Permanganates

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

MnO2 Electrodes etc

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.

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

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]

Patents

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.

Regards, Xenoid

Potassium Permanganate

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.

Regards, Xenoid

@ 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!...

Edit:
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]

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]

MnO2 from MnSO4

Update

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 ....

@ 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]

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

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]

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

Regards, Xenoid

MnO2.H2O by electrolysis of MnSO4

Equilibrium

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

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!

@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]

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.

AAAAARGH!....... DISASTER HAS STRUCK.......
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]

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.
@Chief
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.

Only partial success, again!

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

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

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]

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

Turbulence

From "Industrial Electrochemistry"

Here's where I got my information about turbulence!

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

Regards, Xenoid

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]

Refinement and Optimisation

Optimisation

Current Control

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 ) 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?

@ 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

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

.... 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

Last, final, ultimate effort!

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

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

Quote:

Originally posted by chloric1 nice but chlorate is still more fun to make. I turn edible salt into a poisonous powerful oxidizer. 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