Sciencemadness Discussion Board


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DerAlte - 14-5-2007 at 20:12

I don’t know if there is any interest still in making KMnO4. There was a flurry of activity some time ago. I have developed a method that far excels the traditional KOH plus MnO2 plus (oxidizer +) heat method. Any interest? It is a wet method using easily obtained chemicals. The gist of it is to oxidize an Mn compound (several can be used) to the MnO4+ ion under somewhat carefully controlled conditions. It will take a fair time to explain, so I wanted to see first whether any interest was out there ,,, Der Alte

The_Davster - 14-5-2007 at 20:15

Sure, why not. We have a big thread on this somewhere, I'll merge this with it as soon as I find the other one.

DerAlte - 14-5-2007 at 20:28

I saw the previous thread, but some time back. The industrial method of heating MnO2 with KOH and using air as an oxidizer is doomed to failure. The industrial methods are all secret - but they carefully control temperature and blow air thru the melt. It is a fact that KMnO4 decomposes around 220 C or so. I've tried this, years ago (~50!) and you get a bit of manganate, enough to color the solution green. People do not realize that even a relative low solubility permanganate like KMnO4 when saturated produces a BLACK solution, only the meniscus being purple.

That being said, the amateur must seek other methods to get a reasonable yield. It ain't easy! But it can be done...


DerAlte - 14-5-2007 at 21:43

To make a permanganate, one needs a source of Mn, an oxidant, and a method. The process I am about to discuss is a wet method in H2O solution, at modest temperatures. The Mn may be either an Mn(II) salt, MnO2 (or any other oxide, except Mn2O7 of course! Mustn’t make it too easy). It is always convenient if the reagents are commonly available. So, you can use even MnO2 + crud from old batteries, IF you purify it with care. More anon.

For an oxidant, we have to look at the standard redox potentials of the conversion from Mn(II) or Mn(IV) to the oxidation state MnO4-. These are conveniently shown on WebElements (WE) for Mn and possible oxidants. Using these data, we determine that the oxidation is best carried out in alkaline solution from Mn++ (Mn(OH)2 in the WE figure – Mn will precipitate as hydroxide) at 0.34 volts. From MnO2 we need 0.60 volt. So we need an oxidant with a redox potential of > 0.34 Volt. (or 0.60v for MnO2). The oxidant chosen has to work in alkaline solution – pH TBD.
H2O2, which works like a charm for making chromates, is decomposed catalytically by MnO2 and manganous salts. Chlorates (see WE redox diagram) aren’t quite strong enough in alkaline solution. Forget K peroxydisulphate unless you have it on the shelf. Even then, it's marginal due to H2O2 formation.

Examination shows that hypochlorites might work, redox potential to Cl- of 0.89 in OH-. A bit close for MnO2, perhaps., so expect a slow reaction.

Let’s give it try…

To be continued…



MadHatter - 14-5-2007 at 22:20

Given the recent CPSC(assholes) victory against FireFox, I'm sure that any convenient
method will be appreciated. KMnO4 sales by pyro suppliers are now limited to 1 LB a
year. There's still some OTC sources for now but we don't know how long that'll last.

guy - 15-5-2007 at 00:42

are you suggesting using hypochlorite? Thats not really a good idea.

Zinc - 15-5-2007 at 03:17


Pyrovus - 15-5-2007 at 03:22

I've tried hypochlorites with limited success. I've found it to be very messy and annoying - no matter how much of an excess of hypochlorite I use, the reaction never seems to go to completion. In fact, despite using large amount of hypochlorite I believe that the permanganate solutions prepared this way are probably pretty dilute - after all, the colour of permanganate is pretty intense, yet the solutions weren't all that dark. Then of course, is the problem of purifying it - the biggest problem being the presence of chloride ions in the solution, as in the presence of permanaganate these are no mere 'spectator' ions, and I'm yet to find a viable way of getting rid of them easily.

However, if using hypochlorite as oxidant, use lots of base. The oxidation to permanganate for either Mn(II) or Mn(IV) involves the production of acid, and in acidic solution permanganate will oxidise chloride (which is present from the hypochlorite being reduced) to Cl2. Cl2, of course, is not fun, and in addition your (already probably low) yield will be decreased.

JohnWW - 15-5-2007 at 04:42

Exactly the same occurs, with lack of completion and difficulty in isolating the product as a solid, when one uses NaOCl to oxidize alkaline Fe(OH)3 or NaFeO2 to Na2FeO4, containing Fe(VI). The FeO4-- anion is intensely magenta in color, very similar to permanganate, and is used similarly in water disinfection. Theoretically, FeO4- and FeO4, containing Fe(VII) and (VIII), could also exist, probably being intensely blue or blue-green in color due to the charge-transfer absorption band being at an even longer wavelength, but I have not heard of them being prepared.

I understand that the usual modern industrial method of preparing Na and K permanganates(VII) and ferrates(VI), (and in similar ways chromates(VI), plumbates(IV), bismuthates(V), etc.) is by electrolysis of a cold strongly alkaline solution containing dissolved Mn2O3 or MnO2 and Fe(OH)3, as manganite(III) or manganite(VI) and ferrite(III), at suitable voltages. This avoids liberation of Cl- anions through hypochlorite decomposition, which would be oxidized by the desired products.

[Edited on by JohnWW]

DerAlte - 15-5-2007 at 07:16

Pyrovus, I agree that hypochlorites are messy! See continuation.

But it happens to be a very strong oxidant that is commonly available. The only other wet methods I have tried are electrolysis: ferromanganese (70% Mn) as anode in 30-40% KOH- works quite well, but where do you get the ferromanganese? I have also tried an anode of MnO2/graphite and you can get a pink color.


DerAlte - 15-5-2007 at 07:42

…try the following to show that it can be done. I’ll assume you have the reagents.

(1) Heat MnO2 with 50% NaOH or KOH in a test tube. No reaction is noted (some authorities claim insoluble manganites, xK2O.yMnO2 are produced). On Standing for some time a slight pink coloration may be noticed due to atmospheric O2.
(2) Instead of dissolving the NaOH in water, use household bleach as the solvent. This gives us a weak hypochlorite in about pH 14 solution. Heating produces the green magnate color, due to MnO4— ion.
(3) Repeat (2) using a concentrated solution of sodium carbonate in bleach. On heating the pink permanganate color appear. pH is of the order of 11.

Manganate is only stable in very alkaline solution. If a permanganate is heated with strong KOH, it turns into manganate plus MnO2

The essence of the method, then, is to oxidize MnO2 or Mn++ with hypochlorite at pH 11.

To be continued…


DerAlte - 15-5-2007 at 10:33

Lest you think I’m stalling, let me say that my intent is to give some interesting chemistry, not to show how to produce a ton of KMnO4. This process is NOT 100% efficient. Why any one would need a ton or even a kg escapes me, unless to indulge in dangerous pyrotechnomania or get the DEA very interested. Permanganates are rather unstable, especially with metal powders. Use perchlorates instead. I grew out of pyrotechnomania by the age of 18, managing to escape with no more than singed eyelashes and a few burns.

The main chemical usage of permanganates is in titration of reducing agents and in organic oxidations. It is possible to produce a working quantity – 20 to 50 gm – by using the potassium process to be described.

BUT FIRST, a caution. Strong hypochlorites are nasty. Always use rubber gloves. Read an MSDS on the web for all reagents or products mentioned if you are not sure about them.

If you indulge in chlorine production, remember that the stuff is noticeable at 10-100 ppb so don’t do it in an apartment, condo, row house or flat. If you do and let it leak you may be invaded by Big Brother’s Hazmat Boys (clothe them all in yellow, ho!). Chlorine production demands a shed or porch and good ventilation, an isolated plot and a brisk wind. Always absorb it rather than let it escape. You need good chemical apparatus for this.

Conc. HCl is used as a pool chemical and driveway cleaner and easily available. Be careful with it. Keep it a mile away from hypochlorites, permangantes and other oxidizers, and all metals. Read the MSDS for everything.

Finally, to end this boring preamble, some means of weighing to about 0.1 gm is necessary for success. You can’t just throw quantities of the reagents together and hope. Knowledge of volumetric analysis helps, too. If you’ve got a burette, use it! It helps to know the strength of the hypochlorite solution (iodometry) and also the permanganate (acidic ferrous sulphate or a sulphite). I need a nap…


DerAlte - 15-5-2007 at 20:50

The Process

Manganese carbonate is reacted with an alkali hypochlorite and an alkali metal carbonate in aqueous solution. The putative reaction is

2MnCO3 + 5KClO + K2CO3 --> 2KMnO4 + 5KCl + 3CO2

(An alternative is to use Mn(OH)2 in place of the carbonate. Haven’t tried it but should be effective too. MnO2 can be used but the reaction is much slower. For reasons to be stated, speed of reaction is vital.)

