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Author: Subject: Permanganates
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[*] posted on 19-7-2007 at 23:46


Thats what about I am here: I myself do need the barium permanganate and am trying to get something out of everyone here ...
But maybe as long as the Ba(MnO4) is isolated it decomposes at 100 [Celsius]. With higher temperatures it might be stable under oxidizing conditions, like in a nitrate melt ...
Because of that I was asking about the decomposition characteristics of nitrates: If there exists any decomposition rate (instead of a phase-change-decomposition at a fixed temperature) and thereby a half-life-time (mathematically a consequence of the decay-law) one could trade of the decomposition/formation times by temperature-regulation, which means one would have to melt it not too hot and not too long, and ready would be some 2figure-percentage of Ba(MnO4)2.
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[*] posted on 23-7-2007 at 11:25


Hello there !! I have made some experiment with nitrates: NaNO3 melts at 306[Celsius] and decomposes over a temperature-range up to over 600[Celsius] and still has oxygen in it !!
So its a question of reaction-dynamics to be optimized. Most probably manganates have too some temperature-time-range of decomposition, but if its as wide as with the NaNO3 I don't know yet. Any Ideas?
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[*] posted on 23-7-2007 at 14:31


I've had a melt of dark green sodium manganate (or -ite, as the case may be) up to red heat (~660C) without decomposition. Ditto potassium dichromate, for that matter.

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[*] posted on 26-7-2007 at 04:33


Here's an interesting mixture:

KNO3-LiNO3-NaNO3 44.9-37.3-17.8

melts at 120 °C !
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[*] posted on 5-8-2007 at 10:57


NOTES ON ACETONE EXTRACTION OF KMnO4

To obtain good results from this process the following precautions must be taken:

(1) The acetone must be anhydrous or very nearly so. Possible drying agents are CaCl2 (AFAIK)* or anhydrous CaSO4. Or dry over conc. H2SO4 if you have a proper desiccator. The product to be abstracted must also be dry.

(2) Cool the acetone, apparatus and product to at least 0C (I use a freezer and try for -10C). As indicated by not_important in a post above dated 26-6-2007 and discussed in subsequent two posts by Der Alte, acetone does react with KMnO4, due to enol formation. A low temp. minimizes this.

(3) Evaporation under reduced pressure would be the ideal (and hence recovering the acetone, reducing temp and excluding moisture ). If evaporated in a air stream, especially here where the humidity is high, the acetone solution, cooled by its own evaporation, tends to deposit and dissolve water in the liquid.

(4) Due to the presence of a large amount of other salts in the dried product, which are (AFAIK) all insoluble, you get a large absorption of acetone in the product. This can be minimized by filtering under suction in a Buchner funnel or equivalent. Otherwise up to half the desired product can be lost in wetting the solids. Yields are poor enough without this added complication. (More on this later).


WRT the reaction of KMnO4 with acetone, I have performed the following experiments:

(1) Took the wet solid from the filter and mixed with water to dissolve the other products and remaining KMnO4. After an short induction period the KMnO4 discolored and eventually precipitated MnO2 leaving a clear solution. Tests showed the presence of acetates (smell) and Oxalates (precipitation). Formates may also be produced but I don’t know a simple test for them offhand.

(2) Drying and weighing the MnO2 so precipitated showed that about 50% of the KMnO4 had been retained with the solids and hence lost.

Nobody said this would be an easy preparation of KMnO4! More anon. Next post gives the only bit of information I have been able to find on the solution of inorganic salts in acetone. It’s from an old and obscure source.

*NOT CaCl2! Apparently, I have found out, although it does not form an addition compound, it encourages aldol condensation of acetone. Use Ca SO4 (anh.), K2CO3 (anh.)

Regards,

Der Alte

[Edited on 10-8-2007 by DerAlte]
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[*] posted on 5-8-2007 at 11:00


ACETONE—SOLUBILITY OF INORGANIC SALTS.

(This may be of general interest.}

W. H. Krug and K. P. McElroy have investigated the solubilities of various inorganic salts in acetone :

CHLORIDES.
Insoluble—Sodium, potassium, ammonium, anhydrous nickclous, and mercurous.
Very sparingly soluble—Calcium, barium, anhydrous strontium.
Somewhat soluble— Anhydrous cadmium, Freely soluble—Ferric, zinc, anhydrous cobaltous, crystalline cupric.
Very freely soluble—Mercuric.

