hours, still no sign of permanganate... :(

Hi not_important

This is by no means a final cell design, I just put it together quickley to check out the method. I would envisage a final cell design comprising a tall SS container, with a tightly fitting transparent plastic lid. The cell needs to be sealed not just because of the toxic mist, but also because of the high water loss by evaporation at 70 oC. The electrolyte level in the cell shown in the image is dropping by about 5mm / hour. A properly built cell would require a vent tube, which would also serve to condense the water vapour and feed it back into the cell.

""treating the MnCl2 solution left with alkali and air to produce higher oxidation state hydrated oxides that can be reused to make more chlorine.""

Yes, I am thinking of reacting some pottery grade manganese carbonate with HCl, to try exactly this, in my next attempt.

I am still not sure what the problem is with fusing MnO2 with KOH and KNO3 in say, a SS bowl, and then running the resulting K-manganate through a cell such as I have built. Other than the fact that I don't have solid KOH, only 195g/litre (3.5M) liquid KOH "drain cleaner", I would be trying that next!

Regards, Xenoid

[Edited on 24-8-2007 by Xenoid]

DerAlte - 24-8-2007 at 22:45

@not_important: agree absolutely about getting rid of the carbon and the iron. I show how to earlier in the thread. As regards toxicity, reasonable care should be exercised but the toxicity isn't extremely high unless you ingest or go intravenous in a mad fit! People seem to forget about the spray from reaction. Manganese has a distinct smell to me, like iron compounds. Neither are particularly volatile per se. Of ouurse, all Mn salts and permangantes can stain horribly, like iron salts. I wear gloves with both.

@Xenoid _ wow! You are even faster off the mark than I thought. I misunderstood or misread about anode/cathode so of course you are right there. I worked out the current density thinking the centre post was anode - force of habit! An iron cathode doesn't bother me.

I wouldn't be too happy about using Zn/C stuff as MnO2. (See Not_important's post above) Those old lantern battries - i use them too, but prefer the modern alkaline cells, except you don't get a free carbon rod - the Zn/C black stuff is chock full of all sorts of crap. Including zinc, ammonium, etc. Boiling helps, but I never use it without dissoving out any metals with acetic acid, 5%, or very dilute hydrochloric. There's not much actual MnO2 either - it's hydrated lower oxides, hydroxides, etc when it has done it's job. To get Mn O2, convert to chloride with strong HCl (chlorine produced) and react this with NaOCl. Details in a prior post in this thread, IIRC.

Regards, Der Alte

DerAlte - 24-8-2007 at 22:50

I've had this post on the back burner but I’ve been too interested in the possibilities of an electrolytic method to add to the idea of using PBO2. For completeness I continue on that theme:
Der Alte said

Possible half-reactions are:

PbO2(s) + 4H+ + 2e- --> Pb++ + 2H20 1.46V
Mn++ + 4H20 --> MnO4- +8H+ +5e- -1.51V

Which is close but just shy of satisfactory by 0.05V. Combine the ½ reactions to eliminate charge. We get:

5PbO2 + 4H+ + 2Mn++ <--> 5Pb++ + 2H2O + 2MnO4-

At equilibrium, then,

[MnO4-] = [PbO2]^5/2 [H+]^2 [Mn++] / [Pb++]^5/2 [H2O]

which tells us that increasing [H+] and/or reducing [Pb++] we could tip the balance in favor of permanganate production. And lead salts, such as the sulphate or carbonate, are quite insoluble. I must think about this one a bit more…

OK. I have thought.

Suppose we use MnSO4 as the Mn++ source:

5PbO2 + 4H+ + 2MnSO4 < -- > 2PbSO4(s) + 2H2O + 2MnO4- +3Pb++

or, since Pb(MnO4)2 is soluble (AFAIK), one Pb++ could go to permanganate – not that we want it... In which case, what happens to the other two?

Well, we needed a source of H+ ions, i.e. an acid. So use H2SO4 and precipitate the remaining Pb++ ions? That gives us

5PbO2 + 2H2SO4 + 2MnSO4 < -- > 4PbSO4(s) + 2H2O + Pb(MnO4)2

It might work. Searching the literature available to me, this reaction apparently does work with nitric acid. Solutions should be dilute and boiled for some time. I’d like to try this one. But I’ve got to get (or make) HNO3 to provide me with a soluble salt for lead, and also carry out this reaction. Lack of H2SO4 really hampers one.
Another dream perhaps…

Let’s continue on the electrolytic route or revisit the fusion route, in light of the Japanese patent. Don’t waste any Cl2 either – it can be used to oxidize manganate to permanganate... or produce a hypochlorite.

Der Alte.


ciscosdad - 27-8-2007 at 15:15

I have assembled a list of patents that I've found. I hope they are of interest.
The list is likely not exhaustive, so if there are other relevant patents not included please speak up. Other relevant references would be of interest also.
From my reading, it is difficult to find practical information in any one place. Perhaps a collection of resources would help us solve the problems.

Preparation Of Permanganates

3 986 941 1976 MnO2 to KMnO4 direct in 20% Aqueous Soln 90 Deg C
185 214 1876 Condy’s Patent
326 657 1885 Electrolytic (cell membrane) very vague
1 281 085 1918 NaMnO4 then KMnO4 precipitation
1 291 680 1919 Mn Metal Electrodes
1 337 239 1920 Permanganates by Chlorine oxidation
1 360 700 1920 KmnO4 from Mn Metal electrodes
1 377 485 1921 NaMnO4 then KMnO4 Precipitation.
1 453 562 1923 Ref 1 544 115
1 542 538 1925 Purification of Acetone (using KMO4)
1 544 115 1925 Alkaline Earth Permanganates
1 826 594 1931 Manganate via KOH,Air,180DegC
2 504 129 1950 Mg and Zn Permanganates Via Al
2 504 130 1950 Alkaline Earth Permanganates Via Al
2 504 131 1950 Alkali and Ammonium Permanganates Via Al
Previous 3 are variation on the same theme
2 424 392 1947 Manganate via Mn metal electrode
2 843 537 1958 Electrolytic Manganate to Permanganate
2 908 620 1959 Ref 2 843 537
2 940 821 1960 Manganates..slow MnO2 addition..KOH Melt
2 940 822 1960 Manganates..slow MnO2 addition..KOH Melt
2 940 823 1960 Manganates..slow MnO2 addition..KOH Melt
Previous 3 are variation on the same theme
3 172 830 1961 Removal of Impurities
3 062 734 1962 Cell and Electrode
3 210 284 1965 Stabilization of alkaline MnO4 solutions
3 293 160 1966 Using Manganese Metal electrode
4 085 191 1978 Potassium Recovery using Lime
4 117 080 1978 Regen of Gas scrubbing solution
4 592 852 1984 Regenerating KMnO4 solution (secondary oxidant)
4.853 095 1989 Regeneration of Etchant Solutions
4 911 802 1990 Manganate to permanganate regeneration (cell membrane)
5 011 672 1991 Description of Fused KOH/Air/MnO2
5 660 712 1997 Electrolytic using cell membrane

I agree with Der Alte about some of them being amazingly obscure. I guess one must expect it given that lawyers were involved.
However, it is most productive to read and re read (between the lines!). I note that even the patents that deal with a different basic process will often give insights on the processes and the practical chemistry.
This is rather a big list, so if anyone wants a copy of any of the patents listed I can P2P if you have difficulty with google patents.

Congratulations Xenoid. Lovely work.
I note that the alkali can be NaOH or KOH. I wonder if a mixture could be used (if you wanted to adjust the concentration to very close to 20% with solid NaOH)
Did you put a small amount of (catalytic) KMnO4 in as a starter?
If there is no KMnO4 avilable at first, perhaps the first fusion step with KNO3 to manganate might give better results. I believe there is a small amount of KMnO4 in the resulting equilibrium mixture.

Xenoid - 27-8-2007 at 17:34

Originally posted by ciscosdad

Congratulations Xenoid. Lovely work.
I note that the alkali can be NaOH or KOH. I wonder if a mixture could be used (if you wanted to adjust the concentration to very close to 20% with solid NaOH)
Did you put a small amount of (catalytic) KMnO4 in as a starter?
If there is no KMnO4 avilable at first, perhaps the first fusion step with KNO3 to manganate might give better results. I believe there is a small amount of KMnO4 in the resulting equilibrium mixture.

Hmmm... Lovely work, but it didn't produce any KMnO4. :(

Neither pottery grade MnO2 or MnO2 crud from Zn/Carbon batteries has worked. I was surprised by the 2nd example not producing even a hint of pink after 8 hours.

I didn't add any preliminary KMnO4, but I did add a little KClO4, its only there to condition the electrode(s). I think KClO4 was mentioned as an alternative, wasn't it!

Regarding the "Japanese Patent", note my comments a few posts back, when you wade through the obfuscation (great word) the patent involves 2 "embodiments";

1st embodiment: MnO2 (tetravalent Mn) produced as a by-product of other Mn processes (ie. active form of MnO2). They list several reactions where this is formed as a biproduct for recycling. Nitrogen oxide scrubbing and saccharin production are mentioned. One needs MnO2 produced by an aqueous process, not a pyrolytic one. The pottery grade MnO2 is particularly hard and "gritty", the bottom of the beaker was frosted from the grinding action of the MnO2 slurry!

2nd embodiment: Pentavalent Mn produced by fusing Mn ore with alkali hydroxide and alkali nitrate. This appears similar to the normal, modern commercial process.

I am going to extract some MnO2 "crud" from an Alkaline cell and see if it works.
Failing that, I have been busy building an apparatus for producing Na-manganate (as per the "Japanese Patent", 2nd last page , example 8). I will then run the manganate in the electrolytic cell.

PS. I ordered some KMnO4 from a garden supplier the other day, it has now arrived 1Kg for $15 not too bad, and it looks to be good quality, nice crystals. :D

Regards, Xenoid

Potassium Permanganate

ciscosdad - 27-8-2007 at 20:16

Hi Xenoid,
I had a fit of jealousy when you told us you could get the KMnO4 in Kg quantities (!) from a garden supplier. What did you say you were going to use it for? Here in Oz the only legitimate "gardening/backyard "purpose I have come across is as a disinfectant in aquaculture. Of course none of the aquarium suppliers have got it either. Is there any other? I'm hoping to identify another source.

Xenoid - 27-8-2007 at 21:05

Originally posted by ciscosdad

I had a fit of jealousy when you told us you could get the KMnO4 in Kg quantities (!) from a garden supplier. What did you say you were going to use it for?

Yeah! It's the only place I have found in NZ. Things have really tightened up here in the last year or so, (Hazardous Substances Act or some such). Soon it will be as bad as the OZ police state. Now I can't even buy 25Kg bags of KNO3 from the fertilizer works any more. You have to be a "licensed operator". The Condy's crystals was listed "for treating "club root" in brassicas and can be used for moss in lawns and Carrot Fly deterrent. Can be used to sterilize soil". Apparently a few crystals are dropped in the hole before planting brassicas (brussel sprouts, cauliflower etc.). It was available as 150g for $6 or a whopping 3Kg for $45, too much for me. :o

Regards, Xenoid

DerAlte - 27-8-2007 at 21:07

@Ciscosdad - good grief, enough reading for a year! Might be something among all that obfuscation (the only adjective for patents). A use for KNO4 is for greensand filters for wells. Buggered if I know what greensand is! But it sounds good. Try hardware stores as well. Do-it -yourself places selling insecticides also might have it. May be a bit impure but easily purified by the usual methods.

@Xenoid: Alkaline cells are best. If sacrificed unused, they only have KOH as contaminant. Used you get zinc contaminant and also reduced dioxide as trivalent Mn, so you ought to purify first to get MnO2.

And yes, your efforts are splendid, even if they didn't give the wanted result. You may have saved others endless trouble in believing Japanese obfuscation, which may exceed occidental obfuscation because of the known inscrutability of the Far Eastener. Maybe a little Zen would help...