The KClO needs to be as strong as possible, circa 15%w/w. Now you’re not likely to have KClO but you can make it. More on making the components from easy available stuff later. Same for MnCO3, but it can be found on sites catering to pottery enthusiasts. Potassium carbonate can be obtained from sites providing material for amateur soap manufacture. Making it is not easy except from CO2 and KOH.

You can substitute Na for K but don’t expect to be able to crystallize out the NaMnO4 – it’s very soluble. It’s enough of a problem with the potassium salt, due to the large amounts of chloride produced. And it’s no better if you try to precipitate KMnO4 by adding KCl.

This is Important if you want any worthwhile yield.: Do the reaction in a glass beaker on a water bath and keep the temperature below 65 Deg C but above 50 deg. The reaction is slow and will take one to two hours at this temperature. DO NOT expect all the MnCO3 to react. Some will remain as indeterminate oxides and hydrated oxides of manganese, a brownish black crud. .

The solution should quickly turn red and then BLACK. Even saturated KMnO4 at 0 deg, containing 2,84% W/W is black and opaque, which complicates the separation by fractional crystallization.

When ready, dilute the solution with about 3X volume water and filter warm using a glass plug in a glass filter funnel. Do NOT Use Paper. KMnO4 attacks most organics, especially double bonds and –OH bonds.

That’s the bare bones. Enough for now! Next time, how to separate and get the permanganate and some notes on technique

DerAlte - 16-5-2007 at 20:21

It is very helpful to determine the amount of product you actually have. The solution should be black. If it’s pink or red, you failed. If it still stinks of hypochlorite, the reaction time was not long enough. Other products are KCl, unreacted K2CO3, and whatnot. We can ignore the carbonate because of its high solubility.

If you are competent at volumetric analysis, titrate against ferrous (better, ferrous ammonium) sulphate. Hypochlorite may still be present so you may overestimate the KMnO4. If you don’t have an accurate burette or can’t fabricate one, use a micropipette and count drips the way I do, using rather dilute solutions. Crude but effective! An accurate balance is essential for titration – I use an antique chemical balance that deflects on 1mg.

One thing you do know is the approximate amount of K+ ion you have in solution (assuming you know the strength of the KClO solution used – see later.). This helps to reduce the solubility of the KMnO4 considerably, by the common ion effect.

At 0 C the solubility of KMnO4 is about 2,8% w/w aq. That of KCl, 28%. Without doing a common ion calculation, it can be safely assumed that in saturated KCl solution, KMnO4 solubility will be reduced well below 2%. So, cool in an ice bath Preferably in a refrigerator to delay the melting. You will expect to get a blackish mess containing permanganate and chloride in all probability. Filter this off, keeping the filter cold, using a glass wool plug in a filter, a glass mat filter or a glass frit filter but NOT Paper. Wash with a small amount of ice cold water.

Now is the time to reveal the real problem. I’m sure many have already seen it – THE HYPOCHLORITE DISPROPORTIONATES:-

3ClO- --> ClO3- + 2Cl-

So we also have to deal with an undetermined amount of chlorate. And here is a HUGE snag; it happens that KClO3 has almost an identical solubility curve as the permanganate and it’s next to impossible to separate the two. It would take the patience of Job.

I can’t think of any chemical means. Permanganate is the stronger oxidizer. I cheated – I didn’t promise you’d make Solid or pure potassium permanganate, merely permanganate. You got it, in solution. Now think of a solution. We could use the sodium salts instead of potassium …

2B continued…


Filemon - 17-5-2007 at 04:17

I have synthesized NaMnO4 with 2MnCl2 + 5NaClO (aprox.15%) + 6NaOH. The reaction is very quick. Surprisingly the NaMnO4 had precipitated. For the ions Na+ of the NaCl?

woelen - 17-5-2007 at 11:45

Are you sure that it is not NaCl with some NaMnO4 mixed in? Permanganate is so dark, even a 5% mix with a colorless salt still is very dark, near black.

DerAlte - 17-5-2007 at 13:34

woelen, Agree. Saturated soln of KMn04 @ O C is still very nearly black at 2.8%. MnO4- is about the most intensely colored ion in existence.

Filemon, I doubt the reaction is "quick". Sure, a pink color appears but this only means very, very dilute. Also, NaMno4 is exremely soluble like sugar in a rainstorm, at about 200gm/100 gm aq. What you probably saw was a mixture of oxides and hyroxides of manganese. Mn is so avid for O that even chlorides turn brown on being kept in air. Anhydrous MnCl2 can only be prepared in an atmosphere such as gaseous HCL

And, under pH 14 or so in strong alkali, green manganate shoulf be produced.

MnCl2 crystals, as hydrate, (can't remember which - ), turns acidic when kept. Strong hypochlorite acting directly on such will produce chlorine, so take care. Addition of NaOH stops this by neuralizing the acid. Same is true off all Mn(II) salts, as regars oxidation. They all tun brown in time.

As for using Na instead of K salts in the method I outlined, guess what? The solubility of NaClO3 is of the same order as NaMnO4. Instead of being fairly insoluble it's very soluble.

The only possible way to separate the permangante from the chlorate that I can think of is via the silver salts. Anyone got an ingot of silver and a gallon of nitric acid?

Regards, To Be Contd,.


Filemon - 17-5-2007 at 15:39

Originally posted by woelen
Are you sure that it is not NaCl with some NaMnO4 mixed in? Permanganate is so dark, even a 5% mix with a colorless salt still is very dark, near black.

Probably it's blended with NaCl.

G.i.B. - 17-5-2007 at 15:42

Can you distill the KMnO4 from the mix at reduced pressure ?

UnintentionalChaos - 17-5-2007 at 15:54

That is like asking if you can distill salt out of salt water.

G.i.B. - 17-5-2007 at 15:58

I read somewhere, that you can distill solids from solids, I am not sure how.

I off course meant, first boil away all the water.

[Edited on 17-5-2007 by G.i.B.]

[Edited on 17-5-2007 by G.i.B.]

not_important - 17-5-2007 at 20:28

Originally posted by G.i.B.
I read somewhere, that you can distill solids from solids, I am not sure how.

True, for example you can distill solid CO2 away from sand. So long as the vapour pressure is high enough at a temperature below the decomposition point, you can distill a substance be it solid or liquid. I doubt that those requirements hold true for permanganates and chlorates.

DerAlte - 17-5-2007 at 21:38


I am sure you all think I’m a lunatic by now! Back to disproportionation. Unless I’ve made a gross mistake in calculation , the reaction

3OCl- --> 2Cl- + ClO3-
is only a bit less energetic (SEP ~0.4v) as the oxidation of Mn++ we are trying to achieve (~0.5V). It is said that this disproportionation becomes “significant” at 40 C. As we are using a fairly concentrated hypochlorite solution and a higher temp (55C) – on the premise of speeding up the wanted reaction rate – this is even more likely to be true. I estimate that about 30-40% of the last product I got was chlorate after several fractional crystallizations.

As for using silver salts, mad as it may seem, it should be feasible. A rather expensive idea, perhaps! (Recycle!) I don’t have any Ag salts on hand so haven’t tried it. If you happen to have a bucket of AgNO3 and an adventurous spirit, first do yourself a favor by ensuring all the chloride is gone from the product so you only have chlorate and permanganate, by repeated recrystallizations. Otherwise all you’ll get is AgCl pptd.. AgMnO4 has solubility 0.9g/100g at room temp; AgClO3 something like 18g/100g. AgNO3 is over 200g/100g. Easy separation.

Finally, react the pptd silver manganate with KCl solution to get inslouble AgCl and the desired KMnO4 in solution

If anyone is still interested, I shall expound my CRUD (Chemical Reagent from Utter Dross) method of getting a pure manganese product from old alkaline cells, and also how to make 15% plus NaOCl or KOCl from crap used for pools.

I used, in another place at another time, to always use reagent grade chemicals; but Big Brother and his minions, the Hazmat police, and those who believe that owning an Erlenmeyer flask is the prelude to being a drug czar, have made an experimenter’s life difficult (albeit from sound motives combined with political correctness) to say nothing of the affront to my libertarian principles. So I am forced to turn crap into at least technical grade reagent like the rest of you. And that can be fun, too…


DerAlte - 19-5-2007 at 09:27

How to prepare a strong solution of Hypochlorite.

After writing this offline I discovered that Garage Chemist had already posted an excellent dissertation on this under a new NaOCl thread. Please read it. I agree with all he says.