BROMIDES.
Slightly soluble—Potassium, sodium,
Freely soluble—Anhydrous cadmium.

IODIDES.
Soluble—Potassium, mercuric.

CYANIDES.
Freely soluble—Mercuric.

SULPHOCYANIDES. (i.e. thiocyanates - Der Alte)
Freely soluble—Potassium, ammonium. Soluble—Ferric, cobaltous.
Insoluble— Nickelous.

NITRATES.
Insoluble—Barium, bismuth.
Very slightly soluble—Sodium, potassium, lead.
Slightly
soluble—Ammonium, crystalline nickelous.
Soluble—Silver.

CARBONATES.
Insoluble—Potassium, anhydrous sodium.

SULPHATES.
Insoluble—Anhydrous copper, potassium, anhydrous ferric, crystalline ferrous ammonium,
anhydrous ferrous.

ACETATES.
Insoluble—Magnesium, sodium, calcium.
Slightly soluble—Crystalline copper, crystalline
lead.
Soluble—Crystalline zinc.

MISCELLANEOUS.
Insoluble—Potassium ferrocyanide, mercuric sulphide, ferric pyrophosphate, ammonium
molybdate, ammonium oxalate, ammonium tartrate.
Very slightly soluble—Potassium chlorate.
Freely soluble—Boric acid, malic acid, tartaric acid, oxalic acid.

The solubility has been determined for the following at 25° C.: One
hundred parts of acetone by weight dissolve :
Potassium iodide 2.930;
Potassium bromide 0.023 :
mercuric chloride 50.990;
mercuric iodide 2.090 ;
anhydrous cobaltic chloride 8.620.

Acetone—Solubility in Dextrose Solutions ( Glucose}.—The same authors determined the solubility of acetone in glucose of different strength at 25° C. One hundred gm. of glucose solutions dissolve :

Per cent, of glucose. Grammes of acetone.
10 747.86
20 237.71
30 146.30
40 72.72
5) 32.70 —

Chem. News, 1892, Ixv., 255, from Jour, analyt. appl. Chem.

Regards,

Der Alte.
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[*] posted on 6-8-2007 at 14:45


Thx for the good info!

(1) Purifying the acetone by destillation is most certainly a good idea anyways, to remove other crap like aldehydes. To properly do this, however, you need a strong oxidizer like permanganate :(

(3) Mhmm reduced pressure is a problem for the less sophisticated home chemist like me ;)
If moisture is the problem, maybe slightly heating the acetone in a destilling apparatus would do the trick? This way the acetone could be recovered as well. Question is: How much will the raised temperature increase the rate of permanganate-acetone reaction?

(4) Vaccum filtrating is also a less available method to the amateur experimenter. Taking more acetone, and maybe putting the filter cake into a cloth and wringint it out (using good gloves of course!) might do the trick as well, however...

So: If the acetone could be recovered, all this wouldn't be a big problem!

btw I tried using mangane dioxide instead of mangane sulfate and it didnt react with the sodium hypochlorite at all :( But I still need to try the two-stage fusing process...

Regards

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[*] posted on 6-8-2007 at 19:31


@Ballermatz:-

re reaction rate, probably about double for every 10C rise in temp., the usual exonential rate per theory, unless the formation of the MnO2 catalyzes the reaction, as it does in H2O2 decomposition. I know that one trial with acetone at 30C produced a lot more MnO2 than doing it at 0C or lower.

The products in water/acetone solution appear to be oxalate and acetate, as far as I can detect, as I said. But Hot solid permangante and acetone may go all the way to H2) and CO2, and very rapidly, as suggested by the reference cited by not_important. Haven't tried heating acetone to BP with KMnO4 yet, but sounds like one should be cautious at least. If you've ever seen what it can do to glycerol.... enough said!

Quote:
btw I tried using mangane dioxide instead of mangane sulfate and it didnt react with the sodium hypochlorite at all


You need some OH- ions around, from NaOH or Na2CO3 to get sodium permanganate. Remember, this oxidation is rather slow.

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[*] posted on 7-8-2007 at 18:58


A BIT OF THEORY- for those that like such things.