Der Alte

Xenoid - 27-8-2007 at 22:55

@ DerAlte
I'm afraid I went ahead with my latest attempt, without checking out the electrochemistry of alkaline cells. I picked out an old "D" cell, totally flat (about .2V with no load). I cut it open and peeled off the outer Mn layer. I was expecting it to be black (MnO2 + graphite) but it was a dun brown with a few black streaks. I immediately assumed (wrongly) that it was Mn(OH)2 and that it had been created in the half reaction at the cathode during discharge.
I put this brown material in the cell under the same operating conditions as outlined previously. I had it running with the stirrer on before it heated up and it was just a brown slurry, interestingly when I turned the current on, all the graphite suddenly appeared, forming silvery bubbles on the surface. The cell has been running for about 4 hours now, but there is no sign of pink permanganate.
After a little research, I now know the half reaction is as follows;

2 MnO2 + 2e- + H2O ——-> Mn2O3 + 2 OH-

I always thought Mn2O3 was black, like MnO2.
I am loathe to cut open a fresh alkaline cell at the present time, just to obtain some MnO2.
What is needed for this reaction to work is hydrated MnO2, precipitated from a reaction in solution - MnO2.H2O this is the so-called manganous acid H2MnO3, This is what is used in the Japanese reaction.
I think it has already been mentioned earlier in this thread, that this can be obtained by treating an aqueous solution of a manganous salt with an alkaline hypochlorite solution (aka. bleach). I have 5Kg of MnSO4 which I bought cheaply from a hydroponics store when I was messing with MnO2 plated anodes. I'll treat some of this with some bleach to get some hydrated MnO2. And then try this in the cell, I'm sure it will work!... ;)

Yes, I just tried this reaction, and it works well. Plenty of dirty dark-brown MnO2.H2O, I'll filter it and try it in the cell tomorrow........

Regards, Xenoid

[Edited on 27-8-2007 by Xenoid]

DerAlte - 28-8-2007 at 09:00

Don't bother with dirty alkaline cells if you have MnSO4 in bulk! You'll need something to mop up the SO4-- ions. I suggest sodium carbonate or NaOH:

MnSO4 + 2NaOCl + Na2CO3 --> MnO2 + 2NaCl + Na2SO4 +CO2


MnSO4 + NaOCl + 2NaOH --> MnO2 + NaCl + Na2SO4 +H2O

(one of the rare instances where NaOH acts as an oxidant - as it does in the fusion reaction)

Do this in a large vessel with good headroom with Na2CO3; it froths horribly; with NaOH of course it doesn't. The reaction is fast and exothermic, especially if you use concentrated solutions. No need to. You can use cheap hypochlorite bleach with impunity because the usual contaminant is only NaCl.

The MnO2 will be very finely divided and takes a fair time to settle as a brown hydrated form, probably MnO2.H2O. Probably need to filter with a fine pore filter paper because decanting stirs it up too much. Wash and dry (at about 150C) if you want to keep it: or use directly wet for your slurry...

I have done this process many times - the yield is close to 100%

Glad to see you haven't abandoned the method yet, in spite of the Japanese obfuscation! I am not sure how the MnO2 iselectrolytically oxidized by this method but assume all the action is very close to the anode and not in the solution (suspension).


Der Alte.

[Edited on 28-8-2007 by DerAlte]

Xenoid - 28-8-2007 at 13:04

This is the stuff I think we need! Hydrated MnO2 (MnO2.H2O), aka. Manganous Acid (H2MnO3).
I made this last night, reacting MnSO4 solution with household bleach (4%), I didn't bother to measure anything accurately.

MnSO4 + 2NaOCl +H2O --> H2MnO3 + Na2SO4 + Cl2

The material on the right-hand filterpaper was from about a third of a test tube of reagents. The material on the left was from about 250ml reacted in a beaker, the layer is about 2-3mm thick. Although it looks black, it is actually very dark brown.
I'm about to put this in the electrolysis cell, as is, without drying.

Will report back in a few hours time.... :)

Edit: Well at least something is happening this time, I've taken an aliquot after 1 hour and it is distinctly green. I assume potassium manganate is forming, hopefully this will further oxidize to permanganate!

Regards, Xenoid

[Edited on 28-8-2007 by Xenoid]

H2MnO3.jpg - 11kB

MnO2 from MnSO4

ciscosdad - 28-8-2007 at 15:51

You must have been reading my mind guys, because I was looking for patents that did just this last night. I'll post the numbers if you are interested. The one that interested me said that if you keep the pH above 8, the density of the MnO2 is higher (for use in batteries..also gives greater battery capacity). The idea is to mix the required alkali with the bleach, mix vigorously and rapidly (MnSO4 into the alkaline bleach). Settle and decant. Wash with water if required. They recommend a small excess of hypochlotite (1% to 25%).
I would guess this method will help us because the denser material should settle and filter more easily, but it may be a guide to vary the properties of the MnO2.
Interestingly, from the reading, it seems that some believe that precipitated MnO2 is superior to electrolytic MnO2 for use in batteries.

I have some questions:
How do you convert the "available chlorine" figures quoted on commercial hypochlorite to NaClO?
What is the water of crystallization in agricultural MnSO4?
Would Calcium Hypochlorite work? Would it be an improvement?

Xenoid - 28-8-2007 at 17:24

Originally posted by ciscosdad
I have some questions:
How do you convert the "available chlorine" figures quoted on commercial hypochlorite to NaClO?
What is the water of crystallization in agricultural MnSO4?
Would Calcium Hypochlorite work? Would it be an improvement?

MnSO4 can crystallise with 1, 4, 5, or 7 water molecules. It usually crystallises as the rose-pink tetrahydrate (MnSO4.4H20). The product I have, appears to be anhydrous, it is a slightly off-white powder. Makes sense, not shipping all that extra water around. When heated in a test tube it just remains as a powder with no change in appearance, a little water condenses but it is probably just adsorbed moisture. When dissolved it is a bit murky but is stabilised with a little sulphuric acid. It forms a delightfull pale rosey pink solution.

Ca hypochlorite will precipitate CaSO4 when mixed with MnSO4.

My bleach is 4.2%W/V (42g/litre) Sodium Hypochlorite
Available chlorine is stated to be 4.0%W/V. I can't remember how to calculate this and I can't find my chemistry textbook to look it up. No doubt it will involve Avogadro's number and the volume occupied by 1mole of chlorine at STP. My brain is starting to hurt!

Regards, Xenoid

[Edited on 30-8-2007 by Xenoid]

[Edited on 30-8-2007 by Xenoid]

ciscosdad - 28-8-2007 at 18:05

Thanks Xenoid,

CaSO4. Jeez how did I miss that?!?
I can only blame the accumulating obfuscation of those patents.
The acidification of the MnSO4 is a useful hint. I do intend to make some crystals for my grandson's crystal garden (got the waterglass ready..just need some crystals/lumps of various things.)

DerAlte - 28-8-2007 at 18:20


Your equation (MnSO4 + 2NaOCl +H2O --> H2MnO3 + Na2SO4 + Cl2) balances OK. Did you actually get chlorine? If you want to avoid this, see my last post.

H2MnO3 is an interesting way of looking at MnO2.H2O. Makes it appear acidic!

One could write it as MnSO4 + 2NaOCl --> MnO2 + Na2SO4 + Cl2

AFAIK, "available chlorine" merely means the amount of chlorine that can be disengaged and used compared with the theoretical famount from the active compound - i.e. essentailly the purity of the desired product.

To use Ca(ClO)2 (which can be gotten with up to 65% "available chlorine") use the method I suggested somewhere up the start of this thread to make conc. NaOCl (up to 15%) It doesn't keep well at this strength, especially unbuffered to pH> 13-14. See above.

Regards, Der Alte

guy - 28-8-2007 at 18:31

Originally posted by DerAlte

H2MnO3 is an interesting way of looking at MnO2.H2O. Makes it appear acidic!

It is actually a very weak acid.


Xenoid - 28-8-2007 at 19:59

Another aliquot taken after 3 hours and allowed to settle appears clear!!!!!
The 1 hour aliquot is a distinct, slightly bluish- green, I hope it's not a chromium compound ripped out of the SS anode.

Edit: Another aliquot, after 6 hours, it's still clear!!!! WTF !

@ DerAlte
"Your equation (MnSO4 + 2NaOCl +H2O --> H2MnO3 + Na2SO4 + Cl2) balances OK. Did you actually get chlorine?"
This reaction came from my copy of "Mellor's Modern Inorganic Chemistry" by G.D.Parkes, 1961 Ed. It is now, probably, very much out of date. It has been my chemistry "Bible" since I was a teenager. For some reason, the equation was "doubled up" ie. 2MnSO4, 4NaOCl etc. etc. not sure why, probably an editing mistake. Yes, it certainly produced chlorine, it was left out in the garden overnight. When I came out the next morning, there were two dead cats and a dead hedgehog next to the beaker! No, not really, just joking .... :D

@ Guy
The manganous acid reference was from the above reference, in the section on oxyacids of manganese. The theoretical manganous acid reacts with alkali hydroxide solutions to form manganites. It's a bit out of date, I realise, I'm not sure what the current thinking on this material is.

Electrolysis of MnSO4 will produce a dark brown oxide, which I assume is hydrated MnO2, with H2SO4 as a by-product, perhaps a neater way to go than smelly old bleach!

Regards, Xenoid

[Edited on 28-8-2007 by Xenoid]

Xenoid - 30-8-2007 at 01:53

As you may already have gathered, the optimism expressed in my previous post has proved unfounded!
After 10 or so hours of heating, stirring and electrolysis, the material in the cell remains pretty much as it was when I started, a brown slurry. The bluish-green solution, extracted after 1 hour has remained stable, how it formed or what it is I am not sure.

The only difference, I can see, between my cell and the "Japanese Patent (1st Embodiment) 1 litre experimental cell" is;

1) Scale, mine is smaller, maybe too small, and mixing or electrode spacing is an issue.
2)They use a nickel anode, but state SS is OK.
3) they added a little "catalytic" permanganate at the start, but state perchlorate is OK, I added the latter.
4) they used active or hydrated MnO2, precipitated from various Mn reactions. In my last cell I used what I assume is MnO2 precipitated from MnSO4 using NaOCl.

I have re-read an article I have on the construction of commercial permanganate cells, and I was wondering if too much mixing was occurring in the cell. Without going to the trouble of using divided cells, industrial permanganate electrolysis reduces mixing in different set-ups by;

1) Wrapping the iron rod cathode with a PVC cloth, to physically keep permanganate from being reduced at the cathode.

2) Having short stubs projecting from a sheet-like cathode, which is otherwise insulated (polystyrene) against the electrolyte. This results in a huge anode area to cathode area ratio (150-1). The high cathode current density resulting from this, apparently, mostly goes into producing hydrogen at the (desirable) expense of reducing permanganate.

Accordingly I made some minor modifications to the cell;

1) Stirring rate was reduced to about 100 rpm, just sufficient to keep the slurry in suspension.

2) The steel rod cathode was raised up, so that only about 5mm was dipping in the electrolyte.

3) Raising the cathode obviously increased the cell resistance and the operating voltage jumped up to about 3.5 volts, on top of this I increased the current a little to 2.5 amps.

After several hours of operating under this regime, not the slightest difference was observed in the electrolyte.

Getting a little frustrated at this stage, I decided to see what would happen if K-permanganate was added to the cell. I added about 10mls of quite strong solution (not transparent in a test tube). After an hour or so an aliquot was taken, the electrolyte, after settling was clear!
So, not only doesn't this cell produce K-permanganate, it actually consumes it... :(

I haven't given up on this yet, I'm going to make a batch of oxide (gunk) by electrolysis of MnSO4, and try this in the cell. If that doesn't work then its on to the............ 2nd Embodiment (fusion).