This may be a bit redundant, in consequence, but having written it, I might as well post.....

You can get up to 15% NaOCl from pool supply companies (HTH). It will not keep well except in cold weather when you don’t want it for pools. If you want any other hypochlorite, such as KOCl you’ll have to make it. You can absorb Cl2 gas in KOH or K2CO3 or use the CRUD method:

Calcium hypochlorite is a relatively stable solid form. Get some ‘pool shock treatment HTH’ type. This contains about 60-65% by weight Ca(OCl)2. The rest is calcium hydroxide, chloride, chlorate, etc., all Ca compounds. Read the MSDS for this stuff, it’s nasty. React with sodium carbonate in the correct proportion, assuming that the product is 100% Ca(OCl)2- this is close enough for government work. First mix the two as solids, then add water sufficient to produce a 15% solution of NaOCl. (Remember the sodium carbonate may have up to 10H2O as hydrate – Dehydrate first at 150C in an oven to Na2CO3.) Potassium carbonate is very soluble but sodium carbonate is most soluble near room temp.

You’ll now have a loathsome white mess! Calcium carbonate is precipitated as a very fine powder. Stir it up well and let it stand for a couple of hours in the refrigerator at, say 10C. When settled, first add a few drip of carbonate solution to make sure all Ca++ ions are gone. Then filter over a two stage filter of several fine weave pieces of glass cloth backed up by a glass wool plug in a glass filter funnel. Paper will be attacked by the hypochlorite. Keep well stopped at a low temp. Color should be yellow to greenish.

This strength should be used rapidly. Even a week causes marked deterioration in OCl- content. Adding a little NaOH helps to keep it by buffering the pH.


DerAlte - 19-5-2007 at 09:37

MANGANESE SALTS FROM used batteries: an old favorite. Some are under the delusion that the black mess is manganese dioxide. It does contain some, true. But it also contains zinc oxide, lower oxides/hydroxides of Mn, (NH4Cl and Zn(NH3)2Cl2 and similar Zn complexes in the case of the old Zn/C type), or KOH in alkaline plus potassium zincate and grapitic carbon dust. My CRUD process:

First clean away the solubles by boiling with water for ½ hr. This also breaks up lumps. Next wash in DILUTE COLD HCL or dilute acetic acid (5%)to remove Zinc compounds like ZnO.

You could then try floating off the graphite with a stream of water but that doesn’t work very well. Or you can heat in an iron can to bright red heat (500C+) to burn off the C and reduce the oxides to Mn2O3. Of course you destroy any dioxide present.

IN my CRUD process I just leave it. The next step should only be undertaken if you know and have the apparatus to deal with Chlorine properly. Absorb it in NaOH, KOH, Na2CO3 or K2CO3 solution.

The mixture of oxides will not contain MnO, but contain MnO2, Mn2O3 and possibly Mn3O4 (And C). Dry off at 200C in an oven. React with strong HCl (>30%) in a flask (Round is best, only 1/5 full to allow frothing room). Warning: You now have a chlorine factory. Heat gently on a sand bath or gauze until reaction ceases. The stuff will not clear due to the presence of graphite but has to be judged by cessation of Cl2 emission. I use a second tube in the gas generator flask to blow out the remaining chlorine (Not by mouth!)

Contents of the flask is now MnCl2, excess HCL (and C). Filter off when cool (paper is OK) to remove the C. A fine pore paper is essential: coffee filters let a lot of fine graphite through. You can crystallize the MnCl2.4H20 as nice pink crystals from a solution acidified by HCl. Or, as I do, convert to carbonate, an insoluble pink powder that keeps well, is not deliquescent, and can be used to make any Mn(II) salt.

Qualitative analysis shows that the salt from alkaline cells, the most common these days, is quite pure. That from Zn/C cells seems to have a slight iron impurity.

Finally, if what you want is MnO2, you can regenerate it from the chloride (or carbonate) by precipitating the Mn(II) hydroxide with NaOH, washing it, and adding strong NaOCl and warming. The hydroxide is creamy colored, carbonate whitish pink, and MnO2 black as the ace of spades. Mn compounds have many colors, red, purple, green, blue, brown, black – all except yellow. At least, can’t think of one


Eclectic - 19-5-2007 at 10:21

Does fusion of MnO2 with NaNO3 work?

[Edited on 5-19-2007 by Eclectic]


DerAlte - 19-5-2007 at 11:02

The last time I tried any fusion reactions was about 25 yrs ago, so no guarantees from the management.

It does, according to some scrappy notes I have from that era, but you have to add NaOH. So does KClO3. You get the green mangante in both cases. I suspect KClO4 would also but didn't have any available then. The trouble always is preserving the manganate. Permanganate and manganate decompose at between 200-250 C yet the fusion temps are all higher than this. I don't think you'd get anything without the alkali.

Another oddity is that the probable product of the nitrate oxidation is to nitrite, a reducing agent.

DerAlte - 19-5-2007 at 11:52

Finally, I’ll throw in a method I tried earlier. It differs little from the above and still suffers from the same permanganate/chlorate problem. If anyone can find a solution to this problem, please post it! The method of using Ag is obviously impractical for nearly everyone.

Basically, you mix stoichimetric proportions DRY: 4 moles MnCO3, 2 moles Ca(ClO)2, 1 mole Na2CO3 and hope that

4MnCO3 + 3Ca(ClO)2 + Na2CO3 --> 2NaMnO4 + 3CaCl2 +Mn2O3 +5CO2

(Note: anyone can write an equation like this! The laws of physics determine whether it works. There is no doubt that all the Mn carbonate does not react and a black mess is also left, which may be Mn2O3 or MnO2. Also, the calcium is precipitated as carbonate, which makes it gray. K can replace Na, of course)

Add water and warm to 40-50C in a beaker on a water bath for a fair time, stirring frequently..

Chlorate is produced again, of course. It only avoids the messy step of actually making the NaOCl.

NOTE:- Further research shows that using Lithium instead of K or Na does allow a separation : the permanganate is very soluble at about 70g/100g aq but the chlorate is extremely soluble at about 400g/100g aq. Not too easy but practical, but most of us don’t have lithium (pyro enthusiasts might! I’ve got a few gms carbonate)

I do hope I did not disappoint too many. If there was an easy cheap method it would be used industrially, you bet. Electrolysis in KOH or K2CO3 on to a MnO2 anode should work but making a conductive one and holding it together is a challenge.

A tit bit for Mn enthusiasts. Have you ever seen a hypomanganate like K3MnO4? Azure blue in color. Very careful reduction of pure KMnO4 with sodium sulphite at around 0C in a dilute solution will take you through the steps MnO4- (Red/purple) --> MnO4-- (deep green) --> MnO4--- (sky blue) --> MnO2 (black ppt)


Eclectic - 19-5-2007 at 13:31

OK, so a nitrate/nitrite mixed melt with some -OH melting at about 125 C should work nicely, followed by electrolysis to bring up the oxidation state? Or disproportionation of the manganate?

[Edited on 5-19-2007 by Eclectic]

The_Davster - 19-5-2007 at 14:10

Simple acidification of the manganate, even using CO2, will work causing the disporportionation.

Eclectic - 19-5-2007 at 14:36

But I LIKE electricity! :D

I like the acidification with CO2 though. For making a strong NaMnO4 solution to prep a wide variety of permanganate salts, you could drive the excess Na out of solution as bicarbonate, and get rid of the very soluble nitrite at the same time. Excess NaNO3 should crystallize out by common ion effect, leaving fairly pure NaMnO4 solution.

Agricultural MnSO4 + NaHCO3 + agitation/air --> MnCO3
MnCO3+2NH4NO3 +heat --> Mn(NO3)2 + NH3 + NH5CO3 (volatile)

(If you try to crystalize out the Mn(NO3)2 from solution you get a double salt with some ammonium, not really a problem though if you are going to add NaOH and heat. The whole mess should dissolve in it's own water of crystallization below 100 C.)

[Edited on 5-19-2007 by Eclectic]

It looks like LiOH would work even better for making an extremely concentrated permanganate solution (70%+), and the carbonate has low solubility. :D

[Edited on 5-20-2007 by Eclectic]

DerAlte - 20-5-2007 at 12:05

I Like electrolysis very much too, Eclectic! I guess everyone who has done the famous chlorate process also does..

Anodic oxidation of ferro manganese in KOH or K2CO3, gives manganate or permanganate (plus some MnO2, I believe). I tried it with an old piece of spiegeleisen (bright silvery stuff, Fe80/Mn20) a long time ago and that didn't work very well, if memory is OK.