The aim of the process has devolved around maximizing the yield of KMnO4 whilst minimizing that of KClO3. The process of using the sodium salts is attractive on paper but too awkward in practice – dealing with viscous solutions produced is next to impractical.

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

2MnO2 + 3KClO + K2CO3 --> 2KMnO4 + 3KCl + CO2 simultaneous with
3KClO -->2KCl + KClO3

The equilibrium constants of these reactions can be written as (at fixed p,T)

K1=( [KMnO4]^2)( [KCl]^3)( [CO2] /([MnO2]^2)([KClO]^3)([K2CO3])
K2=([KCl]^2)( [KClO3]/([KClO] ^3)

(Sorry, best I could do in this format)

Dividing these expressions, using a bit of algebra and rearranging, the ratio of concentrations KMnO4 to KClO3 at equilibrium will be:-

[KMnO4]/[KClO3]=
[MnO2(s)][KCO3]^(1/2) [KCl]^(1/2) / [KClO]^(3/2) [CO2]^(1/2)*(K1^(1/2)/K2)

This tells us that the KMnO4/KClO3 ratio varies as the -3/2 power of the KClO concentration – i.e dilute solutions increase the ratio. But they also slow the reaction rate. Instead of using the 15% KClO if one uses say 3%, the yield should be increased by a ratio of 11.2. Increasing the carbonate concentration also has a less dramatic effect, a ratio of 2.2. The same factor is helped by the elimination of CO2 from solution, which requires a higher temperature.

Now, with regard to temperature, I have come to the conclusion that my assumption that higher temperatures would degrade the product was wrong. Both wanted and unwanted products should maintain the ratio constant, as far as I can see, since both ought to have the same exponential rate factor.

I have been unable to test these hypotheses because I’m out of KOH and KCO3. I did a quick check with sodium salts and qualitatively it did seem to be true.

Regards,

Der Alte
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[*] posted on 9-8-2007 at 10:51


FURTHER OPTIMIZATIONS

As well as decreasing the production of chlorate, large amounts of the chlorides inevitably have to be dealt with. In this respect, the process using potassium salts is far superior because it avoids the required addition of KCl as a final step to produce the KMnO4 from the sodium salt.

Looking again at the reaction

2MnO2 + 3XClO + X2CO3 --> 2XMnO4 + 3XCl + CO2

where X is an alkali metal, we see that for every mole XClO is produced one mole XCl and 2/3 mol XMnO4.

(IF X=Na, and you use Clorox, you add another Mole NaCl for every NaClO used up. I have done a weighing of Clorox evaporated to dryness and accounting for chlorate production as well, it is essentially a 50/50 mix of hypochlorite and chloride. The MSDS says it also contains <1% NaOH. Hence do not use Clorox for oxidation of MnO2 to MnO4- ions. It is fine for the first step of making MnO2 from Mn++ ions because the MnO2 can be washed free of chloride.)

The final step, if using Na salts, is conversion to potassium permanganate. If you use KCl (as I have been) you get more mol NaCl

NaMnO4 + KCl --> KMnO4 + NaCl

Better to use carbonate or nitrate (not sulphate, it too is rather insoluble) and avoid this:-

2NaMnO4 + K2CO3 --> 2KMnO4 + Na3CO3.

In the worst case Na --> K scenario (using Clorox) 2 mols KMnO4 incur the production of 7 mols NaCl; but half (roughly) of the NaClO produces chlorate so in terms of overall reactions, 7 mols of chloride and one of chlorate are produced for every mol of permanganate. Pretty piss poor! If you use hypochlorite from the calcium route, as recommended, this is reduced to one of chlorate and 4 of chloride.

In contrast, using only K salts, only about 3 mols KCl and one of KClO3 have to be dealt with. Further, the presence of a large excess of K+ ions reduces the solubility of the permanganate by common ion effect (and, unfortunately, also that of the chlorate). The final separation of permanganate from chlorate in water solution is very difficult, so there we have to resort to acetone.

There are other approaches to permanganate production in aqueous solution. I shall suggest a few next time.

Regards,

DerAlte
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[*] posted on 9-8-2007 at 16:30
Exotic Yellow Manganese Compounds


I was reading over this thread while doing some research on permanganate preparation (I'm still going down the alkali fusion path), and I noticed this.