Regards, Xenoid

DerAlte - 30-8-2007 at 11:50

Keep up the good work, Xenoid! We are all rooting for you... DerAlte

Xenoid - 3-9-2007 at 00:06

Hi there everybody!

Sorry about the length of this post, but I wanted to get all my results and ideas from the last few days, in one place!

Well I have had a MODICUM OF SUCCESS with the "1st embodiment" procedure of the Japanese patent. I now understand basically how it is meant to operate. This may have been obvious to the more astute and chemically knowledgeable members of this forum, who are following this thread, but it wasn't to me!

I sat down and had another look at how my cell differed from the Japanese patent cell. One item stood out, they added a small amount of potassium permanganate, (this is referred to as a catalyst). They also mention that ferricyanate or perchlorate can be used. The oxidising agent is said to "improve current efficiency at the beginning of the electrolysis". Being a home amateur, I was not particularly concerned with "current efficiency" so I did not pay much attention to this aspect of the procedure. I didn't want to add permanganate because I though this would mask any electrolytic formation of permanganate (I have been looking to see the pink-purple colour indicating permanganate formation). Having some K-perchlorate, I just threw a pea size lump in the mix, essentially for good measure.

However, the way I now see the reaction proceeding, the addition of (preferably) permanganate, in a controlled amount at the start, is essential to the operation of the cell.
In a nutshell, the permanganate added at the beginning, "kick starts" the reaction by oxidising the hydrated MnO2, manganous acid, K-manganite, (whatever) to manganate. I'm not quite sure of the reaction, maybe DerAlte can sort it out, but it is something like the following;

2(MnO4)- + MnO2.H2O + 4(OH)- ----> 3(MnO4)-- + 3H2O

Mn(VII) is reduced and at the same time Mn(IV) is oxidised to Mn(VI)

Once manganate (MnO4)-- is present in the solution, oxidising electrolysis can commence.

The overall electrolytic oxidation process, I am reliably informed from a text book, is the following;

(MnO4)-- + H2O ---> (MnO4)- + 1/2H2 + OH-

This produces more permanganate which reacts with the MnO2 to produce manganate which is electrolytically oxidised to permanganate...etc...etc. until all the MnO2 is consumed, and eventually all the manganate is consumed. At which point we are left with a quantity of permanganate comprising that what was added, plus that formed in the cell.
Basically the electrolytic step will only work by oxidising manganate, not insoluble particles of MnO2 in suspension. This latter idea, of course, was patently absurd and should have been obvious to me! So the overall reaction is a two step process, and will not commence without the initial (chemical) oxidation.

The best "feedstock" for this operation is chemically precipitated MnO2.H2O, precipitated from the reaction of a manganous Mn(II) salt with NaOCl (bleach). KMnO4, KClO4 and KClO3 react with this to produce the initial green K2MnO4 (manganate) required for the cell to "start-up". Neither permanganate nor perchlorate react with pottery grade MnO2. Permanganate reacts with both Zn/Carbon battery "MnO2" and discharged alkaline cell "Mn gunk" to produce a green manganate solution, however I am not sure of the extent to which this may be the action of some other reducing agent in the battery mix reacting with the permanganate. Pottery grade MnCO3 reacts with KOH to produce a brownish material which in turn reacts with KMnO4 to produce manganate, so it too could be used directly as "feedstock"

I have verified this operation by putting together yet another simple cell, this time based on a 100ml beaker. The cell was equipped with a 32 cm^2 SS anode and a small SS cathode, with only a few mm. dipping in the electrolyte. The cell was filled with 20% KOH and heated to 75 oC and stirred at about 100 rpm. About half a teaspoon of dark brown MnO2.H2O was added to create a brown slurry. Next, a "tiny amount" (a little pile on the end of a wood icecream stick) of K-permanganate was added, this resulted in the beaker contents turning a very dark green (as manganate was produced, see !st. image). The green colour can be observed in the thin region (<1mm) between the back of the SS anode and the inside of the beaker wall. When the current is switched on (500mA or 16mA/cm^2), pink-purple permanganate can be seen generating in this same region (2nd. image). At the cathode, on the left, green manganate can be seen in the hydrogen bubbles. This seems to continue for several hours (8 - 12) at which point the cell starts to break down, the green colour gets weaker, permanganate production slows and then ceases, and the cell reverts to a slurry of brown "gunk".

Incidentally, in another experiment, I was trying to take a spectacular photograph of pink permanganate streaming of the anode in a green manganate solution. So of course I had to use very dilute solutions to show this, unfortunately, a dilute manganate solution will not electrolyse and remains green whilst oxygen bubbles merrily from the anode.

Interestingly, when you look at the equations above, it so happens that the products of the "chemical" process are exactly those required for the "electrolytic" process. If we combine the two equations the net result is;

2(MnO4)- + MnO2.H2O + 4(OH)- -----> 3(MnO4)- + 3/2H2 + 3(OH)-

Thus 2/3 of the permanganate produced is used in reacting with the MnO2.H2O until it is all consumed. The equation also suggests that more OH- is consumed than produced, so the cell drifts less alkaline. This could result in the cell drifting out of its narrow optimum operating range (8% < KOH < 30%). Unfortunately I have no way of monitoring the pH. I have been topping the cell up with distilled water, perhaps I should have been using KOH.

OK! Now where do I pick up my PhD!

Just as a fusion reactor consumes more power than it generates, this cell consumes more permanganate than it generates!

Regards, Xenoid

[Edited on 3-9-2007 by Xenoid]

Beakers.jpg - 13kB

DerAlte - 3-9-2007 at 12:02

@Xenoid ... Need to re-read your last post carefully to make an intelligent comment, but hey- you are getting a gleam of light among the obfuscation! Must re-read the Japanese patent carefully, too. Pinks, reds and greens - that's encouraging. If you ever see the elusive blue of hypomanganate, I might even believe the patent!

I have used a little KMnO4 in the electrolyic chlorate process, but whether it helped or not I never really determined. I must get back and do some electrolyses, you whet my appetite. I am getting my gear ready... and fusion is't out of the question (again) but I was really trying for a wet process here.

Regards, Der Alte

Xenoid - 3-9-2007 at 12:35

Originally posted by DerAlte

If you ever see the elusive blue of hypomanganate, I might even believe the patent!

Regards, Der Alte

Is this what you mean?........ Na3MnO4.......... Mn(V) :D :D

Regards, Xenoid

Hypo-.jpg - 31kB

MnO2.H2O by electrolysis of MnSO4

Xenoid - 3-9-2007 at 18:04

Best not bother with this unless you have a Pt anode.

When I first checked this out, I popped a couple of SS electrodes in a tiny beaker of MnSO4. I turned on the current and watched as dark-brown ?MnO2.H2O poured of the anode, I didn't take any notice of the current density!

I then scaled this up to several hundred mls. capacity, in a double cell of the "plantpot" type. When I turned on the current (a few hundred mA/cm^2 at the anode) all I got was O2 and H2. I had to turn the current up to about 750 mA/cm^2 to get any brown material forming. Happy with this, I then left the cell running for several hours whilst I did the shopping. When I returned, I found the SS anode had completely dissolved and the inside of the SS bowl, which was the cathode, was completely covered with a hard scaley deposit of pinkish, brownish, white manganese ?hydroxide. The liquid in the anode chamber (the plantpot) had turned a bright yellow (along with insoluble brown-crud). I assume the yellow colour was a mixture of iron sulphates and chromates and vanadates etc.

Any way I ended up throwing it all out onto the wood ashes on the compost heap, where it fizzed away, at least indicating sulphuric acid had been formed!

Regards, Xenoid

[Edited on 3-9-2007 by Xenoid]

Eclectic - 3-9-2007 at 18:10

Lead electrodes should work.

12AX7 - 3-9-2007 at 19:53

Well yeah, SS in acid conditions disappears quite readily...

DerAlte - 3-9-2007 at 21:05

@ Xenoid - the color's right for hypomangante! What the hell is that crap? How did it happen?

Regards, Der Alte (the sceptic)

Xenoid - 3-9-2007 at 23:41

Ah! DerAlte, I knew that would get you all excited!

But I am afraid I have been a little disingenuous toward you.

What you see in the image is the result of my first attempt at an MnO2 - NaOH - NaNO3 fusion.

It is Example 8. on the 2nd. last page of the Japanese patent. Where 0.5 moles of MnO2 is fused with 3 moles of NaOH and 1 mole of NaNO3. The large excess of NaOH (and NaNO3) apparently results in the formation of Mn(V)-manganate (Na3MnO4). It also results in a very fluid melt, and is carried out at 400oC.

I will go into details of my procedure in a follow up post!

Again, I had misread the patent. Believe it or not I had not even noticed the references to Mn(V). I had noticed the references to Na3MnO4, but I thought it was a typographical error as they continued to refer to it as Na-manganate (not hypo- or anything else). Thus I went about the fusion blithely unaware that I was actually supposed to be producing Na3MnO4 not Na2MnO4.
Everything went very smoothly (literally) though I noticed a bright blue melt sticking around the edges of the pot. It was a typical "copper compound" blue. I thought for a moment I had dissolved away the bottom of the SS pot, to the copper laminate which is sometimes present on the bottom of SS saucepans. However this was clearly not the case, and I made a note to ask on the SM forum what it might be. At this stage the melt was a dark forest-green "as I have heard it referred to in other threads". This seemed good, and at the end of the melt time (3 hours) I covered it and set it aside to cool.

When it had cooled, I took the photo in the previous post, in daylight, looking directly into the pot. Now, here is the disingenuous part, to my eyes the material looked very dark green. But the camera thought different, I'm not a photographic buff but maybe Woelen would know the reason this has happened. There is however blue around the edge of the pot which is real!
When this highly alkaline material is dissolved in water the solution is an inky-blue green, not at all like the beautiful emerald green K-manganate solution.
So, the material should be Na3MnO4 and should presumably be blue, but I think I have some sort of mixture.

Regards, Xenoid

DerAlte - 3-9-2007 at 23:51

Fascinating! In aqueous solution Na3MnO4 is very difficult to make and keep. But here we definitely have what looks very like it. So perhaps the Japanese patent was not all obfuscation and BS, but carefully worded to throw one off the scent, a typical patent ploy. You don't want to give any secrets away in a patent, but it can't be all BS. I haven't re-read it yet - but must do so now!

Keep up the good work, looking forward to fusion results....

Regards, Der Alte.

DerAlte - 4-9-2007 at 10:08

......continuation of my last post.....

I think some of the apparent mysteries of these industrial processes for permanganate manufacture are caused by forgetting one very important fact - all chemical reactions are, in fact, reversible.

Many we assume are not. We write the equation with the arrow ---> while we ought to write <--->; the equilibrium just exists far to the right.. As amateurs we tend to think of all reactions as going to completion - many that we try do so for various reasons at least for all practical purposes (FAPP).

An example that sticks in my mind is the production of H2 from strong acid and a metal. If you use Zinc, this reacion goes to completion, FAPP. If you Sn as the metal a pressure of 4 atmospheres will stop the reaction once the concentration of [H+] reaches 0.01. To stop zinc emitting hydrogen at this acidity needs 4*10^21atmosheres, FAPP impossible, but not in theory. (example In Pauling's General Chemistry). It all depends on the equilibrium constant for the reaction.

So what has this to do with permangantes, manganates, hypomangantes, or manganatesIVII}, (VI) or (V), if you wish?

A lot! If you heat a few crystals of KMnO4 carefully in a crucible, they decompose at c. 200-220C. Fact. Try it and listen to the crackle as they lose oxygen. The action. if reversible, lies far to the right. And the oxygen is lost, so it goes to completion, eventually. Manganate(VI) would do the same, at 190C, IIRC. and manganate(V) seems even more unstable.

