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Author: Subject: Permanganates
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[*] posted on 28-8-2007 at 18:31


Quote:
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.
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[*] posted on 28-8-2007 at 19:59
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 It is now, probably, very much out of date.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 .... :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 It's a bit out of date, I realise, I'm not sure what the current thinking on this material is.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]
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[*] posted on 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
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[*] posted on 30-8-2007 at 11:50


Keep up the good work, Xenoid! We are all rooting for you... DerAlte
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[*] posted on 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]

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[*] posted on 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
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[*] posted on 3-9-2007 at 12:35


Quote:
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

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[*] posted on 3-9-2007 at 18:04
MnO2.H2O by electrolysis of MnSO4


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]
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[*] posted on 3-9-2007 at 18:10


Lead electrodes should work.
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[*] posted on 3-9-2007 at 19:53


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



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[*] posted on 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)
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[*] posted on 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
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[*] posted on 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.
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[*] posted on 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
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[*] posted on 4-9-2007 at 15:18
Equilibrium


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]
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[*] posted on 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

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[*] posted on 4-9-2007 at 21:28


Interesting...

From Ullmann's:


Quote:

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]
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[*] posted on 4-9-2007 at 22:34


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

Quote:
... 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
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[*] posted on 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]
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[*] posted on 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]
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[*] posted on 5-9-2007 at 12:42


Not_important said

Quote:
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.
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[*] posted on 6-9-2007 at 09:44


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

Quote:
@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.


Quote:
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:

Quote:
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.


Quote:
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.


Quote:
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.


Quote:
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]
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[*] posted on 6-9-2007 at 12:43


@ Cesium
Quote:
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
Quote:
…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
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[*] posted on 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.
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[*] posted on 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.
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