Sciencemadness Discussion Board - KMnO4 synthesis

KMnO4 synthesis

MephistosMinion - 5-5-2005 at 21:31

Hello all,
I have found a source of KMnO4 but they only have a kilo left, so I am looking into the manufacture of it. I have found this reaction:

2MnO2 + 4KOH + O2 -> 2K2MnO4 + 2H2O

2K2MnO4 + 2H2O -> 2KMnO4 + 2KOH + H2

I think I can do it by making a strong solution of KOH and then adding the MnO2 (from batteries) and bubbling O2 from a welding tank through the whole thing, the O2 bubbling through will also keep the mix stirring so the MnO2 dosen't settle. What are your thoughts?

BromicAcid - 5-5-2005 at 21:46

Thread on this

Slightly covered here on molten oxidizing agents

On manganates

Your second reaction should be a disproportionation reaction which is usually brought about by increasing the acidity of the mixture, CO2 works well from what I have read, it is somewhat different then what your equation describes. Check those links, good information there. The first reaction might work better if the ingredients were molten but it does work in the aqueous phase as some members of this board have tried.

Esplosivo - 5-5-2005 at 21:46

First of, the second disproportionation rxn is not correct. On boiling the K2MnO4 soln in an acid, preferably H2SO4 the following will occur:

2H2SO4 + 3K2MnO4 --boil--> MnO2 + 2KMnO4 + K2SO4 + 2H20

Filtering the solution will give you the purple solution you need, which is an acidified permanganate one. Now somebody can help with the isolation of the permanganate, probably heating until soln is saturated and then after cooling filtering off the ppt. crystals.

Secondly, this rxn is usually carried out in fused/molten KOH and not in a conc. solution. O2 from the atmosphere is usually enough if you are going to mix the stuff regularly. Or else you could add some oxidizing agent such as KClO3 or maybe KNO3. Hope this helps.

Edit: Upss, sry, Bromic got here first

[Edited on 6-5-2005 by Esplosivo]

[Edited on 6-5-2005 by Esplosivo]

Blackout - 6-5-2005 at 06:33

2 MnSO4 + 5 PbO2 +3 H2SO4 ---> 2 HMnO4 + 5 PbSO4 + 2 H2O

Mix MnSO4 and PbO2 together, add conc. H2SO4. Heat at 100°C during approx. 30 sec then dilute with a lot of water. Remove the PbSO4 and add a conc. solution of KCl. The KMnO4 should precipitate.

neutrino - 6-5-2005 at 13:07

HMnO4 will dehydrate to an oily liquid, Mn2O7. This won’t react with KCl.

The_Davster - 6-5-2005 at 13:30

If a KCl solution is added the Mn2O7 will undergo hydrolysis to HMnO4, which would then react with the KCl to form the potassium permanganate.

neutrino - 6-5-2005 at 16:21

On second thought, it might not dehydrate so easily without the dehydrating power of H2SO4, as it does in the preparation of Mn2O7. Still, I doubt you'd get a permanganate precipitate, as the protons liberated when the permanganate precipitates would probably react back to form HMnO4. Remember that weak acid + salt doesn’t give you a strong acid (HCl) and a salt.

The_Davster - 6-5-2005 at 19:55

Potassium carbonate or potassium hydroxide would be a better potassium source then KCl.
I too doubt the potassium permanganate would just precipitate out. However if excess PbO2 was used, the only soluble chem that should be left over after heating and adding KCO3/KOH, would be potassium permanganate, so the solution could then be evaporated to get potassium permanganate (providing that no excess of potassium salt was added which would contaminate the final product).

unionised - 7-5-2005 at 02:02

You might want to look into the properties of Mn2O7 before you think about making it.

Cesium Fluoride - 17-8-2006 at 23:20

I am interested in synthesizing potassium permanganate and this week I came painstankingly close. In short, my sodium manganate and sodium permanganate do not seem to be stable in any sort of solution.

I have looked at several preparations from various preparative inorganic chemistry books on this site. All of these books pretty much say the same thing. A finely ground up mixture of maganese dioxide and potassium hydroxide is heated strongly and the potassium hydroxide melts. Atmospheric oxygen oxidizes the mixtures to the green potassium manganate. Another oxidizer, most commonly potassium chlorate, may be added to speed up oxidation, but is not required. Then, the fused mass is cooled and extracted with boiling water. All three descriptions I have indicate the use of boiling water, which if you continue to read what I did, is quite interesting. The green liquid may be filtered at this point to rid of any insoluble manganese dioxide. The solution is then kept at a boil and carbon dioxide is bubbled into the hot solution. The solution turns from green to purple as the manganate ion dispropotionates into manganese dioxide and the permangante ion. The purple solution is then filtered once more, evaporated down, and cooled on ice so that crystals of potassium permanganate precipitate.