Ferromanganese is an alloy 20Fe/80Mn which should be available on the web, as steelworks make it by the tons, but a Google serch revealed no sources. Mn metal is expensive but you can find it, but ferromanganese ought to be cheap. Anyone know a source by the kg instead of the ton?

I was hoping, by opening this thread, to see if some bright young spark had any radically new ideas on MnO4- production. This tired old brain is now out of ideas!


Eclectic - 20-5-2007 at 12:52

If you are determined to play around with metallic manganese, you could probably electrowin your own from MnSO4 with an excess of MnCO3 on the bottom of the plating tank to neutralize acidity.

55lbs of agricultural MnSO4 can be had for $30-40, $2-3/lb in small quantities.

I wonder what you would get running electricity through a Mn(NO3)2 and NaNO3 solution?

DerAlte - 5-6-2007 at 19:41


I am still messing around with KMnO4 synthesis. I noticed an old thread was resuscitated a day or so ago. It’s a hardy perennial!

I have a question that has been bugging me for a few days. Does anyone have any reputable NUMBERS for the solubility of NaMnO4 – repeat SODIUM permanganate? I’ve done a CRC and Googled it to hell and gone, and only land up with useless statements like “very soluble” or “extremely soluble” – hell, I know that! I have an old notebook from a past era that states 220g/100g aq at 25C. No reference – I’m bad at that. If anyone has any numbers, please post…



not_important - 6-6-2007 at 04:38

The notes I have are similar, only numbers are " greater than 300 g/100 cc" I suspect it is difficult to get a pure dry sample - "the sodium salt is so soluble and hygroscopic that its solubility is too difficult to be measured accurately". Note that it is commonly sold as a 40% by weight solution, which is only 400 g/l

DerAlte - 6-6-2007 at 20:35

Thanks, not_important! It confirms what I have, near enough for government work.

The reason for asking was because I have been looking at various permanganates and chlorates, to see what ratio of solubilities exist to differentially separate the chlorate produced by the hypochlorite -> chlorate degradation ( disproportionation) inevitable in the process I have been suggesting.

I find the following ratios of solubility, permangante to chlorate - at approx RT -

K 9:10.1: Ba 65:34; Ca 338:209; Na ~260:96; Li 372:71; Ag 15:0.9 g/100g aq

Well, nobody is going to waste silver on a permanganate synthesis, nor are they likely to waste Rb/Cs, which also have rather insoluble permangantes, less than K.

Li looks a promising candidate, but that also sounds a bit too rare (not as rare as it used to be in my youth) so the next best bet is sodium.

I have just performed an all sodium reaction of MnCO3. NaOCl, Na2CO3 as outlined earlier in this thread. So I now have a (well filtered) mixture of NaClO3, NaCl, NaMnO4 plus possibly some Na2CO3. I have given it a brisk boiling to convert all NaClO to NaClO3. At least there is no hypochlorite smell (but that is muted in alkaline solution - it needs actual HClO, as in acid solution, to really give that bleach smell).

I have evaporated the solution to precipitate that Na Cl, the major contaminant. Twice, right down to a fraction of the original volume. One more go has also precipitated considerable amounts of the Na ClO3 if the numbers above are correct. I should have the NaMnO4 as the major component of my solution, which is deep black and impentrable by a focussed flashlight. I shall add KCl tomorrow and - we'll see!

The results of an earlier all K run seemed to produce chlorate and permanganate almost equally. I have not estimate the amount of either except by viewing the crystals under a microscope - needles of the permanganate are interspersed with plates of chlorate.



JohnWW - 7-6-2007 at 01:38

As a cheaper substitute for permanganates, for industrial technical purposes such as water disinfection, how about making ferrates(VI) instead, in much the same ways, although preferably by electrolysis of a cold alkaline solution of Na or K ferrite(III)?

BTW An "holy grail" of iron chemistry would be to try to make, under more extreme such conditions, perferrates(VII) and FeO4. A theoretical study I have read somewhere suggests that they could exist. I wonder if anyone has succeeded.

not_important - 7-6-2007 at 05:07

Theory on FeO4(-) : Inorg. Chem. (1999) 38(22): 4942–4948

DerAlte - 10-6-2007 at 23:16


This will be my last contribution to this thread, unless there is something worth commenting on. The results of attempting via the all Na route still resulted in significant chlorate in the product on adding KCl and messing around with the very concentrated solutions was a pain.

I tried a one pot last effort, using MnCO3, calcium hypochlorite (HTH approximation as above) plus the calculated amount of sodium carbonate to precipitate the calcium AND satisfy the stoichiometric equation given earlier. This seemed as efficient as bothering to make sodium hypochlorite separately. The solid components I ground together in a mortar before adding water. Heated at about 55-60C for two hours followed by a 3/4 hour boiling to disproportionate remaining hypochlorite to chloride and chlorate. Filtered of the calcium carbonate and unreacted Mn products using glass filter. A titration of diluted solution (acidified by HCl) against ferrous sulphate gave the MnO4- in the product as 39% of the ideal yield.

Then reduced to half volume to precipitate NaCl, refiltered, and repeated to precipitate sodium chlorate. Very tedious! Adding a saturated solution of KCl actually gave a product that looked like KMnO4 but with very small elongated crystals when cooled to 0C for a while…

Haven’t yet got to drying but a small dried portion acted like KMnO4 when mixed with carbon and ignited. But then again, chlorate might do the same except this crackled in a way more reminiscent of permanganate. There’s still chlorate there, but how much I haven’t attempted to find out.

If 39% of the hypochlorite went to making MnO4- and the rest disproportionated, the ratio of MnO4- ion to ClO3- ion would be 3.9: 3.7 I calculate. That is, prior to attempts to separate. I think this effort to make permanganates has left me exhausted!



[Edited on 12-6-2007 by DerAlte]

JohnWW - 11-6-2007 at 14:16

With regard to my post above about ferrate(VI), FeO4(2-) and the possible Fe(VII) and Fe(VIII) oxidation states of Fe as FeO4(-) and FeO4, and possible uses as a cheaper substitute for permanganate, Solo has come though for me with the following post, with a PDF article, on a different thread in the References section at: :

posted on 12-6-07 at 08:19

Theoretical Studies on the Higher Oxidation States of Iron
M. Atanasov
Inorg. Chem. 38(22): 4942–4948 (1999)

Density functional theory (DFT) and multiconfiguration self-consistent field (MCSCF) calculations on the oxo FeO42- (FeVI) and the hypothetical oxo FeO4- (FeVII), and FeO4 (FeVIII) and peroxo FeO2(O-O)z [z ) -2 (FeIV),z ) -1 (FeV), z ) 0 (FeVI)], Fe(O-O)2z [z ) -2 (FeII), z ) -1 (FeIII), z ) 0 (FeIV)], and FeO(O-O)2z [z ) -2 (FeIV), z ) -1 (FeV), z ) 0 (FeVI)] clusters are presented and discussed. The results show the potential of stabilizing FeVII and FeVIII in tetrahedral oxo coordination. On the basis of absolute electronegativities calculated using DFT, it is predicted that FeO4 will be rather oxidizing, even stronger than Cl2 and O2. On the basis of a comparison between total bonding energies of M1M2FeVIO4 (M1, M2 ) Li, K), MFeVIIO4 (M ) Li, K), and FeVIO4 clusters, possible synthetic routes for electrochemical preparation of FeO4- and FeO4 species are discussed.

Attachment: Theoretical Studies on the Higher Oxidation States of Iron .pdf (141.03 KiB)
This file has been downloaded 1 times

[Edited on 12-6-07 by JohnWW]

12AX7 - 11-6-2007 at 17:53

I would think the Cs or Ba salt would be most stable; maybe cold anodic oxidation of a suspension of BaFeO4?

Hmm, how does the solubility of Cs2FeO4 compare to the K salt...


DerAlte - 12-6-2007 at 19:05

Above, On 10-6-07 I wrote:

"A titration of diluted solution (acidified by HCl) against ferrous sulphate gave the MnO4- in the product as 39% of the ideal yield."

Looking over my calculations, I find I omitted the 7H2O of the ferrous sulphate.

The corrected figure is 21% yield based on Mn, pretty poor. 39% didn't sound too bad! This accounts for the difficulty of separation.


ciscosdad - 12-6-2007 at 20:59

Fascinating Thread DerAlte.

What about some other means of reducing the xs ClO that will not produce ClO3?
IIRC Urea will reduce it forming N2 and CO2.
If the Urea does not interfere with the MnO4 that may work. It also has the advantage of not introducing any extra ion species to the final mix (with care).
Can anyone else think of some selective reducing agent should the Urea be unworkable?

On further reading about Permanganates, it seems very unlikely that any reducing agent will be unreactive with it.