Yet we keep hearing that in the fusion reaction all these manganates are being produced at temperatures much higher than these. But in this case the oxidants (air, nitrate, chlorate, and the OH radical) are in vast excess - at least when the reaction starts. Once all are used up then obviously no more managanate can be produced. If all the MnO2 is used up, similarly t the reaction stops for lack of 'fuel'.

Now if all the MnO2 has been consumed, then raising the temperature would cause decomposition of the managanate so produced, back to MnO2 which would then be oxidized by any oxidant remaining.

What we have here is a classical equilibrium reaction. In such reactions the exact conditions of temperature, pressure and phase of the moon matter fairly critically - as they do in many industrial processes, which is the reason we need Chem Engs as well as professional chemists..

.. and of course in a patent, these are the things you want to obscure as much as possible, while revealing and demonstrating the novelty of your approach.

Regards, Der Alte


ciscosdad - 4-9-2007 at 15:18

DerAlte's comments are most illuminating
The trick will be to juggle all of the disparate factors, positive and negative so that the reaction is driven to the right. The analysis figures given show that there is always some of the lower oxidation states left at the end; only a couple of percent perhaps but still there.
The process specifies a steel rod cathode and an area ratio of 100 to 200 to 1, presumably to limit the reverse reactions likely to happen in its vicinity. Perhaps we also need to take steps to limit the amount of solution cycling back into the area of the Cathode? Perhaps the stirring needs to be carefully adjusted so that the slurry is maintained, but the recirculation effect is minimised.
Some approaches come to mind:
1 .. Play with the speed of the stirrer until the correct rate is found. Nice and quiet in the middle (cathode) but turbulent around the outside (Anode). Thats where the MnO2 is oxidised.
2 .. Might need to put the rod a little off centre if there is too much of a vortex dip in the middle, but the smaller cathode area in the solution may actually help.
3 .. Put some sort of perforated baffle around the Cathode to limit fluid motion in its vicinity. Not sure if it should be conductive or not.
Perhaps its as simple as keeping most of the MnO2 near the Anode (using some combination of the above approaches).
Bear in mind that there may be deposites of some sort that accumulate on an electrode if the solution in its vicinity is moving too slowly. I recall reading another patent that specified minimum fluid velocities over the electrodes to obviate this problem.
I note that very little is said about the stirring regime in the patent.
Brilliant work Xenoid BTW.

Oops. Xenoid has already covered this. How embarrassing.
My Memory is going.

[Edited on 5-9-2007 by ciscosdad]

[Edited on 5-9-2007 by ciscosdad]

Xenoid - 4-9-2007 at 18:05

Having had enough of the "1st. embodiment", I have turned my attention to the "2nd. embodiment" This part of the patent involves the use of MnO2 in the form of pyrolusite ore (aka. pottery grade MnO2) and it's subsequent alkaline fusion to Mn(V) manganate, Na3MnO4. This reaction is carried out over 3 hours, with continuous stirring at a temperature of 400 oC.

The reaction is apparently; 2MnO2 + 6NaOH + NaNO3 ---> 2Na3MnO4 + NaNO2 + 3H2O

In practise, the proportions of NaOH and NaNO3 are increased significantly over those required by the stoichiometry of the reaction, to increase yield and the fluidity of the melt. The reagents are in the proportion 0.5 moles MnO2 + 3 moles NaOH + 1 mole NaNO3 with the NaOH in the form of a 50% solution. I'm not exactly sure why they use the NaOH in the form of a solution, perhaps it is to emulate industrial conditions where there would be solutions available from other processes. Perhaps it increases the yield and improves mixing, I'm not sure. Anybody attempting this fusion would be well advised to leave it out, because the boiling mess at the start splatters stuff everywhere, especially up the side of the vessel.

Stirring molten NaOH at 400 oC. for three hours required some specialised equipment as I was not prepared to do it by hand. In place of the no doubt expensive "SUS-27 reactor vessel" used by the Japanese boffins, I used a collection of components obtained from recycling centres and costing almost $6. These items included, an old Black & Decker drill stand, a stainless steel saucepan, a burner and regulator, an old microwave turntable motor, a 5 mm steel rod from an old printer and a few bits of scrap metal. These were assembled as shown in the image below. The steel rod stirrer was bent into an "L" shape, and the height of the stirrer assembly was adjusted so the bent part of the stirrer rod scraped nicely around the bottom of the pot. I added an offset heat shield below the motor, but in retrospect it probably wasn't needed. The whole setup worked without a hitch, the highly geared turntable motor, has a surprising amount of "grunt" for applications like this, not that much torque was really needed! Operating at about 6 rpm it was perhaps a little slow, but seemed to do the job. Temperature was monitored using a DMM with a thermocouple attached. Temperature was kept in the range 370 - 430 oC.

The fusion process was very simple, and carried out in open air. There were however no noxious fumes or evil smells evolved. After the water had boiled off, the melt just sat there being gently stirred. After cooling, the "rock hard" material was broken up with a hammer and an old chisel and placed in a sealed plastic bag, ready for the next procedure (electrolysis).

The stainless steel pot was filled with cold water and the material remaining, dissolved to a highly alkaline inky-blue green colour which seems have been stable for the last 36 hours or so!

Image 1; "Brewing up" showing apparatus and H2O boiling off, note black MnO2 and splatter.
Image 2; About halfway through the process, note the high fluidity of the melt and the obvious blue colouration.
Image 3; Final product after cooling, looking directly into the pot.

Regards, Xenoid

MangFus.jpg - 44kB

Cesium Fluoride - 4-9-2007 at 21:28


From Ullmann's:


The price of sodium permanganate is about 5 to 8 times that of KMnO4. This is mainly due to the fact that NaMnO4 cannot be made in the same way as KMnO4, because the oxidation of MnO2 in a NaOH melt does not lead to the required Na2MnO4 (with hexavalent Mn) but only to Na3MnO4 with pentavalent Mn. The latter is very unstable in dilute NaOH solution (and therefore cannot be converted electrolytically to the desired NaMnO4). Even if electrolytic oxidation were possible, there would still be the difficult problem of isolating the extremely soluble NaMnO4 from the alkaline mother liquor

Did your melt at first turn forest green (from Mn(VI))? I've tried fusing NaOH and MnO2 and I end up with a green mass which seems then to convert to MnO2 and Mn(III) when lixivated with water. The above text and the Japanese patent certainly agree with your results though.

Edit: Scratch that I now read your previous posts more carefully! At first you thought it was green...and then it turned blue.

[Edited on 4-9-2007 by Cesium Fluoride]

DerAlte - 4-9-2007 at 22:34

@Cesium: another piece of info to throw into the blender. What date is that reference?

... the oxidation of MnO2 in a NaOH melt does not lead to the required Na2MnO4 (with hexavalent Mn) but only to Na3MnO4 with pentavalent Mn. The latter is very unstable in dilute NaOH solution (and therefore cannot be converted electrolytically to the desired NaMnO4). Even if electrolytic oxidation were possible, there would still be the difficult problem of isolating the extremely soluble NaMnO4 from the alkaline mother liquor
...emphasis mine

It appears that Xenoid has achieved this. That blue color is unique - it's some time since I made hypomanganate (years!) by a process about the same as in Brauer (q,v.) No color is quite like it. Further, I know of no compound of Mn, in any other oxidation state, that is remorely blue like that.

Hydrated Cu ions do not quite have that blue blue color, and although camera images can lie, it looks like the real thing to me. Ullman apparently confirms it. I do agree Na3MnO4 is very unstable in solution. In NaOH the pH ( or at least the concetration) is very critical - but that's exactly what the Japanese chemists said, is it not?

If they don't make the NaMnO4 by electrolysis, then how do they make it? The solution is an industrial chemical used in bulk conc. solution as a substitute for the K salt.

Looking at the Redox potentials, it's a lot harder to oxidize MnO2 to MnO4 - - -
(0.96V) than to MnO2 - - (0.62V) or to MnO4 - (0.60V) ( all in standard conditions, of course [OH-] =1N, assumedly, at STP). Yet it requires a mere 0.27 volts to convert MnO4 - - - to MnO4 - -, much less than to convert MnO4 - - to MnO4 -

Curiouser and curiouser, said Alice...

Regards, Der Alte

not_important - 4-9-2007 at 23:37

To really judge colours in image you need some reference samples to compare; that's why film photographers would take a picture of the Kodak test target or similar at the start or end of a roll of film. In this case having a bit of copper sulfate in the image, or some blue and green samples that are common around the world (think Pepsi logo and the like) are useful.

There are other possibilities for that colour, for instance I've made ceramic glazes in that colour range using nickel and an "alkaline" glaze with a fair amount of Li2CO3 in it. Obviously if some of that lovely blue stuff in the photo were diluted a bit, and slightly acidified managanese would drop out but any nickel would go back into solution and be detectable. Using acetic acid or ammonium sulfate to acidify should work, if no nickel shows then the other possibilities are chromium and Mn(V), and I don't think chromium gives that sort of colour under those conditions.

[Edited on 5-9-2007 by not_important]

Xenoid - 4-9-2007 at 23:44

I gather the main thrust of the "Japanese patent" is all about reducing the number of steps involved in permanganate production ie. 2 steps to 1 or 3 steps to 2 etc. Combined with this is the ability to handle various waste or by products of permanganate oxidation reactions and at the same time simplify the handling of caustic alkali solutions.

The Mn(V) manganate I have apparently produced is next hydrolised in water according to the following reaction;

2Na3MnO4 + 2H2O ---> Na2MnO4 + MnO2 + 4NaOH

Also part of the Mn(VI) undergoes the following reaction;

3Na2MnO4 + 2H2O ---> 2NaMnO4 + MnO2 + 4NaOH

This "slurry" is essentially the same as the mixture in the "first embodiment" and the cell I have been trying to get to work! It is of course though the sodium equivalent, they seem to use K and Na interchangeably. I have used NaOH because I have plenty of it in solid form, but no solid KOH.

I assume K-permanganate can be made from Na-permanganate by reacting with KCl but a perusal of other threads suggests people have had problems with this.

I'm now building, yet another small cell (about 400 - 500mls) using a squat bottling jar with a 10mm thick perspex sealed lid. Hopefully it can be run without topping up every half hour.

Incidently, the beaker of beautiful emerald-green K-manganate which I left siiting on the bench after carefully filtering all the gunk out, completely reverted to brown gunk overnight. :(

Oh well, I'm off for a days skiing tomorrow, see you all later! :D

@not_important. I just noticed your post. I had exactly the same idea about using some standard colours, but it was too late for the images above. The two images of the stuff in the pot, the large one earlier and the small one in my last post are two seperate images, one with flash the other without, not sure which is which. I guess the slightly brighter one was flash. I also thought about Cr and Ni from the SS pot. The inside has turned a brownish coppery colour (from oxidation?) but it is not pitted or corroded.

Regards, Xenoid

[Edited on 4-9-2007 by Xenoid]

DerAlte - 5-9-2007 at 12:42

Not_important said

To really judge colours in image you need some reference samples to compare; that's why film photographers would take a picture of the Kodak test target or similar at the start or end of a roll of film. In this case having a bit of copper sulfate in the image

The closest color to how I remember Mn(V)O4 - - - appears is azurite, basic copper carbonate. But to me that is a shade too green. I seem to have a different visual threshold between blue & green to others.

It’s even more difficult to describe in words. The best I can do is call it the color of the sky looking south (not north) on a relatively cloudless day. It’s a light blue, a very unsaturated hue. A rare coloration in a pure compound.
Nickel is nearly always a nice green, apple to deep in color. Chromium does give hydrated ions in the purple range and perchromic in the blue, but those blues are much deeper (more saturated), IIRC.

I’m convinced Xenoid has Mn(V). And I’m a skeptic of the worst kind!