Quote:

Mn compounds have many colors, red, purple, green, blue, brown, black – all except yellow. At least, can’t think of one.


Well, I happened to read in Ullmann's about these two compounds and I just had to post:

Quote:
Manganese ethylenebis(dithiocarbamate)[12427-38-2] , Maneb, (CH2NHCS2)2Mn, Mr 265.3, is a yellow powder. It is prepared by the addition of an aqueous solution of ethylene diamine and ammonia to carbon disulfide, followed by neutralization with acetic acid and precipitation with MnSO4 or MnCl2. The compound is an important fungicide.

Methylcyclopentadienylmanganese tricarbonyl[12108-13-3] (MMT), Mr 218.08, r 1.39 g/cm3, bp 233 °C, is a yellow liquid that is insoluble in water but soluble in organic solvents. Several synthetic routes exist for this compound. For example, manganese(II) chloride may be allowed to react with cyclopentadienylmagnesium bromide, C5H4MgBr, to form biscyclopentadienylmanganese, an intermediate that reacts with carbon monoxide to give the tricarbonyl. This is then methylated in the presence of Friedel – Crafts catalysts. The product finds limited use as an antiknock additive in motor fuels and as a combustion aid in heating oils.


[Edited on 9-8-2007 by Cesium Fluoride]
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[*] posted on 9-8-2007 at 17:18


@ Cesium Fluoride - OK! I think that covers the spectrum! Regards, Der Alte
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[*] posted on 11-8-2007 at 17:19


I will make permanganate with this reaction:

2 MnCl2 + 5 NaClO + 6 NaOH => 2 NaMnO4 + 9 NaCl+ 3 H2O

The problem is separate the permanganate of the salt. I would use a barium salt but I don't have. I had planned to use ammonium chloride alternatingly to precipitate the permanganate. Then, I mix to the ammonium permanganate with KOH to make the KMnO4. But did the KMnO4 react with the ammonia? Does somebody know some acid but weak that MnO4-? The CO3--? Is it dangerous the NH4MnO4?
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[*] posted on 11-8-2007 at 21:53


@ Filemon

You have the gist of a method, but using NaOH in the manner you suggest will raise the pH too high and form manganate, and also saddle you with excess NaCl because some is required for the Mn++ --> MnO2 step. Read my previous posts carefully!

I am going to post a method via the manganate route shortly.

The idea of using the ammonium salt had occurred to me and it's not a bad one. Ther are a few caveats, however. One, I have to rely on Wiki, and Wiki is sometimes slightly tainted by bullshit. For instance, the solubility of barium permanganate given in their solubility tables is quite inaccurate - however, the manganate is fairly insoluble, which gives one a possible route.

NH4MnO4 is somewhat unstable, likely to decompose at 70C or thereabouts. Like dichromate, it decomposes fairly gently, especially in solution. But the Chlorate (which you'll get if you use any hypochlorite) explodes when dried, so care has to be taken. The NH4MnO4 is less soluble than the potassium salt ( about 0.8g/100g aqua @ 20C) IF Wiki is right. I have no confirmation on this. Hence it should be precipitated preferentially from the highly soluble sodium salt if you don't heat it too much. Haven't tried this, but seems a good idea worth trying...

Regards,

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[*] posted on 15-8-2007 at 22:51


ALTERNATIVE METHODS OF PERMANGANATE PREPARATION....

(1)The alkali + oxidant method – see the various threads on this site, and Google for the rest. I have never satisfactorily made it work other than pale colorations of manganate.

(2)Making the manganate in a wet process and converting to permanganate. This may have promise from a rough trial I carried out. This mimics the fusion process in the wet.

The idea is to do the oxidation at pH 14 (roughly 1N in OH-) in a similar fashion to the one discussed above:

MnO2 + KClO + 2KOH --> K2MnO4 + KCl + H2O

The result is a very dark green solution that appears almost black due to suspended MnO2, at least until the dioxide settles. The Manganate so produced easily converts to permanganate by any path that manages to reduce the pH. Any acid will do this: even one as weak as H2CO3

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

with CO2 gas we get 3K2MnO4 + 2CO2 --> 2KMnO4 + 2K2CO3 + MnO2.