[Edited on 13-6-2007 by ciscosdad]

[Edited on 13-6-2007 by ciscosdad]

DerAlte - 13-6-2007 at 18:42

Ciscosdad, thanks for the compliment. But I suspect that KMnO4 would oxidize urea all too easily.

You raise an interesting point. In acidic solution both permanganate ion and hypochlorite are much stronger oxidants than in alkaline solution, but still almost equal in redox potential. In alkaline solution, hypochlorite is marginally the better oxidant – hence the difficulty in carrying out the suggested synthesis rapidly and efficiently. Hypochlorite would just as easily oxidize itself to chlorate as do what we want.



ciscosdad - 14-6-2007 at 15:40

Interesting Patent.

Not a method to be attempted in the backyard I guess

I'm still looking for details of the the method that Condy used in the last part of the 1800's. The encyclopedia references I've seen say that the method was simple enough that he had problems with patent infringers. There's got to be some clues there.

I wonder if a pot of 40%KOH/MnO2 kept at ~200 DegC and with air bubbled through for a few days would be possible?
If the pot was a tallish tube, well insulated and heated by resistance wire, the heating costs should not be excessive.
The fumes/mist from the hot KOH would be a problem, but not insurmountable. Evaporation could be checked on by inspection and (careful ) topping up

Just mental doodling. I need to think more on this.
I also need to check if urea solution decolourises KMnO4.

12AX7 - 14-6-2007 at 16:14

As I recall, I once tried to oxidize urea with permanganate, getting nowhere, having let it sit for a week or two at room temperature. I don't remember if I did anything to the pH, but it shouldn't need any if it's as good an oxidizer as it claims to be!


DerAlte - 14-6-2007 at 23:44

ciscosdad , I believe Condy used the time honored fusion with KOH. and oxidation by air. Apparently even the alchemists knew of the process.



Condys process

ciscosdad - 17-6-2007 at 15:27

You are almost certainly right about the basic method, but as they say, the devil is in the details. There should be a patent listed where he specifies exactly what he does to get from Pyrolusite to the solid crystals. Buggered if I can find it though.

One further thought re Urea: It will be of llimited use even if Permanganate does not oxidise it as there will likely be a significant amount of ClO3 formed during the initial synthesis (ClO + heat). So we're still stuck with the main impurity.

I still like the idea of the steel reactor with liquid KOH and MnO2 +O2. Next is a boiling point/concentration graph for KOH to see what temps will work.
More doodling!

DerAlte - 17-6-2007 at 20:27

ciscosdad, you have it. We do not want to reduce the hypochlorite but the chlorate. And chlorate is a weaker oxidant than hypochlorite.

I have been doing some thinking and have come up with a method to somewhat ameliorate the situation. I will hold my water on that one until I have completed my latest experiment.

If any are curious as to how I manage to do these somewhat lengthy tria runs, the answer is I am retired and have been for years (and years and years!) I happen to be alone at present, which makes it even easier... and chemistry is still fun, even at my age.

Notice that (provided you start with potassium hydroxide and DON"T use chlorate as an oxidant) the traditional fusion does not give any products that interfere with the production of the manganate. Not so with the wet methods here. Chlorate production is inevitable and must be minimized. Other oxidants such as persulphates might work but it seems pointless to use a more expensive one to make a permanganate.

The other method that I would like to try is electrolysis in a divided cell using a MnO2 anode. I have tried this in a hit and miss fashion and you do get the purple coloration. The problem is making a suitably conductive anode. Any ideas, anyone? The batteries using MNO2 use graphite. Electrolysis is a favorite of mine because almost anyone with a bit of knowledge can do it, and it allows us to achieve 'reactions' that otherwise require high energy reactants. Just what the amateur requires!



12AX7 - 17-6-2007 at 21:10

How about some anode in KOH with a suspension of MnO2?

Would graphite break down (I know it makes a mess in sulfuric acid), needing say, PbO2? (Hmm, but that would want to try to make K2PbO3, at least without voltage...)


DerAlte - 17-6-2007 at 22:02

You need the KOH for sure. Conc. to induce the KMnO4 to precipitate off the anode (common ion effect) But just suspending the particles - I suppose the OH- ions near the anode might just get them. As for graphite, does H2SO4 make a mess of them or only when it's really concentrated? PbO2 seems to have a cult following, 12AX7. Unike most, I have made perchlorate sucessfuly with even graphite. You just have to tolerate the anode erosion and clean up the mess. If it's inefficient, what does it matter? Carbon and electricity are cheap compared with the prices one sees chaged, if you can get it al all.

Der Alte

12AX7 - 18-6-2007 at 13:41

I believe graphite forms an oxide (intercalated or something) when subject to anodic conditions. As I recall, it works with any sulfate, and probably nitrate as well. Seems to me the only way you can possibly get perchlorate with graphite is by cheating the reaction with high voltage and current density pushing past the erosion regardless.

Speaking of graphite oxide, a lot of times I've had the graphite sludge from my chlorate cell rise to the top due to adherent oxygen. Now, I would ordinarily attribute this to hypochlorite decomposing slowly, but that only works when the smell of chlorine is strong. Sometimes it happens to low-hypochlorite solutions. Graphite oxide as the erosion product, with a high oxidation potential (above chlorate, but below perchlorate, persulfate, etc.), would seem to make sense, and if it's decomposing in suspension, that would explain the adherent oxygen bubbles.

Anyway, applying to this thread, you have to determine if the oxidation of whatever mechanism operates -- direct oxidation of manganite, production of intermediate peroxide or superoxide, etc. -- if it's lower than graphite's erosion potential.


DerAlte - 18-6-2007 at 21:19

12AX7, both boiling nitric acid and I think I read somewhere, permanganates, will oxidize carbon to mellitic acid, which is benzenehexacarboxylic acid - I'd draw it if I knew how to in this format. The anhydride of this is mellitic anhydride, C12O9, another weird 'oxide' of carbon.

Sure, you use a high voltage and high current density to make the perchlorate. The rods get eaten away like there's no tomorrow! To make the chlorate, it's essential to make sure anode and cathode products mix - the action is chemical as well as electrolytic; for the perchlorate, this is not as essential although hypochlorite, still produced, is capable of converting chlorate to perchlorate. I use different electrode configurations (Fe cathode). Carbon is the sacrificial lamb. Current efficiency be damned! You are spending more effort to make your PbO2 anodes than it's worth (but I'd love to try one - or Pt!)

With regard to permanganates, I'm working on it (still!)



12AX7 - 19-6-2007 at 16:10

Mellitic acid is interesting stuff. I bet its first proton or two is strongly acidic (much as EDTA's first has a pKa in the unity range). And yes, not to mention the anhydride being an aromatic oxide. Apparently it occurs naturally as the aluminum salt (mellite).

But I digress... maybe titanium doped with sun-dried shyte would work as an anode. :P (Hey, don't poo-poo this idea -- excrement contains trace elements!)


ciscosdad - 19-6-2007 at 17:33

Patent 3652417 (US) is interesting reading.
It refers to Alkaline Permanganate solutions used for descaling. It uses anionic surfactants to decrease the self destruction of the KMnO4 (as I understand it). It may be applicable to synthesis if it shifts the equbilibrium appreciably.

I have found the boiling point of KOH 40% to be ~132 degC. Has anyone seen a table / graph for higher concentrations? I'm thinking of aerated digestion of KOH and I expect the temp will need to be significantly over 132 DegC.
Perhaps a Eutectic mix of KOH/NaOH?

More searching.

ciscosdad - 19-6-2007 at 18:38

More searching has turned up Patent No 7 056 424 (US)
This refers to the manufacture of MnO2 electrodes on a stainless steel substrate by initially applying a porous non conducting layer, which may be a fabric, or some deposited porous ceramic or plastic.
What about a porous pot filled with Mercury! I love Mercury.:P
It might make the Electrochemical method a lot more appealing. The key to the method appears to be the electrodes.

I wish they would not use the convoluted legalise in these
things. Just getting my head around the chemistry is bad

Patent No 3986 941 (US)

The process is for the electrolytic production of KMnO4 directly from a KOH and MnO2 slurry. (or NaMnO4)
Electrodes can be Stainless Steel (amongst other things)!

20%KOH (aq) at 80 - 90 Deg C. Much more friendly than the process I was thinking of.
Process time is less than a day assuming reasonable currents.

The patent even describes a run using a glass 1 litre beaker :D

[Edited on 20-6-2007 by ciscosdad]

DerAlte - 20-6-2007 at 08:23

12AX7, your anode ideas, although a shade far fetched, may indeed have merit! Is not PBO2 a pleasing velvety brown color? But, I fear, sun-baked shyte may have poor conduction characteristics. Liberal addition of titanium powder may help this. I am sure its oxygen overvoltage is suitably high!