But one has to be a little careful with the manganate family. If you take a dilute alkaline (c. 1N in OH-) solution of permanganate and slowly reduce it with sulphite drip by drip, it turns murky (MnO2) green (MnO4 - -). If you do it slowly enough, you have both manganate and permanganate ions present. This can produce a blue color which isnot hypomanganate, but due to the combined filtering effect of the green passband of the manganate and the red passband of the permanganate. Once all is converted to manganate, a clear green is seen (if you let the MnO2 settle). Then, at a low enough temperature, you can effect the conversion MnO4 - - to MnO4 - - - and get the sky-blue but transient color – if you are lucky and the moon is full, etc. . It is, to me, amazing to see the stuff lying at the bottom of a beaker. Of course, aqueous solutions and fused conditions cannot be compared except remotely.

Xenoid points out a very important fact – when you chemically oxidize a lower manganate to a higher, you always get a loss of Mn as MnO2. Unless you do it electrolytically, by removing one of the excess electrons. Electrolysis is a powerful and cheap way to oxidize (or reduce) for the amateur, but conditions are often tricky.

Great work, Xenoid. We’ll let you have a few days off. Lucky devil, skiing while we bake in the sub-tropical lowlands here (jealous!). Come back refreshed and rise to new heights of achievement.

Regards, Der Alte.

Cesium Fluoride - 6-9-2007 at 09:44

I just love this thread! Soon....oh soon I hope...I'll be joining the quest!

@Cesium: another piece of info to throw into the blender. What date is that reference?

The encyclopedia was published in 2002. Unfortunately, there's no reference given for that piece of information.

There is an entry on potassium manganate(V), but not the sodium salt.

Potassium manganate(V) [12142-41-5] , K3MnO4, Mr 236.24, r 2.78 g/cm3, decomp. > 1000 °C, occurs as fine blue-green to turquoise crystals. In the presence of water it is readily hydrolyzed, disproportionating to K2MnO4, KMnO4 , and MnO2. Solutions of K3MnO4 in 40 % potassium hydroxide have limited stability even below – 10 °C; however, in the presence of 75 % KOH and under nitrogen, potassium manganate(V) is stable up to 240 °C. Pure potassium manganate(V) can be heated to over 900 °C without decomposition. It is an important intermediate in the production of potassium permanganate.

I also assumed that one could make sodium permanganate by fusing NaOH and MnO2, but these experiments have shown us why this is not possible. I still to do not understand, chemically, why there is a difference between the oxidizing powers of fused sodium hydroxide and potassium hydroxide. Does anyone have a guess?

Furthermore, before discovering the information in Ullmann's encyclopedia, I had previously gathered several (albeit unreliable and old) sources that suggested that one can make NaMnO4 directly by fusing NaOH and MnO2. Here they are:

Sodium manganate (Na2MnO4), prepared by fusion of a mixture of natural manganese dioxide and sodium hydroxide; green crystals, soluble in cold water, decomposed by hot water.

Sodium permanganate, NaMnO4, is obtained in a similar way to the potassium salt, and is distinguished from it by being deliquescent, and therefore, crystallizing with difficulty.

Sodium manganate, Na2MnO4, is formed when a mixture of equal parts of manganese dioxide and soda-saltpetre is heated for sixteen hours; the mass is then lixiviate with a small quantity of water and the solution cooled down, when the salt separates out in small crystals isomorphous with Glauber’s salt, and having the composition Na2MnO4 + 10H2O. These dissolve in water with partial decomposition, yielding a green solution.

For disinfecting purposes it is not necessary to employ the pure, well-crystallized salt [potassium permanangate], which is used in the laboratory, but a commercial article consisting of a mixture, more or less pure, of manganate and permanganate of sodium is used. The substance is obtained by mixing the caustic soda obtained from 1,500 kilos of soda-ash with 350 kilos of finely-divided manganese dioxide in a flat vessel, and heating this mixture for forty-eight hours to dull redness. The product is then lixiviated with water, and the solution either boiled to the requisite degree of strength or evaporated to dryness. If the manganate is to be completely converted into permanganate it is neutralized with sulfuric acid, the solution concentrated until Glauber’s salt separates out, and these crystals are then removed and the liquid further evaporated.[2]

[2] Hofmann's Report Exhib. 1862, p. 109.

So you can see why I thought NaOH and MnO2 would work! Am I misinterpreting the information given in these sources?

The last 2 sources I gave are particularly enticing. It is interesting that both of them use "soda" or "soda-ash", presumably sodium carbonate. Perhaps, if I heat a mixture of NaOH and MnO2 for "48 hours to dull redness"? Red heat is what? 500C? Maybe in this process sodium carbonate decomposes to sodium oxide, but that doesn't happen until 800C and the oxide doesn't melt until 1100C.

What am I missing here?

[Edited on 6-9-2007 by Cesium Fluoride]

DerAlte - 6-9-2007 at 12:43

@ Cesium
I just love this thread! Soon....oh soon I hope...I'll be joining the quest!

Hell, Cesium, you’re already part of the team. Your last post was most illuminating.

I am, I feel, beginning to see the light in this permanganate business. Xenoid’s blue product amazed me, and now you quote a reliable source saying
…Potassium manganate(V) [12142-41-5] , K3MnO4, Mr 236.24, r 2.78 g/cm3, decomp. > 1000 °C,…. . In the presence of water it is readily hydrolyzed, disproportionating to K2MnO4, KMnO4 , and MnO2.

The second part of this statement is my experience, and the first Xenoid’s (true, not at 1000C, but at elevated temperature). Further, overnight in water it deteriorated to MnO2. All the comments I made and Xenoid made are entirely valid and compatible with the above…

I’ll restate what I said in an earlier post: Of course, aqueous solutions and fused conditions cannot be compared except remotely. I was doing just that previously…

Manganese has an incredible affinity for oxygen. MnO can be made but only out of the presence of oxidizing agents. Greenish solid. It oxidizes readily to Mn2O3 (brownish) or even MnO2 (black or brown, depending on crystalline nature, hydration, etc). MnO2 is a good oxidizer and stable to c. 500C. It seems to have very strong covalent links to its oxygen.

The oxyanionic MnO4 moiety can exist as Mn(V), Mn(VI) or Mn(VII) with the appropriate charge..

The Mn(IV) state only exists stably with oxygen and fluorine. The Mn(III) state (manganic) is strongly oxidizing, but Mn2O3 and MnO(OH) are quite stable. Mn(II)(OH)2 (white) rapidly oxidizes to MnO(OH) (brown) in air… and most of the manganous salts also easily oxidize to MnO(OH). Mn thus likes oxygen in all ratios.

Permanganate has probably 3 covalently bonded oxygens and one having a single bond plus a spare electron. Less stable in aqueous solution is the Mn(VI) ion, two covalent double oxygen bonds and two singly bonded oxygens plus 2 electrons; and even less stable is the Mn(V) variety.

Then you have the manganites, compounds of Mn(IV) oxidation state, insoluble, looking very like MnO2 (black), stable, but capable of evolving chlorine with conc. HCl. For example K2MnO3, i.e. which might be written as K2O.MnO2 as the Victorian chemists often did.

On heating alone, anhydrous Mn(VI) and Mn(VII) manganates decompose in the 190-220C region. Not so the Mn(V) manganate, we learn. The industrial processes use long heating at up to low red heat (I’d call that 500C) and use atmospheric oxygen to oxidize. The hydroxide has to be in excess. Only under these conditions can the manganate be stable – any tendency to revert to MnO2 will be balanced by reconversion to the manganate. I.e. we have a classic equilibrium case – the exact temperature is critical for maximum yield. And that temperature, the amount of hydroxide, etc., are closely guarded trade secrets.

It is no wonder, then, that amateur efforts to produce manganate by the fusion method founder. Not only is the heating time required rather too long, but temperature conditions are probably far from optimum. Good results could probably be attained in a temperature controlled electric furnace with added oxidant, such as nitrate, to speed it up a bit.

The fusion processes seem to produce the Mn(VI) or, as we have learned from Cesium’s references and Xenoid’s fine work, Mn(V). The latter is a new twist. However, both these lower manganates easily hydrolyze to the permanganate, or can be oxidized electrolytically, or with Cl2, even CO2, or any acid that permanganate do not oxidize.

We lose MnO2 in the process, if done chemically. But that can be recycled. There’s a sound method somewhere in all this madness…

Regards, DerAlte

not_important - 6-9-2007 at 19:07

Incipient red heat is in the low 500s C, dull red heat is 650 to 700 C, cherry red is around 850 C.

The higher oxides of manganese are 'acidic' and will react will molten Na2CO3 to release CO2 and form salts.

DerAlte - 6-9-2007 at 20:59

Thanks for that info., not_important. I'd always put it a bit lower, but I guess it's a bit subjective and depends heavily on light level and maybe even on eye response. We'll take your figures as representative of typical in texts and references. Oh for a decent pyrometer!

Regards, Der Alte.

not_important - 6-9-2007 at 21:51

The numbers I gave are sort of an average, some metallurgical references run higher once you're out of the dark/dull/low red heat; a few ceramics refs run a bit cooler, possibly because those are almost always in a kiln and the enclosed space reduces heat loss from the surfaces - hot stuff in the open will have a cooler skin than bulk temperature.

chief - 8-9-2007 at 07:18

Has anyone any source for the solubilities of
- KMnO4
- Ba(MnO4)2
- Ba(NO3)2
in dependency from the temperature ? I can't find anything on the net except values without the temperature or at 25 [Celsius] or something. But I go for Ba(MnO4)2 from Ba(NO3)2 and KMnO4,
and need those data for optimum performance.

chief - 8-9-2007 at 07:44

Besides: How sounds 15 $/kg for KMnO4 ? 10 kg at 10$/kg ??

DerAlte - 8-9-2007 at 09:25

@ chief - the price is good, but who needs more than a kg, except for an indusrial user?? Watch what the laws say about possession in your state. The Fed limit, DEA, is Ikg/yr, IIRC. State laws may be more restrictive.

For the solubilities see, the table in Wiki. They are a usually reliable source.

The figure for barium permanagnate is incorrect, however. It is quite soluble. I think the compiler mistook the manganate figure for that of permanaganate. The mangante is almost insoluble, about the only manganate that is stable. Barium permanganate is in the region of 30-60 g/100g aq. at RT, AFAIK, but I'm rather unsure about that. I haven't been able to find it either.


Der Alte

[Edited on 9-9-2007 by DerAlte]

chief - 8-9-2007 at 13:03

No such limits for me (Germany)(not that I knew); just forbidden to export it to Honduras and e few other exotic countries, because they make drugs there (oxidizing organic materials). (Sorry for you americans (?) to have such stoneage-politics)

The wiki-site has only one solubility for KMnO4, at 20 [Celsius], but it dramatically increases with temperature, and that would be interesting. (
In the same manner i need the _curves_ for the other 2 salts, to know how to drive the temperature; wiki is not much help there.

Mainly until now I mix concentrated solutions of Ba(NO3)2 and KMnO4, watch the temperature and try to find the best point of harvesting the Ba(MnO4)2, but always I get some (50+%) amount of useless KNO3/KMnO4/BaNO3/Ba(MnO4)2 - mixture, which I need to separate ...

Cesium Fluoride - 8-9-2007 at 15:34

Mr. Cheif,

Here is some information I dug up for you.

Solubility of barium nitrate in 100g H2O:
0C: 5.0g
10C: 7.0g
20C: 9.2g
30C: 11.6g
40C: 14.2g
50C: 17.1g
60C: 20.3g
80C: 27.0g
100C: 34.2g

Solubility of potassium permangate in 100g H2O:
0C: 2.83g
10C: 4.4g
20C: 6.4g
30C: 9.0g
40C: 12.56g
50C: 16.89g
60C: 22.2g

Probably at higher temperatures, it decomposes too readily.

No info about BaMnO4 except that it's soluble. Sorry it's rather exotic.

Xenoid - 8-9-2007 at 18:32

Well, I'm back on the job!

