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Author: Subject: Permanganates
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[*] posted on 6-9-2007 at 21:51


The numbers I gave are sort of an average, some metallurgical references run higher once you're out of the dark/dull/low red heat; a few ceramics refs run a bit cooler, possibly because those are almost always in a kiln and the enclosed space reduces heat loss from the surfaces - hot stuff in the open will have a cooler skin than bulk temperature.
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[*] posted on 8-9-2007 at 07:18


Has anyone any source for the solubilities of
- KMnO4
- Ba(MnO4)2
- Ba(NO3)2
in dependency from the temperature ? I can't find anything on the net except values without the temperature or at 25 [Celsius] or something. But I go for Ba(MnO4)2 from Ba(NO3)2 and KMnO4,
and need those data for optimum performance.
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[*] posted on 8-9-2007 at 07:44


Besides: How sounds 15 $/kg for KMnO4 ? 10 kg at 10$/kg ??
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[*] posted on 8-9-2007 at 09:25


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

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

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

Regards,

Der Alte

[Edited on 9-9-2007 by DerAlte]
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[*] posted on 8-9-2007 at 13:03


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

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

Mainly until now I mix concentrated solutions of Ba(NO3)2 and KMnO4, watch the temperature and try to find the best point of harvesting the Ba(MnO4)2, but always I get some (50+%) amount of useless KNO3/KMnO4/BaNO3/Ba(MnO4)2 - mixture, which I need to separate ...
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[*] posted on 8-9-2007 at 15:34


Mr. Cheif,

Here is some information I dug up for you.

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

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

Probably at higher temperatures, it decomposes too readily.

No info about BaMnO4 except that it's soluble. Sorry it's rather exotic.
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[*] posted on 8-9-2007 at 18:32


Well, I'm back on the job!

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

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

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]

Permanganate.jpg - 28kB
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[*] posted on 8-9-2007 at 21:39


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

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

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

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

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


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

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

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

http://en.wikipedia.org/wiki/Solubility_table

regards,

Der Alte
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[*] posted on 8-9-2007 at 21:50


Keep up the good work Xenoid!

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

http://en.wikipedia.org/wiki/Solubility_table


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

[Edited on 8-9-2007 by Cesium Fluoride]
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[*] posted on 9-9-2007 at 15:46


I was looking over some patents on the weekend, and found that there are a couple that appear to be further developments of the Japanese patent.
The idea is to replenish spent permanganate etch solutions.
I believe the numbers are 4835095 and 4911802. Not entiirely sure (my copies are at home). If you are interested look at the last few entries on the list I posted (page 5 this thread). The main difference in the new development is a diaphragm around the cathode containing concentrated alkali. The main solutions are about 1 molar alkali, and the current efficiencies look a bit poor at first glance.
@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.:(
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[*] posted on 11-9-2007 at 15:59


Quote:
Originally posted by Cesium Fluoride
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



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

Isolation is perhaps unnecessary if you're planning on using it in solution anyways, and if some base is OK, and it's concentration can be determined by any number of analytical methods.
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[*] posted on 11-9-2007 at 16:43


@ S C Wack
Fascinating but confusing.
Do you think it is concentrated enough to get a precipitate of KMnO4 with (say ) KCl or aybe KOH? Would it survive boiling long enough to concentrate?
Could it work with KOH?
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[*] posted on 11-9-2007 at 17:38


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

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

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

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

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

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

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

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

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

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

Regards, Der Alte
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[*] posted on 11-9-2007 at 17:45


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

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

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[*] posted on 11-9-2007 at 20:09


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

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

I have a hunch that Ullmann's is closest to the truth as its the only solid explanation I've seen given for the prevalence of KMnO4 over NaMnO4.
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[*] posted on 11-9-2007 at 20:27


@ Cesium Fluoride

You said

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


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

Regards, Der Alte
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[*] posted on 11-9-2007 at 20:42


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


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

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

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

[Edited on 11-9-2007 by Cesium Fluoride]
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[*] posted on 11-9-2007 at 21:08
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
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[*] posted on 11-9-2007 at 21:16


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

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

Tim

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




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[*] posted on 11-9-2007 at 21:20


Yes, I have even seen the greens of manganate form when boiling a NaOH solution with MnO2. Of course, this color signifies almost nothing from a practical standpoint.
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[*] posted on 11-9-2007 at 21:56


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

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[*] posted on 11-9-2007 at 22:38


Yeah! The solution is effectively black!

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

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

Regards, Xenoid
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[*] posted on 12-9-2007 at 10:43


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


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

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

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

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

Regards, Der Alte
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[*] posted on 12-9-2007 at 23:15


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

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

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

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

Regards, Xenoid

Permanganate3.jpg - 26kB
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[*] posted on 15-9-2007 at 08:23


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

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

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

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

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

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

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

[Edited on 16-9-2007 by Antwain]
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