So the net result is that for 3 moles MnO2 and 3 of KClO and 6 of KOH plus 2 of CO2 gas we get 2 Moles of KMnO4, 2 of K2CO3 plus 3 of KCl (– and one of MnO2 for recycling!)

If you compare this carefully with the original process done at pH ~11 as earlier you find they are identical in result except for the addition of carbonate. Remember that the chlorate bogey is still around. The excess of K+ ions helps to reduce the solubility product of both chlorate and permanganate, as before.

The reaction appears to go much faster, however; but I wouldn’t swear to it – I only did a cursory trial with small quantities using sodium salts and actually using sodium bicarbonate to do the acidification step, in spite of the fact that the pH of a bicarbonate solution is around 8. Extraction with acetone is still required. I intend to do this one again.

For manganate, however, whereas Barium Permanganate is fairly soluble, the managanate is quite insoluble and could be isolated at this stage..

(3)Use a different oxidant that does not give the chlorate problem. A search of possibilities in alkaline solution did not yield anything to me but there is a possibility in acid solution. PbO2 is a strong oxidizer with a SRP of -1.69 volts in acid solution; sufficient to oxidize Mn++ ions to MnO4- ions (1.52 v).

Possible half-reaction 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. Combining the ½ reactions to eliminate charge, we get:

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

At equilibrium, then,

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

(K is the equilibrium constant) - 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…

(4) Filemon said, above, quote:

“The problem is separate the permanganate of the salt. I would use a barium salt but I don't have. I had planned to use ammonium chloride alternatingly to precipitate the permanganate. Then, I mix to the ammonium permanganate with KOH to make the KMnO4. But did the KMnO4 react with the ammonia? Does somebody know some acid but weak that MnO4-? The CO3--? Is it dangerous the NH4MnO4?”
.
Indeed, the ammonium permanganate is poorly soluble and would be precipitated. However, it is very unstable. If temperatures were kept low enough and no effort made to dry it, one could get some product, I’m sure. However, remember that in the method I have outlined, chlorate is also produced. Ammonium chlorate is notorious for instability – it explodes at c. 100C. The permanganate decomposes, how violently I don’t know, @ C. 70C. Doesn’t sound like a safe mixture to recommend to any novices!

Enough for the time being,

Regards, Der Alte.
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[*] posted on 16-8-2007 at 16:43


@Der Alte

thx for the info!

As for the ammonium permanganate, heres a video of its explosiveness:

http://www.youtube.com/watch?v=ARXBzPhoCRg

The NH4MnO4 route might be an option anyways: If you precipate it at low temp and react with KOH while it is still wet, the ammonium chlorate formed will not be a problem. It is only dangerous when dry (a similar route through ammonium chlorate is used to produce barium chlorate) and will be converted to potassium chlorate anyways...

As for the reaction of MnO2 with KClO/KOH: Maybe the process can be speed up by using MnSO4 instead. MnSO4 should form fusing MnO2 with NaHSO4, which is available by the kilo as a pool water additive. So at least we could have a cheap way of turning the stubborn pottery-grade MnO2 into a very soluble manganese salt ;)

But still, the precipation of produced permanganate remains the main problem here.

This might be of interest also:

http://www.informaworld.com/smpp/content~content=a713704496~...

"A simple and easy preparative route to obtain highly pure permanganate salts via aluminium and barium permanganate is described. Aluminium permanganate is prepared by a known method from KMnO4 and excess Al2(SO4)3, then converted to barium permanganate by reaction with excess barium hydroxide. The residual KMnO4 content is co-precipitated together with BaSO4 or adsorbed on solid Al(OH)3 (which are formed in large amounts during synthesis). The excess Ba(OH)2 is transformed into an insoluble precipitate during heating of the Ba(MnO4)2 solution. Ammonium, zinc, cadmium, magnesium and nickel permanganates were prepared in high purity from pure barium permanganate and sulfate salts. "

Any info on the solubility of aluminium permanganate?