Talking of sun-baked shyte reminds me of the layer of crud on the bottom of my graphite perchlorate cells – but there it is the pulverized remains of many a carbon rod that gave it’s life for the cause…

Ciscosdad, nice thinking and research. Damn, before this thread dies of inanition, we’ll have a method, by hook or by crook. A clever and persistent chemist can do anything… Sorry no data on KOH or NaOH/KOH mixtures

When I have time I’ll report on my latest effort and suggest a few ideas – at present I’m midway through testing…



ciscosdad - 20-6-2007 at 15:47

For those of you interested, the Patent Search site I'm using is:

The patent Number I'm currently obsessing about is 3 986 941.

As I read it, the procedure can be:
In a (say) 2 litre steel container....
Make a slurry of MnO2 in 20% KOH solution so that there is 4.2 moles of KOH per Mole of MnO2. Add a catalytic quantity of KMnO4 to improve current efficiency in the early stages (recommended but not essential as I understand it). Heat to 80 to 90 Deg C. Insert stainless steel elctrodes and pass 10A through the solution for ~18hrs while maintaining the temp.
Conversion is of the order of 98 - 99 %.

Little is said about subsequent extraction of the KMnO4 (the solution was apparently merely analysed to determine conversion ratios).
Perhaps it can be as simple as hot filtration followed by crystallization of the KMnO4. A quick bit of mental arithmetic implies a concentration of ~150g of KMnO4 in the approx 1Litre of solution.

To be determined:
1......Filtration and cooling will certainly recrystallize the KMnO4 , but is it soluble enough at 80 - 90 deg C to allow uncomplicated filtration of the xs MnO2?
2......Does the solution need to be neutralized or acidified to work up?
3......Will the electrodes give problems with the MnO2 deposits mentioned elsewhere? Ref Patent No 7 056 424
4.......Current density required?
5.......Will agitaiton be necessary?

There is the possibility that the mother liquor (after KMnO4 recrystallization ) can be reused as is for the next cycle. Simply add approx 1 mole each of MnO2 and KOH and repeat. It even has some KMnO4 left in it to help start the process.

If you guys have read the patent, this post is probably unnecessary, but I'm trying the get the small scale / improvised procedure straight in my own mind. Have I missed anything?

Ref Item #4

Current density is quoted as 5 - 50 mA/cm2 (converted)
Current Concentration is 3 - 30 A/Litre

At 10A the area of the anode is: 250 - 2500 cm2.
I assume the iron rod (or whatever ) cathode area is not an issue, but the problem would be easily rectified by using another cylindrical electrode.

I'm hoping that the resistive heating of the cell will supply a significant amount of the heat required to maintain the required 90 Deg C. Insulation and an idling hotplate under the reactor should do the job.

[Edited on 21-6-2007 by ciscosdad]

DerAlte - 20-6-2007 at 15:52

A few thoughts on the production of KMnO4

The paramount thing is to minimize the production of chlorate during the process. So I have been racking my brains as to how best do this. I was looking at the possible kinetics of the process when it struck me that the equation I had given supra was in fact a combination of a fast reaction and a slow one.

2MnCO3 + 5KClO + K2CO3 --> 2KMnO4 + 5KCl + 3CO2

On addition of the hypochlorite, the MnCO3 quickly turns (brownish) black due to (hydrated) MnO2 production:

MnCO3 + KClO -> MnO2 + KCl + CO2

This is followed by oxidation of Mn(IV) to Mn(VII) as permanganate, a very slow process as noted before. Hence the driving potential is quite small, not the approx 0.34 volt redox potential from Mn++ to MnO4- but rather the 0.6 volt needed to convert MnO2 to Mn-. (The hypochlorite provides about -0.89 volt to achieve this)
(Note: none of these values apply to the conc. solutions we are using. I am too lazy to apply the Nernst equation to get realistic values)

So there is actually no point in using the carbonate unless you have it. You can start with the chloride directly from the manganes purification (but neatralize to acid to prevebt chlorine production first), or any Mn(II) you happen to have.

So, lets call it Mn++ + ClO- + 2e --> MnO2 + Cl- to keep it general.

The great thing about this is we can separate the MnO2 as a ppt. and use it for the following step. This has three advantages: (1) It avoids the addition of unnecessary chloride ions because the dioxide is precipitated and can be separated and dried; (2) if you have MnO2 that is relatively pure to start with, use that. (3) you can use cheap old bleach, NaOCl, for this step and not care about the extra load of NaCl it will introduce.

The MnO2 can be dried and weighed. It is likely hydrated to some extent as X(MnO2), Y(H2O). Dry at about 200C to minimize hydration, or even higher.

The (very!) slow reaction is then
2MnO2 + 3KClO + K2CO3 --> 2KMnO4 + 3KCl + CO2

Or, in general, 2MnO2(s) + 3ClO- + 2OH- --> 2MnO4- + 3Cl- + H2O
(We could have used hydroxide instead of carbonate, which in strong solution produces the OH- ions: carbonate is used to keep the pH so that permanganate is produced instead of manganate – this is why, I guess, that CO2 is used to convert manganate to permanganate in the fusion process)

The burden of Cl- ions is thus reduced in the ratio 5:3. We haven’t done anything to avoid the disproportionation of the ClO- ions yet.

This side reaction is 3ClO-  ClO3- + 2Cl-. (Alkaline solution understood throughout). The standard redox potential for this (also slow) auto-oxidation is about 0.49 volt, which is actually less than the 0.6 volt needed for the Mn(IV) to MnO4- transition. So my original assumption that the desired oxidation ought to beat the undesired were wrong.

My original assessment that about the same amount of chlorate was produced as permanganate bears this out. A word now about the kinetics. To produce one ion of permanganate needs 3/2 ClO- ions; but for the chlorate, 3 ClO- ions. This should favor the permanganate in less concentrated solution, so my assumption that the concentration should be high may also be at fault.

Finally, given that chlorate is inevitable, is there a method to separate it? from the permanganate? We have to think of some solvent other than water, and that means organic. Well, in another thread, mericad193724 in the thread "KMnO4 Synthesis" gave us the clue. Acetone dissolves KMnO4. Now it seems the stuff he was using destroyed his product, but it should not have, as far as I know. Permanganate converts alcohols to Ketones, and IIRC is then stable toward them. Acetone has a dielectric constant near 20, not too polar. But permanganate is a far stronger acid than chloric, although both are strong. Question: does acetone dissolve KClO3 to any marked extent. As far as I can gather, it is almost insoluble, <0.1 %, versus KMnO4 at about 11%. Anyone got a comment on this – acetone should separate the two. I don’t have any acetone at present, but must try this on my impure samples.

Sorry about the length of this but I just had too many thoughts and it helps to get it down.



ciscosdad - 20-6-2007 at 21:12

The acetone looks good. What about a continuous extraction with a soxhlet apparatus? I assume the extraction would be done on the dry product.
Don't know about KClO3 solubility in the acetone, but I would expect it to be low.
Acetone has a low BP, so it should all happen at nice low temps.

I like the idea of prior generation of MnO2 from Mn++.
I personally would use MnSO4 for its easy availability.

not_important - 20-6-2007 at 23:46

Simply exposing Mn(OH)2 to air will give MnO(OH), a mixed Mn(III) and Mn(IV) hydrated oxide, or even hydrated MnO2. So you could use the even cheaper air to save some on the MnO2 prep.

You don't really want to dry the MnO2 that much, just vacuum filter suction should do well. Heating it reduces its reactivity, heating to 200 might be enough to cause the lose of some oxygen.

Permanganate attacks acetone, although slowly. I suspect it may be through the enol form, a C=C double bond and a free OH group are prime targets for permanganate. Washing the glassware with a strong acid, then rinsing well might reduce this effect, as it is base catalysed. Technical grade acetone generally has reducible impurities in it as well. If you use it as a solvent, you don't want to have it sit around for too long to avoid too much loss of the product.

See doi:10.1016/j.jpowsour.2005.03.178

and interestingly enough this

[Edited on 21-6-2007 by not_important]

Attachment: reaction_hazards.pdf (110kB)
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DerAlte - 22-6-2007 at 07:28

Not_important, you always have something interesting (and important!) to say.

WRT to the keto-enol tautomerism of acetone, I had forgotten that. I am not much of an organic chemist. There’s no doubt KMnO4 eats up double C=C bonds like Pacman, and also attacks -OH bands in alcohols. I am sure you’ve seen the impressive action against glycerol. Looking it up in my organic text book, I find that KMnO4 can be used to create vicinal diols from alkenes, but only if you keep the temp low, below 25C. Above that you get splitting into a couple of acids.