The new permanganate electrolytic cell is completed and is up and running. Based on the figures for the 1 litre cell in Example 8. of the Japanese Patent (20Amps for 4.5 hours at 400A/m^2) my 450 ml cell should have a run time of about 13 hours. Although less than half the size, I am operating it through my constant current lab power supply which has an upper limit of about 3.5 Amps. The patent cell operates for 90Ah so for my cell I need about 45 Ah, so at 3.5A that is about 13 hours. My cell is operating at about 25mA/cm^2 based on an anode size of 142 cm^2.

I am using 100g (the yield was 216g) of the fused material, which turned an intense green in the cell when dissolved. It is so intense that only in the very thinnest part of the meniscus, was the green visible. Interestingly, after operating the cell for only about 30 mins. the meniscus had already turned purple. Probably due to evolved oxygen in the headspace.

I just went down to check the cell, which had been very slowly heating up to 70 oC. which was the maximum temperature I was going to operate at, only to find that it had cracked. I think perhaps the tight fitting perspex lid had expanded more than the glass. A couple of bits of glass have broken from the neck, fortunately above the electrolyte line, but large cracks extend down the side and intense green manganate solution has run onto the hot plate etc. I'm very surprised and disappointed by this, these type of bottling jars will normally withstand over 100 oC.

This will set me back a bit!

In the mean time I would appreciate any suggestions on how to convert Na-permanganate to K-permanganate. I'm thinking of adding hot saturated alkaline KCl solution in the right stoichiometric proportions and cooling to 0 oC or less.

EDIT: What a f**king mess! Concentrated alkaline manganate/permanganate solutions are incredibly messy. Browns, greens, blues and purples and thats just my hands! Bad choice of materials all round, the perspex (acrylic, PMMA, plexiglas) lid, features radial cracks around the outside after only about 1 hour operation. The "bottling jar" also had another crack, extending almost half way around the jar, just above the base. The reason I used glass and perspex was because I wanted to see the activity in the cell. Given the fact that the solutions are so dark and it is impossible to see anything going on, I might revert to my original idea of using SS container, and fit a PVC or PE lid.

Image 1: Cell components, showing jar, lid, electrode assembly and vent tube. The height of the cathode can be adjusted but will be immersed only about 10 mm in the electrolyte.
Image 2: Cell in operation, note hydrogen vent tube (PVC) which also serves to condense any water vapour being evaporated from the cell, and return it.

Regards, Xenoid

[Edited on 8-9-2007 by Xenoid]

Permanganate.jpg - 28kB

DerAlte - 8-9-2007 at 21:39

@ Xenoid - welcome back and hope your sojourn in the mountains has left you refreshed.

If I ever had a fleeting doubt left about your lovely blue hypomanganate (I did not!) you have now dispelled it. The fact that your got the intense dark green manganate proves it. Instant hydolysis.

WRT the breakage of jars, differential temperature shock does it. I'd use a water bath for the process. Pyrex Beakers are much more reliable but if one doesn't have one one has to improvise.

Your set up looks positively professional. I believe you are on the verge of success.

In the mean time I would appreciate any suggestions on how to convert Na-permanganate to K-permanganate. I'm thinking of adding hot saturated alkaline KCl solution in the right stoichiometric proportions and cooling to 0 oC or less.

That should work fine. The solubility of the Na salt is very large ~300g/100g aq at RT; that of the K salt quite low ~2.8g/100g at 0C - an easy separation.

I always use gloves dealing with manganates or Mn salts. They stain everything, clothes, hands, benchtops etc. They deposit MnO2 or MnO(OH) - a nasty brown - use a hard scrub or peroxide on skin, conc. HCL on glass or porcelain, to remove.

@chief - Cesium has given you the figures. You want to search under Solubility Table in the search engine to find the comprehensive table, not under each compound. Try


Der Alte

Cesium Fluoride - 8-9-2007 at 21:50

Keep up the good work Xenoid!

@chief - Cesium has given you the figures. You want to search under Solubility Table in the search engine to find the comprehensive table, not under each compound. Try

I got my figures from Seidell's Solubility of Inorganic and Organic Compounds (1919) so there may be some deviation. That's a useful link I didn't know of, but as you said it's solubility for Ba(MnO4)2 is definitely wrong.

[Edited on 8-9-2007 by Cesium Fluoride]

ciscosdad - 9-9-2007 at 15:46

I was looking over some patents on the weekend, and found that there are a couple that appear to be further developments of the Japanese patent.
The idea is to replenish spent permanganate etch solutions.
I believe the numbers are 4835095 and 4911802. Not entiirely sure (my copies are at home). If you are interested look at the last few entries on the list I posted (page 5 this thread). The main difference in the new development is a diaphragm around the cathode containing concentrated alkali. The main solutions are about 1 molar alkali, and the current efficiencies look a bit poor at first glance.
Have you tried the route through Aluminum Permanganate for your Barium Permanganate synthesis? If I recall correctly, the idea is to add sufficient aluminum sulfate to the permanganate solution so that you can crystallise potash alum from the mix. The resultant conc Aluminum Permanganate solution can be treated with an oxide, hydroxide or carbonate of the required metal to give the permanganate. The aluminum precipitates as the hydroxide.
I can find a relevant patent if you are interested, but we have discovered that that may not be as much help as you think.:(

S.C. Wack - 11-9-2007 at 15:59

Originally posted by Cesium Fluoride

From Ullmann's:


The price of sodium permanganate is about 5 to 8 times that of KMnO4. This is mainly due to the fact that NaMnO4 cannot be made in the same way as KMnO4, because the oxidation of MnO2 in a NaOH melt does not lead to the required Na2MnO4 (with hexavalent Mn) but only to Na3MnO4 with pentavalent Mn. The latter is very unstable in dilute NaOH solution (and therefore cannot be converted electrolytically to the desired NaMnO4). Even if electrolytic oxidation were possible, there would still be the difficult problem of isolating the extremely soluble NaMnO4 from the alkaline mother liquor

A simple quick experiment shows the truth. Whatever it is that is precipitated from an impure Mn salt by addition of hydroxide was heated with NaOH and enough water made into a paste, which was spread in a thin layer on the bottom of a steel container. Heating produced the manganate within minutes and this was leached out with water. After standing for a time the Fe impurities precipitated as Fe2O3 or something else, and the dark blue green solution with little crystals of sodium carbonate glistening in the light was observed to be perfectly red by sunlight reflecting off of the bottom of the glass through a small amount of liquid. Dilution with water produced a swiftly changing number of colors ending at permanganate. Heating the filtered original solution some soon produced a somewhat concentrated solution of permanganate.

Isolation is perhaps unnecessary if you're planning on using it in solution anyways, and if some base is OK, and it's concentration can be determined by any number of analytical methods.

ciscosdad - 11-9-2007 at 16:43

@ S C Wack
Fascinating but confusing.
Do you think it is concentrated enough to get a precipitate of KMnO4 with (say ) KCl or aybe KOH? Would it survive boiling long enough to concentrate?
Could it work with KOH?

DerAlte - 11-9-2007 at 17:38

@ S.C. Wack - the Victorian chemists nearly all say that a 'sensitive' test for Mn compounds is to heat on a Pt foil with NaOH (or KOH) with (or without - i.e. using oxygen in air) a oxidant. They say that permanganate ( sometimes manganate) is produced, just like in your experiment.

Why, then, is the fusion reaction nearly always disappointing?

When Xenoid made the hypomanagate, I was really surprised. I had been lulled into thinking it was necessarily unstable because of it's easy hydrolysis in all but solutions of pH~14. Yet the claim is that as a solid it's stable to >1000C. (Not sure I believe that!).

Permanganate never seems to be produced in fusion oxidations. We've seen hypomanganate and manganate. Both of these easily pass into solution to permanganate. But there's a price to pay - at each oxidation step from MnO4 - - - to MnO4 - - or MnO2 - - to MnO4 - , we lose some MnO2.

3MnO4 - - - + 4H+ - -> 2MnO4 - - + MnO2 + 2H2O

3MnO4 - - + 4H+ - -> 2MnO4 - + MnO2 + 2H2O,

so that 9 mols MnO4 - - - are needed to give 4 mols MnO4 - and 5 mols MnO2. Waste of oxidant! Or MnO2, but that could be recycled.

As for any reaction working with NaOH and not with KOH, I cannot believe that. True, the solubility of Li, Na, and {K, Rb, Cs] often differ - carbonates, for example - Li is sparingly soluble, Na soluble, and {K,Rb,cs} very soluble.

I am not familiar with Ullmann. The chemistry of Mn is often somewhat perplexing. That's what makes it so interesting. AFAIK Mn produces no complexes only oxyanions (except organic, and the halogen analogs of MnO4, if you call them complexes) yet all the other transition elements positively bristle with them.

Consider also manganous permangante, which should be a stable compound, Mn(MnO4)2 = Mn3O8 - another oxide of Mn??? Reminds me of mellitic anhydride, an 'odd' oxide of carbon, or carbon suboxide, C3O2 (if you call it 1,2 propene, 1,3 dione per IUPAC then you get an idea of what it does; O=C=C=C=O).

Regards, Der Alte

12AX7 - 11-9-2007 at 17:45

As I've mentioned somewhere before (if not in this thread, then elsewhere lost to the sands), I've fused MnO2 + NaOH, obtaining a dark, deep green melt stable up to red heat (~600C).

Chromate is stable up to the same temperatures; is it so suprising that lower states of Mn are? Mn(7) is pretty high and not a good representative for this decision, remember.


Cesium Fluoride - 11-9-2007 at 20:09

Yes, I've produced that same green melt, but in solution I never could get adequate yields of permanganate. I now think that actually the result of the fusion of MnO2 + NaOH is actually mostly MnO4-3 with enough MnO4-2 to color the mass green. That's why I was extremely interested in Xenoid's hypomanganate experiment because he at first said that the mass turned a forest green color and then only later did it turn blue.

I suppose that the dilute solution of permanganate that I was able to procure was a result of the hydrolysis of hypomanganate to manganate and finally to permanganate, with, of course, lots of MnO2 crud being precipitated.

I have a hunch that Ullmann's is closest to the truth as its the only solid explanation I've seen given for the prevalence of KMnO4 over NaMnO4.

DerAlte - 11-9-2007 at 20:27

@ Cesium Fluoride

You said

I have a hunch that Ullmann's is closest to the truth as its the only solid explanation I've seen given for the prevalence of KMnO4 over NaMnO4.

Not true, friend! NaMnO4 is trucked around the country as a 40% solution for industrial use in water treatment plants. Goggle it. I have no reason to suppose the Na salt is any more difficult to produce than the K salt. As a lab reagent the sodium salt is far too soluble and deliquescent to be easiy used. Hence the prevalence of the K salt in labs - it has no water of crystallization, is not deliquescent, and it reasonably easily obtain pure to ACS standards - or to purify oneself.

Regards, Der Alte

Cesium Fluoride - 11-9-2007 at 20:42

Not true, friend! NaMnO4 is trucked around the country as a 40% solution for industrial use in water treatment plants. Goggle it. I have no reason to suppose the Na salt is any more difficult to produce than the K salt. As a lab reagent the sodium salt is far too soluble and deliquescent to be easiy used. Hence the prevalence of the K salt in labs - it has no water of crystallization, is not deliquescent, and it reasonably easily obtain pure to ACS standards - or to purify oneself

Of course, DerAlte. I'm all too familiar with this and have googled/"libraried" NaMnO4 scores of times.

Yes I realize the deliquescent nature of NaMnO4 and that it is hard to crystallize but I cannot find any legitimate reference to the production of NaMnO4 by fusion of NaOH. (If someone can, please show me!) Hence I think the K salt is mainly more common because it can be made more easily!

Once I produce some KOH, perhaps I'll be able to back up my claims!

[Edited on 11-9-2007 by Cesium Fluoride]

Only partial success, again!

Xenoid - 11-9-2007 at 21:08

Sorry, but I've had a SMART failure on my hard drive, so I will have to get a new one!
I'm typing this on my son's laptop, so can't be bothered putting any pictures together, they will have to wait until I get my own computer running again!