For my source of manganese dioxide, I opened up an alkaline battery and extracted the black paste inside. According to the MSDS for a duracell battery, this should be anywhere from 85-95% manganese dioxide with some added graphite. The problem is that I do not have any potassium hydroxide so I substituted in sodium hydroxide. This is seemingly advantageous for a couple of reasons. It is cheap and it also melts at a lower temperature than potassium hydroxide does. My theory was to make sodium permanganate and then displace this with potassium chloride to give potassium permanganate. Sodium permanganate is very soluble in water, but only about 6g of KMnO4 dissolve in 100mL of water at room temperature.

I mixed 20g of MnO2/graphite with 20g of NaOH and heated for 2 hours and 15 minutes. Within 10 minutes, areas of green could be seen. I stirred the mixture with a spoon and ground with a mortal and pestel every 15 minutes or so. By the end, I was left with a beautiful dark green powder that was coloured throughout. I have done several trials of this and each time I have gotten the green sodium manganate with or without the use of potassium chlorate.

Here is the problem:
I suppose it is correct to assume that I have a mixture of sodium manganate with some left over manganese dioxide, sodium hydroxide, and graphite. When I dissolve this powder in a substantial amount of water, I at first get a green liquid, but this quickly turns completely brown as the sodium manganate is reduced to manganese dioxide. At first I thought that perhaps I did not let the powder cool sufficiently and that hot sodium manganate in solution likes to be reduced. I actually found literature to back this up: “Sodium manganate (Na2MnO4), prepared by fusion of a mixture of natural manganese dioxide and sodium hydroxide; green crystals, soluble in cold water, decomposed by hot water”. So, then I proceeded to add ice water to the mix and this seemed to help a bit, but after 15 minutes of passing carbon dioxide, the solution had turned brown again. There was no sign of permanganate because I had not added enough carbon dioxide and the pH was still greater than 13. I then added one drop of ice water on top of some of the green powder and a huge brown spot formed. So, in conclusion my sodium manganate mix is very unstable in water at nearly all temperatures.

I did some more tests. I took a dilute solution of sodium manganate and while it was still green I passed in lots of carbon dioxide. It eventually turned red, not the purple color of potassium permanganate. I found literature to back this up as well, however. “Fusing the wolframite with sodium nitrate, for example, produced a sodium manganate (green color) which changed to sodium permanganate (red color).” This red color, however, is very transient. I immediately could see the brown manganese dioxide that formed from the dispropitionation. After a few minutes, the mixture completely is transformed to the brown sludge of manganese dioxide. Heat (boiling) also immediately destroyed both the green sodium manganate and the red sodium permanganate.

Sources I've read mention that weak acids are used for the conversion of manganate to permanganate. They typically mention carbon dioxide (carbonic acid) or acetic acid so I decided to give vinegar a shot. With dilute solutions of sodium manganate, the addition of vinegar gives the red permanganate but this is again quickly converted to the brown manganese dioxide.

Using boiling water as the text's instructions say results immediately in manganese dioxide! In fact, I did a test and placed some of the green powder in boiling vinegar and I saw a flash of red and then tons of frothing and brown all over the place.

I also tried various dilute strong acids, namely hydrochloric acid and sulfuric acid. The transient red color was again noticed, but a gas was also produced. I initially though this was chlorine produced by the reaction of manganese dioxide and hydrochloric acid, but when the gas also was formed with sulfuric acid, I concluded that it must be oxgen that is released while the permanganate is being reduced. Concentrated strong acids faired far worse.

So, pretty much I am following the texts except I am substituting sodium hydroxide for potassium hydroxide. I've also tried adding potassium chloride to whatever weak acid I was using in a hope that potassium permanganate would precipitate, but I just get the brown manganese dioxide again. Why is it that the respective sodium and potassium salts have such different properties? Is it true that sodium manganate and sodium permanganate are really this unstable (NaMnO4 is used industrially in at least one company and is advertised as a more soluble alternative to KMnO4)? Is there anything that I can do to stabilize the permanganate ion?

I've tried pretty much all I can think off and don't really want to go about making potassium hydroxide so I hope you all can help. I have about 30 grams of this powder stored so I'm willing to try any brilliant suggestions. Thanks!