I had proposed to use the acetone on dried product and avoid water, which may catalyze the enol formation. I’ll give it a try. Keeping it cool may help to slow any reactions.

The reason for drying the precipitated MnO2, xH2O was to be able to weigh the product. It’s easier, I decided, to then weigh the unreacted product later, to determine the amount of permanganate produced, than to do a ferrous ammonium sulphate titration. (One has to assume all the reacted Mn goes to permanganate, of course. It could go to Mn++ ions too, I suppose, but I don’t see how, in the presence of the hypochlorite)

As you can imagine, all this measurement is very tedious but necessary to get an idea of how worthwhile any method is. One then assumes that the chlorate comes from the ureacted part of the hypochlorite, the rest to chloride.

AS to heating, CRC says MNO2 decomposes at 525C; permanganate at around 200C. So it should be pretty safe to heat the hydrated stuff to about 200C. But one doesn't nkow how much water is left, unfortunately.

That doi abstract was interesting. Couldn’t download the pdf for some reason.



DerAlte - 22-6-2007 at 19:30

Managed to download the pdf, Not_important. Germane to the problem. I assume it refers to KMnO4 and acetone in a highly acidic medium from the number of H+ ions. What a hideous equation – too many prime factors in it!

Well I got my acetone, industrial solvent. I evaporated some on Al foil – very slight residue, probably organic (charred on rapid heat?) and a small less volatile liquid residue, probably water.

I now have a small pile of glistening black acicular crystals. It worked.

I cooled the acetone to about -10C first and poured it over a gram or so of my previous dried solid product, a mixture of chlorate, chloride and permanganate. This was in a small funnel with a glass plug. The permanganate dissolved readily without incident. The solution was just like a water solution – very deep color. I re-poured a few times and finally gave it a small wash. Then evaporated at ambient (30C+ here) in a shallow dish in an air stream from a fan – quite rapid. (Don’t mess with acetone indoors). The final result of this was a mixture of black goo in with some liquid and solid and no smell of acetone or ketone. The liquid was water, I believe. It evaporated too, leaving little black needles with no sign of cubic or plate type crystals typical of chloride or chlorate.

There were signs of MnO2 but not serious – slight brown stains. With permanganates this is par for the course. There may have been some reduction. If I were going to do this in any quantity I would not leave the acetone solution hanging around, would keep everything cool and evaporate under suction.

On the basis of this I claim to have made solid tolerably pure KMnO4 by a CRUD method – (Chemical Reagent from Utter Dross). I am not suggesting it is very practical, however. Yields are low and acetone isn’t that cheap, unless you re-condense it.



ciscosdad - 24-6-2007 at 18:59

Congratulations DerAlte!!!:cool:
Brilliant piece of work.

It seems a pity to evaporate the acetone off at the end and effectively lose it.
Do you think it will be feasible to vacuum distill to recover the acetone? What condition was the chlorate residue left in?

DerAlte - 24-6-2007 at 20:58

Ciscosdad, I don't see why one could not distill under reduced pressure. Acetone boils at about 59C, (IIRC) at atmospheric. Might want to use cooled water in the condenser.

Flash point of acetone is -20C. almost as bad as ether at -40C ( - in the climate here ether has to be kept under refrigeration). I cooled the solvent to minimize any reaction with it or impurities. Main impurity seems to be water, which could be removed by a drying agent such as anhydrous CaSO4.

I find from a manual I have that KMnO4 is used as a test for organic impuities in acetone. It should not, apparently, reduce KMnO4, assumedly at room temp.

The Chlorate (plus chloride) was heavily stained purple. of course. Some MnO2 obvious also in the insoluble remainder in the filter funnel.

I have a few more thoughts on this type of process but need to do a few more tests, if only to avoid making a fool of myself...



Ballermatz - 17-7-2007 at 08:35

Awesome work so far, "Alter" ;)

I found this VERY interesting article which provides detailed information on the production of various manganese components:

The industrial process is also outlined. One interesting fact is, that oxidation from Mn(+IV) to Mn(+VII) does not only go through Mn(+VI) but through Mn(+V) as well. The conditions favouring formation of Mn(+V) do not favour the further oxidation to Mn(+VI) however! This is why the industrial process is split into THREE steps; the first two are seperate roasting processes which bring the Mn from +IV to +V and then to +VI; the last one is electrolytic oxidation to permanganate.

Thus when you simply fuse MnO2 with KOH (oxidizing using air or KNO3 or whatever else), yields can never exceed 60% of theoretical. Only if you fuse twice, first favouring formation of Mn(V) and then formation of Mn(VI) you will get better yields.

It is also interesting that decomposition of KMnO4 is highly dependant on pH. Thus if you solve the fused mass in boiling water and the mass is too alkaline, decomposition will be greatly accelerated.

Direct oxidation to permanganate by fusing with KNO3 (similar to direct oxidation from Cr2O3 to dichromate) is impossible because KMnO4 decomposes at 240°C; much lower than the melting point of KNO3. NH4NO3 would be perfect - except it will most probably explode due to formation of highly instable ammonium permanganate :o

One encouraging fact is that electrolytic oxidation of manganate to permanganate does NOT require any special gadgets like nickel anodes or diaphragms. Even industrial electrolysis uses plain steel electrodes and most processes dont need a diaphragm.

DerAlte - 17-7-2007 at 22:02

@ Ballermatz: Thank you for the compliment. I have been very interested in Mn compounds for years. The range of colors, oxidations states, oxyanions and general chemistry of the transition elements has fascinated me since youth.

Welcome to the forum (though it’s not my place as a novice member to say!). I found your post very interesting.

I have never got the fusion with KOH to work other than to give a faint green manganate color. Even with the aid of oxidants. (Look for other permanganate threads on this site with the search engine for several efforts in this direction).

Yet older authorities claim that any Mn compound, fused with KOH and an oxidant like nitrate on a Pt foil will produce a very sensitive test for Mn due to permanganate formation! Must be the magic of Pt! Matches its price! No industrial process I have ever read about ever manages Permangante in one step, they always get manganate which is electrolytically oxidized to permanganate.

I long ago decided they must use a very special process. KOH melts at 406C; KNO3 @ 337C; KClO3 @ 368C with decomp.; and KClO4 at 525C. All these are way above the decomp temp of KMnO4, (240C, you quote) or manganate (190C). As for NH4NO3, no go, I’m afraid. It explodes at 210C. Worse, NH4MnO4 decomposes at 70C! (Figures from CRC).

AS far as the oxidation states go, in alkaline solution the direct step Mn(IV) to Mn(VII) has standard reduction potential of 0.60Volt versus about 0.62v for Mn(IV) to Mn(VI). However, the intermediate hypomanganate Mn(V) state MnO4--- has a potential of 0.96 volts to MnO2. I.e. permanganate is slightly easier to produce than manganate and much more so than hypomanganate. The ion MnO--- can be made by carefully reducing KMnO4 at ~ ph 14 with a sulphite at about 0C via the manganate. It is a light blue color and quite unstable. I have done this – the sequence purple Mn(VII) to green Mn(VI) to blue Mn(V) to black (MnO2) is quite remarkable.

I am still tooling around with these processes but am temporarily sidetracked. Hope to report some more later.



Ballermatz - 18-7-2007 at 05:48

Hi "Alter" ;)

Permanganates all have relatively low decomposition temperatures; KMnO4 seems to be the most thermally stable on my list; much more stable than say barium or copper permanganate which decompose below 100°C. Decomposition of the formed manganate or permanganate seems to be the most likely reason for the reported failures to produce them by fusing MnO2 with hydroxides. However, a liquid phase oxidation process IS used industrially, and involves fusing MnO2 ore with a large excess of KOH (1:5 molar ratio) for 4-5h. However, the temperature is said to be 250°C, yet the mixture is said to be liquid at all times (WITHOUT water). So they must use some kind of trick to lower the melting point of KOH.

Your hypochlorite oxidation method is clever in this regard, because it does not need such high temperatures. However it suffers from unwanted auto-oxidation of hypochlorite to chlorate. I remember reading somewhere that NaOCl can be composed into NaCl and O using catalysts like cobalt oxide, but I dont know to which extend it avoids the chlorate formation.