Just a brief report!

Well, I found a nice SS canister for a new cell, made a new lid (used polystyrene this time, another poor choice!). Used solution recovered from previous attempt. Ran for 14 hours at 75 oC and 3.5Amps. Everything went smoothly. At the end of the run I lifted the lid and was greeted with intense purple. I thought this looked pretty good so I added the stoichiometric amount of KCl as solid to the hot solution and continued the heat and stirring for another 10 mins. or so. Then allowed to slowly cool to room temperature. I then transferred the container to the refrigerator and slowly cooled to -5 oC. After a few hours, I decanted the liquid, expecting to see lots of chunky K-permanganate xtals. The yield should have been about 39 g. Unfortunately, although the bulk of the solution was intense purple, there was a lower layer of intense green and still brown crud on the bottom. There were K-permanganate crystals in the brown sludge, and sticking to the sides but only a few grams in total!....:(

I now have 500mls of totally useless gunk! It has Cl- in it, so I can't electrolise it any more.

I am going to give this one more shot. I will use the second half of my fusion mix. I think perhaps I need a longer electrolysis time and more vigorous stirring. I was only stirring at 100 rpm which in retrospect wasn't enough!

Regards, Xenoid

12AX7 - 11-9-2007 at 21:16

Colors are great, but they suck when you're actually working on the thing that's making the color. I can add KClO3 to hot Cr2O3 all day, and the melt will turn yellow and stuff, but only when it's pure blood red is it actually nearly converted.

Moral is, it's probably better to sample and titrate than to trust the color.


[Edited on 9-12-2007 by 12AX7]

Cesium Fluoride - 11-9-2007 at 21:20

Yes, I have even seen the greens of manganate form when boiling a NaOH solution with MnO2. Of course, this color signifies almost nothing from a practical standpoint.

DerAlte - 11-9-2007 at 21:56

@ 12AX7 & cesium - second that. Titrate. Then you'll know how much MnO4- you have and how much KCl you need. Permangante solutions can look black and still be dilute.
Hells bells, the forum is busy tonight!

Der Alte

Xenoid - 11-9-2007 at 22:38

Yeah! The solution is effectively black!

When I talk about intense colours, I'm referring to thin films adhering to the thermometer and the inside of the container above the liquid level. I'm not sure what else I can do to get this to work, I'm running out of ideas!

Not sure about you guys, but there's nothing on TV here tonight!

Regards, Xenoid

DerAlte - 12-9-2007 at 10:43

@ Xenoid ….. a couple of tips, although you probaly know them.

Use colorimetric titration – a true volumetric analysis takes a lot of time weighing, making up standard solutions, etc., even if you have decent volumetric equipment. You do need reasonably pure KMnO4 to start. Make a solution of about 0.01 M – I can’t remember the figure but it’s something like that – to give a clear transparent red permanganate reference solution. You need to be able to weigh accurately for this.

Using an identical test tube against a white background, fill the tube to about ½ the same level as the reference with water acidified to about 0.05 – 0.1 M – virtually any acid not attacked by KMnO4 can be used, even HCl. Run in the test solution drip by drip from a micro –burette, improvised or otherwise, until the color approaches the reference. Judiciously add acidified water and solution to be tested until there is a good color match and equal volumes. It sounds more difficult than it is. The amount of MnO4- ion present is then equal in both tubes and a quick calculation gives you what you have in the solution to be tested. If the solution is highly alkaline it must be first brought to near neutrality to convert manganate to permanganate.

The estimate will them include both Mn(vii) and Mn(v1) manganates (and even the sweet sky-blue hypo!) per the reactions I wrote a few posts back.

A quick test for the presence of Mn(vi) and Mn(vii) is to use crude chromatography. Drip a drop on to a filter paper. The permanganate diffuses faster than the manganate. Two separate rings form, one green, one red. The size of these rings is a rough indication of the presence and amount of each. The paper will go brown in a minute or so due to oxidation, but there’s enough time to see distinctly.

Regards, Der Alte

Xenoid - 12-9-2007 at 23:15

Well I'm up and running again, thankfully no data was lost!

Just a few images from my last attempt, outlined previously.

1: New SS cell and electrode assembly, much simpler arrangement as the SS canister is the anode. Lid is made of a strange highly compressed polystyrene board, it can be worked like wood but it deformed a little at 75 oC. and did not stand up to the oxidising conditions very well, not recommended for this type of application!
2: Cell up and running, operating at 2.9 Volts and 3.5 Amps.
3: A clump of K-permanganate crystals, more than this formed, these were just residual after transferring the solution a few times.

I'm now making a PVC lid for my final attempt, I'm going to use a 5V computer SMPS and run about 10 Amps for 24 hours. Depending on the results, I will also try concentrating the solution.

Regards, Xenoid

Permanganate3.jpg - 26kB

Antwain - 15-9-2007 at 08:23

Sorry to post a half-informed post, but I just read the last couple of pages, and haven't looked at the patent... but it occurs to me that the reaction at the anode may be reducing the permanganate? I could just be plain wrong about that, but has anyone tried separating the anode and cathode with a salt-bridge?

Also, whilst I don't trust wiki as far as it should be able to be thrown, the figure of 0.015g/100g @20*C for the solubility of Ba(MnO4)2 may in fact be correct. The Merck index, which I do trust, says it is "sparingly soluble in water" which in my experience translates to "as near as counts is insoluble, but does dissolve SLIGHTLY" I will do this test... I will add a small amount of saturated solutions of Ba(NO3)2 and KMnO4 then add say 5 times the water to it. *If* a precipitate forms, I will collect it and treat it with some dilute sulfuric acid (leaving it for a while so that equilibrium can be established) and see if in fact BaSO4 is produced. This will at least settle whether it is almost not soluble or at least almost as soluble as KMnO4. I will get back to you all within a week, since I may not have time to do it tomorrow.

Ok, perhaps you are right and the wiki is (again) wrong. I have just found a synthesis of the acid following the equations:

KMnO4 + AgNO3 -> KNO3(aq) + AgMnO4(s) red precipitate : wiki claims 0.9g/100g water @ 20*C, but don't trust it ;)

2AgMnO4 + BaCl2 -> Ba(MnO4)2(aq) + AgCl(s)

Ba(MnO4)2 + H2SO4 -> BaSO4(s) + 2HMnO4(aq)

edit- second paragraph
edit- third para
[Edited on 16-9-2007 by Antwain]

[Edited on 16-9-2007 by Antwain]

Xenoid - 15-9-2007 at 12:09

Originally posted by Antwain

Sorry to post a half-informed post, but I just read the last couple of pages, and haven't looked at the patent... but it occurs to me that the reaction at the anode may be reducing the permanganate? I could just be plain wrong about that, but has anyone tried separating the anode and cathode with a salt-bridge?

Yeah! You are just plain wrong!... :(

If you HAD read the previous posts and the Japanese patent, you would realise that reduction at the CATHODE is apparently not a big issue. The japanese boffins don't even seem to bother about it in their lab scale experiments. Industrial electrolytic oxidizers use short stubby cathodes with very high current/surface area, this inhibits reduction. In the older style electrolytic cells, porous material was wrapped around the cathode! In my first, ridiculous, half-hearted attempt at oxidising alkaline MnO2 I used a separate anode/cathode chamber (plantpot), (see my first post up the thread).

In my later cells I have used a short stubby cathode (the iron rod) it only dips about 10 mm into the electrolyte.

I'm not saying reduction doesn't occur, but we are trying to emulate the Japanese patent, and they don't even mention it (from memory). I suppose, that on a very small scale (< 1 litre) it could become important, but at this stage, unless someone can come up with some convincing evidence, I'm not prepared to build a double cell, what with all that near boiling NaOH, etc....

Regards, Xenoid

Eclectic - 15-9-2007 at 12:45

It was my understanding from reading a book on industrial electrochemistry (from the 1950's, and speaking from memory as I don't have it with me) that commercial permanganate production involved making sodium permanganate by electrolysis of a sodium manganate solution, and the producing the potassium salt by metathesis with KCl, all due to solubility issues.

I'll recheck the book when I find it.

Xenoid - 15-9-2007 at 12:57

Originally posted by Eclectic
It was my understanding from reading a book on industrial electrochemistry (from the 1950's, and speaking from memory as I don't have it with me) that commercial permanganate production involved making sodium permanganate by electrolysis of a sodium manganate solution, and the producing the potassium salt by metathesis with KCl, all due to solubility issues.

I'll recheck the book when I find it.

Most modern cells, electrolyse K-manganate to K-permanganate, the cells are specially designed to have high Reynolds numbers (I assume that means high turbulence) to inhibit crystallisation of the poorly soluble K-permanganate.

What I am trying to achieve is basically what you have stated above, but only because I have plenty of cheap solid NaOH, but no solid KOH for the fusion process.

Regards, Xenoid

[Edited on 15-9-2007 by Xenoid]

12AX7 - 15-9-2007 at 21:02

Wouldn't high turbulence tend to cause excessive nucleation and a dispersion of KMnO4 rather than crystals?


Xenoid - 15-9-2007 at 21:21

Originally posted by 12AX7
Wouldn't high turbulence tend to cause excessive nucleation and a dispersion of KMnO4 rather than crystals?


I guess the high turbulence is to stop crystallisation building up on the anodes and other parts of the system ie. cathode separators, pipe work, etc. and "clogging up" the cell, with presumably disasterous consequences.

Regards, Xenoid


ciscosdad - 17-9-2007 at 20:01

I got the impression the turbulence was to control deposites other than crystallization. "Oxidic deposites" I think was the term used in a patent I read. MnIII and MnIV I guess. The mother liquor spend a lot of its time a cruddy mess.

From "Industrial Electrochemistry"

Xenoid - 17-9-2007 at 20:41

Here's where I got my information about turbulence!

Sorry it's only a screen dump from Google Books!

Regards, Xenoid

PM-3a.gif - 31kB

ciscosdad - 17-9-2007 at 23:14

Fascinating. I had no idea that violent movement would inhibit crystallization. That must make the operation of a crystallization tank rather simple. The decrease in velocity on entry into the larger volume would seem to do it. It seem these amateur cells will be better with somewhat frantic stirring until harvest time.

Xenoid - 17-9-2007 at 23:24

Well I've just completed my latest attempt at making K-permanganate.

The cell is the same as last time except I have made a new lid from 15 mm. thick, grey, PVC sheet which can withstand the high temperature and oxidising conditions better. I also added a 100 mm. long, 6 mm. diam. SS tube to the vent hole, to stop the clear PVC vent hose softening and distorting. Initially I started the cell operating from a 5 Volt, constant voltage computer SMPS. The cell was running at about 7.5 - 8 Amps for about 3 hours but I found it was impossible to maintain a constant temperature, there was a lot of thermal runaway. Since I was operating at 80 oC. and I had to leave the cell running overnight, any thermal creep could have resulted in the cell boiling, which would have been a bit of a disaster to say the least. I was forced to revert to my constant current supply which has a maximum output of 3.5 Amps. I used this for the rest of the run, which lasted for 33 hours, nearly 3 times the length of the last run. The cell was stirred at about 550 rpm., about 5 times faster than the previous attempt.

At the end of the run, 25g of solid KCl was added and heating and stirring continued for a further 15 mins. The cell was then slowly cooled to room temperature and then placed in a refrigerator at about -5 oC. After several hours the cell was tentatively opened, and the dark supernatant liquid was poured off. Lining the bottom of the container to a depth of about 15 mm. was a mass of aciculate K-permanganate crystals along with some remnant brown MnO2.H2O "crud". I tried to get some pictures, but the whole mass was slowly collapsing into the murk every time the container was moved. The flash photo reflected off the insides and a shot from further out wouldn't focus properly. Any way, I scraped it all out and left it in a fine kitchen sieve, in the fridge, to drain overnight.