Regards

Der BallerMatz ;)

[Edited on 16-8-2007 by Ballermatz]

[Edited on 16-8-2007 by Ballermatz]

[Edited on 16-8-2007 by Ballermatz]
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[*] posted on 17-8-2007 at 04:15


Mhmm Im not so sure anymore that MnSO4 can be prepared by fusion of MnO2 with NaHSO4/KHSO4, but I found a preparation from another off the shelf-product: iron(II)sulfate. Its available as "iron fertilizer", "lawn fertilizer" or moss removal agent. Strong heating with manganese dioxide will produce the sulphate:

2MnO2 + 2FeSO4 → 2MnSO4 + Fe2O3 + ½ O2

PbO2 could be made from electrolysis of lead metal sheets in dilute H2SO4. Reaction with MnSO4 could be carried out in H2SO4 to precipate lead sulphate.
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[*] posted on 17-8-2007 at 13:35


@ Ballermatz - good video! I assume the decomposition is something like:

2NH4MnO4 --> N2 + 4H2O + 2MnO2

As for the solubility of Al(MnO4)3, it's probably soluble - most permangantes are but it's devilishly difficult to find data. The only insoluble ones, or fairly insoluble, are the sliver, ammonium and at low temps, our friend KMnO4 plus the Rb and Cs salts.

It is unusual to find (nearl) insoluble ammonium salts (except for complexes). The only other one I have is ammonium magnesium phosphate.

According to Wiki, a usually reliable source but not infallible, the solubility of NH4MnO4 is 0.8g/100g aq at 20C

As for 'technical grade' MnO2, it's often pyrolusite ore and contains a fair bit of iron. Iron is a menace. Reagent or 'practical' grades are much purer. The MnO2 used for alkaline batteries starts off quite pure, though less so for the older Zn/C types. If you purify it as I have suggested in the above posts, you'll have a pretty pure product as approx MnO2.H2O with good activity as a fine powder.

For making other manganese(II) salts it's best to produce the carbonate which keeps fairly well in a sealed container. You can then react with virtually any other acid to get the salt you want.

PbO2 -

electrolysis lead anode & cathode in H2SO4 works well. Or you can use NaClO and a Pb++ salt, such as nitrate or acetate (acetate might have a problem). Most other Pb salts are not too soluble.

Regards,

Der Alte.

[Edited on 19-8-2007 by DerAlte]
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[*] posted on 18-8-2007 at 09:15


@Der Alte

You might like the following thread

http://www.sciencemadness.org/talk/viewthread.php?tid=6882


[Edited on 19-8-2007 by Rosco Bodine]
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[*] posted on 19-8-2007 at 09:22


@ Xenoid - having trouble with U2U, have read your message but failed to reply successfully. I haven't seen the quoted text, but offer the following comment for general consumption.

The industrial processes as described in many sources vary and are always vague. Air is always quoted as oxidant - it's free! The oxidadtion is carried out at 'fusion temp' i.e. 400C for KOH, 'low red heat', vague, 'incipient red heat', vaguer, or I have seen 350C. In one case (IIRC) 'wetted and rotated in drums', in another 'spread on a floor and sparged with heated air.'

The only common factor is 'heated for 10-24 hours' and electrolytically oxidized to permanganate.

I believe that amateurs fail because they can neither keep the temperature at the right point, nor for that long. The lab processes suggest the use of an oxidant like KNO3 or KClO3 which should speed things up a bit. Nevertheless, all the times I've tried the yield has been very poor. The reaction is very slow (and so are the methods suggested in this thread). It seems that forming the extra covalent links of Mn to O - essentially adding one to the two already in MnO2- is slow and arduous.

Regards,

Der Alte
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[*] posted on 19-8-2007 at 12:01


@Rosco - had a quick look at that thread. It seemed to me from Hilski's excellent photos that in fact permangante was being produced - the color is quite unique. The manganic sulphate is hard to make but if you dissolve MnO2 in strong H2SO4 it is produced as a GREEN color, very difficult to crystallize out and unstable if you manage it.

The Mn(III) chloride seems to be produced if ice cold conc HCl is added to MnO2. Maybe even MnCl4, transiently . Whatever, you get a deep brown solution. MnF4 is quite stable, IIRC. MnF3 is too, color red says CRC. I don't know the color of the Mn(III) alums. Like the ferrous and ferric ammonium alums, it seems the Mn alums are also more stable than their constituents.

The SEP at low pH of Mn++ --> Mn+++ and Mn++ --> MnO4- are almost identical at 1.5 v.

