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Author: Subject: Permanganates
Xenoid
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[*] posted on 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

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DerAlte
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[*] posted on 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.

Regards,
Der Alte.
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[*] posted on 23-8-2007 at 13:11


Quote:
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
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[*] posted on 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.

Regards,

Der Alte
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[*] posted on 23-8-2007 at 15:45
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.:(
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[*] posted on 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.

Regards,

Der Alte.
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[*] posted on 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…

Regards,
Der Alte

[Edited on 24-8-2007 by DerAlte]
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[*] posted on 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

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[*] posted on 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.
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[*] posted on 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
http://ceramic-materials.com/cermat/education/139.html
http://ceramic-materials.com/cermat/education/147.html
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[*] posted on 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]
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[*] posted on 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
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[*] posted on 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:
Quote:
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.

Regards,
Der Alte.
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[*] posted on 27-8-2007 at 15:15
Patents


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.
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[*] posted on 27-8-2007 at 17:34


Quote:
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
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[*] posted on 27-8-2007 at 20:16
Potassium Permanganate


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.
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[*] posted on 27-8-2007 at 21:05


Quote:
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
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DerAlte
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[*] posted on 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...

regards,

Der Alte
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[*] posted on 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!... ;)

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]
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[*] posted on 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

or

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).

Regards,

Der Alte.

[Edited on 28-8-2007 by DerAlte]
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[*] posted on 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
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[*] posted on 28-8-2007 at 15:51
MnO2 from MnSO4


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?
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[*] posted on 28-8-2007 at 17:24


Quote:
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]
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[*] posted on 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.)
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[*] posted on 28-8-2007 at 18:20


@Xenoid

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
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