Next morning, I washed most of the "crud" off the crystals, rather stupidly using room temperature water instead of ice cold water. This resulted in some of the finer grained "product" being washed through. The crystals were then put in a desiccator under vacuum for a few hours. With the sun coming out, I then air dried them for a few hours, they are now back in the desiccator. I haven't weighed them yet as they are still damp, it won't be a great yield, but at least I'm making some progress.

Maybe it just requires more current for a longer time.

I still have the option of concentrating the remaining solutions and trying to extract a bit more K-permanganate, I'm still thinking about the best way of doing this.

Image 1: Cell with new lid, operating at nearly 80 oC.
Image 2: K-permanganate crystals, still slightly moist, after air drying. Watchglass is about 100mm. diameter.

EDIT: Well, the K-permanganate has dried overnight in a desiccator, the crystals have a slight "dusting" of brown MnO2. The weight was 6g woohoo! The theoretical amount of K-permanganate from this run should have been over 30g. In retrospect, if I had known there was a cluster of crystals in the bottom of the cell, I would have left it in the refrigerator for a much longer time. The dry air would have promoted evaporation and there would have been a bigger yield. I have added a few seed crystals to my remaining solutions to see if I can extract some more K-permanganate!

Regards, Xenoid

[Edited on 18-9-2007 by Xenoid]

Perm2.jpg - 34kB

DerAlte - 18-9-2007 at 21:59

As usual, congratulations, Xenoid. You got the spiky crystals. The Japanese gentlemen have been vindicated.

Sure, the yield seems poor. But the principle is proven. Same as with my earlier wet methods up top of this thread. I intend to do a carefully measured run again on that, using what I consider the best methodology in light of my experience to date. I hope you will do the same with your fusion/electrolysis method.

As my old drill sergeant used to say, "The easy way ain't easy, but the 'ard way's bloody 'ard" Producing prmanganate is like that.

Regards, Der Alte .

Refinement and Optimisation

ciscosdad - 18-9-2007 at 23:31

@DerAlte and Xenoid

Lovely work.
I look forward to seeing the refinements we all hope you will be making.
The electrolytic process looks the most promising for small scale prodution. Perhaps the secret may be to find a way to precipitate the KMnO4 without contaminating the solution with chloride or whatever so the solution can be reprocessed.
Would it be better to do the precipitaion with KOH? (Assuming it is available). I have a little solid, but have not seen any easily available sources unless one counts chemical suppliers and their high altitude pricing. Did someone here mention an OTC source of 40% solution?

not_important - 19-9-2007 at 00:57

K2CO3 is another alternative to KCl. It's sold in pottery supply stores as potassium carbonate or pearl ash.

Remove some of the solution from the cell, add a hot saturated solution of K2CO3 in excess, which will drop out any Mn(II) present. Filter while hot, return solids to the cell, cool filtrate and filter off KMnO4, return filtrate to the cell.

Xenoid - 19-9-2007 at 07:28

Originally posted by not_important
K2CO3 is another alternative to KCl. It's sold in pottery supply stores as potassium carbonate or pearl ash.

If only this was the case. I don't think any pottery suppliers in NZ carry K2CO3. I've tried to find it in the past. It's actually only rarely used by potters, I think it causes cracking of the glaze. They tend to use K-feldspar (orthoclase) in their formulations!

Regards, Xenoid


ciscosdad - 19-9-2007 at 17:29

Do you think the KMnO4 produced was hydrolysing while the cell was cooling? If so, and if it is an equilibrium mix, then I expect so, then rapid cooling after K addition would help.
Further on the equilibrium issue, if the hydrolysis is proceeding in the solution , then the oxidation needs to exceed the hydrolysis rate by as large a margin as possible. Could the current per volume be a factor? ie, the oxidation is not proceeding fast enough to stay ahead?

Of course, the higher current may give heating issues and require active cooling.

Xenoid - 19-9-2007 at 18:53

@ ciscosdad

Hi, I'm not convinced that K-permanganate hydrolises all that quickley in alkaline solutions. I would prefer slow cooling to try and get reasonable sized crystals which make separation quite simple (with a fine sieve). I think the electrolysis should be done at a higher current. The Japanese boffins used 20 Amps in a 1 litre container with an anode current density of 40 mA / cm^2. In contrast my current density was only about 12 mA / cm^2, (3.5A and anode area of about 300 cm^2). I wasn't too worried about this, because the commercial cells operate as low as 5 or 10 ma / cm^2. But since this procedure is a little different, higher current densities may be called for. That is why I tried to run the last cell at 8 Amps. Unfortunately, one really needs a constant current supply for this type of electrolysis, otherwise it is impossible to keep the cell at a stable temperature. I'm currently (no pun intended) looking out for a suitable supply!

Regards, Xenoid

Current Control

ciscosdad - 20-9-2007 at 20:52

This problem of Xenoids applies to chlorate making as well.

The computer power supplies we all have are great for supplying nice fixed 5volts, but there is a problem controlling the current. A variable resistor in series comes to mind, but the resistance range we want is a few tenths of an ohm.
It occurred to me that a water resistor may do the trick to lose that critical 2 volts or so.
Has anyone seen anything that will accept a couple of hundred watts if required at around 20 - 30A?

Google gives lots of hits, but all High voltage or high power (motor start etc).

How stable can these devices be? They will not be much use if they vary too much over time (heat or electrode erosion).
Stainless Steel electrodes and Sodium Carbonate or Hydroxide would be better than salt I would think, with a little Dichromate as a depolariser.
I have visions of a 20 litre drum (or 200l :o) full of water with a kilo or so of Na2CO3 and 2 long SS electrodes hung off the sides. Vary the position and/or the electrolyte concentration to change resistance.
Do any of our electronics guys have any input?

Xenoid - 20-9-2007 at 21:16

Actually, constant current isn't really required for chlorate making because the temperatures are relatively low. Sure the chlorate cell may get hot, but the temperature rises until the cell is at equilibrium with the surroundings. Some people put their cells in water, but personally i've never had a problem, running a 5 litre cell at about 15 - 20 amps from a 5 volt computer supply, the cell tops out at about 50 - 60 oC.

The problem with a permanganate cell is that you have to apply external heat as well as the heat from the electrolysis. With the hotplate cycling, any increase in external heat causes a decrease in resistance of the cell, so it draws more current, causing more heating, etc. etc. It is like trying to balance something on a knife edge. Slight increase in heat and the cell runs away and starts boiling, too little heat and the cell drops below optimum permanganate production.

I'm actually kicking myself, because a few months ago there were a couple of 100 Amp constant current power supplies for sale on the local online trading site. I decided not to bid because at that time I couldn't really see a use for them. Now I'm just hoping something similar will come up soon!

I think I've mentioned it before, in another thread, but if you do need a resistor for your chlorate cell, a stripped gouging rod has a resistance of about 0.2 ohms and makes a useful high wattage variable resistance for limiting current!

Regards, Xenoid

[Edited on 20-9-2007 by Xenoid]

12AX7 - 20-9-2007 at 21:51

OMG... a cell... in series with a water resistor...

It's like...well fuck...why don't you put another chlorate cell in series with it then!?...

I use a chunk of steel coat hanger. Brass braze copper wires onto the end. Solder won't do, for what should be obvious reasons. Of course I don't have as nearly stable a power supply (OC = 8V or so, and the characteristic probably isn't perfectly resistive due to the choke input filter), but the ballasting effect is notable.

A stick or TIG welder, set to DC output and relatively low current (many go down to "only" 40A), would suffice.


ciscosdad - 20-9-2007 at 23:02

Point taken Tim.
However, the voltage across any extra cell(s) is fixed, and is probably inconveniently just a bit much to fit into the available 5v. I have played with all sorts of combinations of series and parallell arrangements, without success, unless I have a variac in there. Perhaps I'm just too fussy.
The variable resistor suits my frugal approach to this. The use of a gouging carbon @~.2ohm (Thanks Xenoid!) is perfect.

Xenoid - 20-9-2007 at 23:44

@ ciscosdad

Here are a couple of my 0.2 ohm resistors, not sure what wattage they would handle, at least a 100 W I guess. I mounted the rods in tool clips (Terry Clips) and made some connectors from strips of SS squeezed around 9.0 and 9.5 mm. drills using a vice and hammer. Make sure the connectors grip the rods tightly when the screws are done up. Actually, in retrospect, 2 or 3 mounted side by side on a single board would be a better idea than the individual mounting.

Regards, Xenoid

Resistors.jpg - 18kB

Well, I'm a sucker for punishment.....

Xenoid - 22-9-2007 at 11:24

.... and against my better judgement I decided to have another go at K-permanganate.
I've brewed up another batch of (hypo)manganate, I have however, made a couple of changes to the procedure used in the first fusion.

Firstly, I dispensed with the water, (NaOH as 50% solution). The only reasons I could see for doing this were that it produced better mixing or demonstrated that recycled NaOH solution could be used in the procedure.

Secondly, I replaced NaNO3 with KNO3 this should provide sufficient K+ ions in the system to alleviate the need to perform the double dissolution at the end of the electrolysis.

The procedure is the same as given earlier in the thread, and proceeded smoothly, without any fumes or smell. The final yield was about 250g from .5 moles MnO2 (pottery grade), 3 moles NaOH and 1 mole KNO3. This is enough for 2 runs in my 500 ml. cell.

Image 1: Melt proceeding smoothly.
Image 2: Product after cooling.

Regards, Xenoid

Manganate.jpg - 23kB

ciscosdad - 23-9-2007 at 15:12

Thanks Xenoid.
I like the look of your resistor design. Simple and effective.

We are all looking forward to the results from your new batch. Go to it.

Last, final, ultimate effort!

Xenoid - 28-9-2007 at 14:44

I've finally run a batch of the last manganate fusion through my cell! This is my final effort!

The results were mildly encouraging, but it would seem to me that one would have to be pretty desperate to make your own permanganate. If you can aquire the MnO2, the sodium or potassium hydroxide, and the sodium or potassium nitrate required for this process, then you can probably aquire the potassium permanganate as well!

Half the quantity of manganate made in the last fusion ~125 g. was put in my 500 ml. cell (see above). Stirring was set for 500 rpm and the temperature ranged from about 70 - 80 oC.

I put together a simple transformer based power supply, running of a variac and set the current for 10 Amps. This current is nearly 3 times what was used in the previous run and upped the current density from ~ 11 mA/cm^2 to ~ 30 mA/cm^2. The cell was run for about 24 hours.

The cell was slowly cooled overnight and then placed in a refrigerator for 2 days at about -5 oC. When the supernatant liquid was poured off there was a mass of fine crystals on the bottom of the cell! The crystals were treated as described in the last procedure (see above) however this time I washed them with ice-cold water. During the drying process some of the crytals seemed to partially redissolve and formed fine crusty material on further drying. This remained K-permanganate however, as can be seen when a few tiny grains were dropped in water to test it.

The watchglass is 150 mm. in diameter. Final dried weight was about 17 g. almost 3 times the previous effort!

So, it would seem that more electricity is better for a simple cell of this type!

Regards, Xenoid

K-Permanganate.jpg - 14kB

chloric1 - 28-9-2007 at 17:48

nice but chlorate is still more fun to make. I turn edible salt into a poisonous powerful oxidizer:D. Last time I looked the Menard's by my house still has 5 pounds of KMnO4 for $17.

Xenoid - 28-9-2007 at 18:14

Originally posted by chloric1
nice but chlorate is still more fun to make. I turn edible salt into a poisonous powerful oxidizer:D. Last time I looked the Menard's by my house still has 5 pounds of KMnO4 for $17.

Duh.....! I too make chlorate (and perchlorate) and can buy permanganate for a similar price.... so what! The point of the thread was to try and find a practical way for the amateur chemist to make it at home. I'm not entirely sure this procedure is all that practical, though!... :(

Regards, Xenoid

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