Electrolytic methods for producing permanganate from MnO2 in high pH environment (KOH) seem eminently practical. I am thinking if trying it once I can figure how to make an anode structure (without the heroic efforts of plating by Dann2 and others which I admire!) MnO2 is conductive but poorly 1-10 ohm-cm.

Regards,

Der Alte

[Edited on 19-8-2007 by DerAlte]

[Edited on 19-8-2007 by DerAlte]
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[*] posted on 22-8-2007 at 18:59


Der Alte,
Ref Electrolytic Permanganate
I found a Japanese patent that looks extremely useful

US 3 986 941

The basics are to electrolyse an MnO2 slurry in a ~20% KOH Solution at ~90 Deg C.
There is an example given using a quantity of 1 litre in a glass beaker.
4.2 moles of excess KOH to 1 mole of MnO2.
The electrodes used are nickel and iron, but stainless steel will also work. Conversion rates of ~98% seem possible. This looks the way to go for the amateur.

In other references I note that alkaline MnO2 slurries can become very viscous as the MnO2 particles swell to many times their original size. There was another patent that overcame this by adding the MnO2 in small quantities at ~15min intervals such that the MnO2 is converted as fast as it is added. In any case a good stirrer is certain to be needed.

No mention is made of the possibility of the KMnO4 crystallizing as the concentration increases, though I expect that would be inevitable, and may in fact be the preferred method of separation, or perhaps the way is to rapidly cool the solution at the end of the process to minimise hydrolysis, then filter ASAP. I have still to decide the best way of filtering the solution, given its extreme alkalinity and the oxidising power of the KMnO4. Suggestions on this potential problem would be most welcome.

Interestingly, it recommends that a small quantity of KMnO4 be added to improve initial electrode efficiency. This raises the possibility of reusing the spent solution with its remaining permanganate and xs KOH for another run. Simply add a suitable quantity of KOH and MnO2 and repeat.

I hope this is of interest.
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[*] posted on 22-8-2007 at 20:57


@ ciscosdad Very much so! I had been thinking and an MnO2 anode (somehow encased) in a diaphram-separated cell using fairly strong KOH. Several problems occur to me with this. The MnO2, although conductive, has resistivity 1-10 ohm-cm. Using graphite to decrease this is contraindicated. since KMnO4 will oxidize graphite when hot. Second, KMnO4 is very poorly soluble in KOH of any strength, due to common ion effect, so it might foul the anode. So, if it works using a slurry, two problems at least might be overcome.

A new problem is you'd probably have to stir it. Also, I think you still need a diaphram unless the KMnO4 falls to the bottom of the cell (and the MnO2 doesn't!) . Else the permaganate will be reduced at the cathode. A nickel anode also is easily available and non-exotic. Few people seem to realize that a nickel anode can be used to electrolyze water (with an iron cathode) and give hydrogen and oxygen without noticable oxidation of the anode, provided you use enough volts.

Can I get that patent on Wiki? Or can not_impotatnt help. He seems to read patents by the ton and know all the sources.

Best lead I've seen for ages!

Interesting about 'priming' with KMnO4 to get the thing started. I use it in chlorate production cells instead of dichromate. I've never really been sure it does much, though.

Regards,

Der Alte
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[*] posted on 22-8-2007 at 21:05


Might be able to use compressed air to keep the mixture agitated, and as a lift pump to move some of the mix into a settling tank about the electrolysis tank, with the slurry getting gravity return to the electrolysis tank and solution getting filtered, cooled to crystalise out KMnO4, and then returned.

You also might have to go with NaOH instead, adding KOH or K2CO3 to the clarified warm solution to get KMnO4 as a ppt.

One more high tech approach to the reduction problems would be to use a cathode that catalysed the reaction of H2 with O2, and/or was permeable to H2 or O2 with a stream of air down it's center. Convert the H2 back to water, keeping the hydroxide concentration relatively constant.


[Edited on 23-8-2007 by not_important]

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[*] posted on 22-8-2007 at 23:02
Google Patents


@Der Alte

I do all my patent searching at Google Patents.


http://www.google.com/patents

The actual patent:
http://www.google.com/patents?id=oFgrAAAAEBAJ&dq=3172830

Enjoy

BTW. The cell does not use a diaphragm, unless I have seriously misread it.




[Edited on 23-8-2007 by ciscosdad]

[Edited on 23-8-2007 by ciscosdad]
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