"I long ago decided they must use a very special process. KOH melts at 406C; KNO3 @ 337C; KClO3 @ 368C with decomp.; and KClO4 at 525C. All these are way above the decomp temp of KMnO4, (240C, you quote) or manganate (190C). As for NH4NO3, no go, I’m afraid. It explodes at 210C. Worse, NH4MnO4 decomposes at 70C! (Figures from CRC). "

Do you have decomp. temperatures for the hypomanganate at hand? I would guess they`re much higher because the industrial process uses 400°C during the first roasting (conversion to hypomanganate, Mn(+V)) and 200°C in the second, which forms the manganate. The first roasting kiln is fed with a slurry of KOH and MnO2 ore, but from the description it becomes obvious that the water is evaporated during the process. It seems like its only used to produce an intimate mixture and be able to spray the slurry into hot air. During the second roasting, however, water is continously sprayed onto the mass to keep it wet at all times.

This whole process sounds feasible for the amateur chemist as well. The first roasting process sounds a bit difficult to copy because it involves spraying a highly corrosive slurry into a rotating kiln. However I imagine that spreading out a intimate KOH/MnO2 mixture on a baking sheet so that a very thing layer is formed, then spraying water onto it and treating it with a heatgun might do the trick. It will be necessary to heat the baking sheat by an additional heat source because it has a large surface that will cool down rapidly. The hot air (>300°C) from the heat gun will provide the necessary O2 without cooling the mixture. Heat until it is dry, maybe mix yet again, wet with water and repeat the step. Important question is: How can the formation of K3MnO4 be observed? If figure it would be blue in aqueous solution but how does it look when dry?

In a second roasting step, one would carefully reduce the heat to 190°C and spray water onto it to keep it wet during the process. Dont forget that the whole process will be a matter of 4h or more! The hypomanganate formation is a matter of minutes but turning it into manganate requires more time.

Final oxidation to permanganate could still be accomplished using chlorine from TCCS or calcium hypochlorite.

Ballermatz - 18-7-2007 at 06:19

I just found that an eutecticum of 90%mol NaOH and Na2CO3 has a melting point of 286°C so there is hope ;)

Best regards

Der Ballermatz

DerAlte - 18-7-2007 at 09:00

Most fascinating! I believe the hypomangante Mn(V)O4--- forms blue crystals (according to Bauer). I've never made them. The solurion in 30% NaOH at around 0c is light blue and on keeping rapidly decomposes, especially on heating. Must try to repeat this.

In this context many of the older references talk about manganites of the general structure xK2O.yMnO2 which are generally insoluble and black in color, hence difficult to distinguish from the dioxide.

There appear to be several forms of crystalalline MnO2, possibly with different reactivity. I do know that technical grade, probably pyrolusite with 75-90% MnO2, is less reacive in the processes described than freshly made MnO2 as above. The color of MnO2 is generally blackish brown as precipitated, probably due to hydration. The precipitated stuff is also very finely powdered giving a large surface area.

It is obvious that the commerial processes use considerable art and experience and chemistry alone will not lead one to them!

I have some further comments but will reserve those for later, afyer I have done one further experiment to test the idea.



chief - 19-7-2007 at 13:07

Hello. I'm new here; why not make the KMnO4 from Ba(MnO4)2 by mixing solutions of the latter and of K2SO4: BaSO4 would precipitate, leaving KMnO4 in the solution.

The potential advantage to be hoped for would be the higher temperature-stabilities of the Ba-compounds, although yet unknown to me about the Ba(MnO4), but Ba(NO3)2 is stable up to ~ 595 [Celsius]. Thereby it might be possible to easily create the Ba(MnO4)2 and after the reaction with the K2SO4 reuse the BaSO4 by cooking it with concentrated Na2CO3-solution (which works!, although not quantitatively) and going from the carbonate again.
Only thing not solved for me so far: Ba(NO3)2 also does not melt below the 595 [Celsius];
so how then make the Ba(MnO4) from MnO2 (from old cells, it's the by easiest source !!!) and Ba(NO3)2 or a combination with other nitrates ??

Another route would be to anodically oxidize Mn in a solution of Ba-hydroxide (which gives at least the manganate)

You see: the way would be to somehow create Ba(MnO4)2, and from there on it should be possible to create a lot of manganates by just mixing the solutions of Ba(MnO4)2 with the sulfates of the elements -- and Ba(MnO4) should be easier to create because of the higher temperature-stability of the compounds.

Any ideas ??

chief - 19-7-2007 at 13:18

Another Idea from me about the Ba(No3)2-melt: should it not decompose immediately above the 595 [Celsius] (it somehow seemingly "boils" [releasing fumes] at 607-615 [Celsius], but I didn't measure too exactly) one might trade of the decomposition-time of it versus the creation/re-decomposition time of the thereby created Ba(MnO4)2 but temperature-control.
How is it: do nitrates decompose at a specified temperature or is it a temperature-range with temperatue-dependent decomposition-rate ? If it were the latter, a half-life-period exist (adjustable by the temperature). And it might be the same for the permanganate. so if these facts were known: 1-time purchase of a Ba-salt (carbonate or whatever), and just by putting K2SO4 and nitric acid into it obtaining the KMnO4 !!

Ballermatz - 19-7-2007 at 15:57

"The potential advantage to be hoped for would be the higher temperature-stabilities of the Ba-compounds, although yet unknown to me about the Ba(MnO4), but Ba(NO3)2 is stable up to ~ 595 [Celsius]"

What we need is actually the opposite - a nitrate that decomposes below 200°C so that the formed manganate will not be decomposed. Ammonium nitrate would do the trick but explosive compounds will be formed.

It also doesnt work because batrium permanganate has an even lower dec. temp. (100°C) than KMnO4. Ba(No3)2 decomposes >550°C - much higher than KNO3.

What we really need is a mixture that brings the melting point of KOH or NOH below 200°C.

chief - 19-7-2007 at 16:23

Is it really the case that Ba(MnO4)2 decomposes at such a low temperature? Is it not the general trend for Ba-compounds to be quite stable (see the oxide)?
Besides: Manganese nitrate is said to decompose at above 140 [Celsius] ...

I could have bet the Ba(MnO4) is at least stable up to 500 [Celsius]

DerAlte - 19-7-2007 at 20:34

I have Ba(Mno4)2 decomposing @ 200C (CRC)

@chief : The double decomp of potassium sulphate with barium permanganate would work like a charm -IF you have barium permanganate. Since the thrust of this thread is to find an alternative to the fusion process by a wet process at <100C. how do you propose to make Ba(Mno4)2? See my posting for 6-6-2007 above in this thread.

More later!



chief - 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.

chief - 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?

12AX7 - 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.


Ballermatz - 26-7-2007 at 04:33

Here's an interesting mixture:

KNO3-LiNO3-NaNO3 44.9-37.3-17.8

melts at 120 °C !

DerAlte - 5-8-2007 at 10:57


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.)


Der Alte

[Edited on 10-8-2007 by DerAlte]

DerAlte - 5-8-2007 at 11:00


(This may be of general interest.}

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

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.

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

Soluble—Potassium, mercuric.

Freely soluble—Mercuric.

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

Insoluble—Barium, bismuth.
Very slightly soluble—Sodium, potassium, lead.
soluble—Ammonium, crystalline nickelous.

Insoluble—Potassium, anhydrous sodium.

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

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

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.


Der Alte.

Ballermatz - 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...



DerAlte - 6-8-2007 at 19:31


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!

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.


Der Alte

DerAlte - 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:-

[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.


Der Alte

DerAlte - 9-8-2007 at 10:51


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.



Exotic Yellow Manganese Compounds

Cesium Fluoride - 9-8-2007 at 16:30

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.


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:

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]

DerAlte - 9-8-2007 at 17:18

@ Cesium Fluoride - OK! I think that covers the spectrum! Regards, Der Alte

Filemon - 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?

DerAlte - 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...



DerAlte - 15-8-2007 at 22:51


(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.

Ballermatz - 16-8-2007 at 16:43

@Der Alte

thx for the info!

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

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:

"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?


Der BallerMatz ;)

[Edited on 16-8-2007 by Ballermatz]

[Edited on 16-8-2007 by Ballermatz]

[Edited on 16-8-2007 by Ballermatz]

Ballermatz - 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.

DerAlte - 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.


Der Alte.

[Edited on 19-8-2007 by DerAlte]

Rosco Bodine - 18-8-2007 at 09:15

@Der Alte

You might like the following thread

[Edited on 19-8-2007 by Rosco Bodine]

DerAlte - 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.


Der Alte

DerAlte - 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.


Der Alte

[Edited on 19-8-2007 by DerAlte]

[Edited on 19-8-2007 by DerAlte]

ciscosdad - 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.

DerAlte - 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.


Der Alte

not_important - 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]

Attachment: US3986941A1.pdf (919kB)
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Google Patents

ciscosdad - 22-8-2007 at 23:02

@Der Alte

I do all my patent searching at Google Patents.

The actual patent:


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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