Also, just as a note, sodium manganate does not seem to be soluble in acetone (like KMnO4 is), methanol, or ethanol to any appreciable extent.

not_important - 18-8-2006 at 01:33

First off, I don't think you had nearly enough NaOH, you should have it in slight excess. You should have gotten a pasty fusion mass that slowly hardened. That's part of the reason boiling water is used to extract; the permanganate to manganate reaction is fairly slow, normally the water will cool down quickly enough to not lose much that way.

Did you wash the raw MnO2? There is going to be other stuff in it besides carbon.

The graphite/carbon is a reducing agent, it will have to be oxidised away before you can oxidise much of the MnO2.

If you have K2CO3 you can use a mix of it and NaOH.

Cesium Fluoride - 18-8-2006 at 10:17

After heating, my mix was not one huge solid mass. It was more like a bunch of dry crumbs. I will try to repeat the experiment with twice as much sodium hydroxide and then extract with boiling water.

I did not take any steps to purify the MnO2. I was assuming it was carbon, MnO2, and perhaps some KOH electrolyte. The carbon does not seem to be a problem- at least visibly the powder turns very green throughout.

I do not have any K2CO3, but I do have KNO3 and KClO3. I tried with KClO3 and I couldn't see any visible difference, perhaps the reaction went a bit faster. When KClO3 + NaOH is used, will I get a mixture of Na2MnO4 and K2MnO4?

[Edited on 18-8-2006 by Cesium Fluoride]

not_important - 18-8-2006 at 10:50

What you would get would depend on the amounts of permanganate formed and of K+ and Na+ around. If there is enough K+ to allow most of the permangante to crystallize out, then you'll get KMnO4. NaMnO4 is so soluble that it shouldn't play a role so long as there is enough K around.

Calculate the amounts of MnO2 and NaOH, needed, then add a bit more NaOH. You might also try splitting the fusion result, extract one half with boiling water and the other with just hot ( 50C ?) water.

If the batteries are used you will have some Mn(III) and/or Mn(II), those will still oxidise up to manganate. I'd also expect a bit of zinc.

If there is a pottery supply store near you, not the little greenware and glazes type but one that sells clay and kilns, you may find that they sell MnO2 and K2CO3.

Cesium Fluoride - 18-8-2006 at 11:24

4NaOH + 2MnO2 + O2 --> 2Na2MnO4 + 2H2O

For 20g of MnO2, stoichiometry dictates that about 18.4g of NaOH are needed. Therefore, in my first experiment I used 20g of MnO2 and 20g of NaOH (for excess). Nevertheless, since I did not get a "solid fused mass" like the texts say so I am now trying another run with only 10g MnO2 and 20g NaOH.

My thinking was that I would make NaMnO4 in solution and then add KCl and then boil it down and crystalize KMnO4 (only 6.4g/100mL at 20C). NaMnO4, as you mention, has a much higher solubility.

The problem is that as of now whenever I add any heat (or even when a solution is left alone for long enough at room temperature) my sodium manganate/permanganate is immediately reduced to MnO2. Perhaps the carbon is doing this? If this is the case, then I should be able to extract the Na2MnO4 with boiling/hot water and then filter it immediately to rid of any carbon that may be present.

Cesium Fluoride - 18-8-2006 at 14:11

Ok, I got pretty much the same results as last time with 10g MnO2 and 20g NaOH. The mix was heated again for 2 hours and 15 minutes with occasional stirring/grinding. The only difference was that this time I could see more of the molten lye. The mix was still rather fine and I was able to extract it with a spoon and pestle. The mix was cooled to room temperature and then vinegar was added to part of the mass. The red color soon transformed (within minutes) completely to MnO2. Now, thinking about, I do not think I have got the permanganate ion at all. I think the pink/red color is just Mn+2 ions floating around. Why my Na2MnO4 would be reduced to Mn+2 I have no idea. Also, what is this gas I see when I add my acid (in this case vinegar) to the green solution?

Crude Na2MnO4
Right after adding vinegar
2 minutes later (notice the bubbling)

12AX7 - 18-8-2006 at 14:14

Which has a lower melting point, NaMnO4 or KMnO4? I'm thinking toss KCl into the melt.

I may just try that...although, will it work with a gas heater? I don't happen to have anything that gets reasonably hot without direct or indirect flame...er well I suppose I could go and do it in the induction heater, which is back on the bench...hmmm

Tim

Cesium Fluoride - 18-8-2006 at 14:22

Potassium permanganate decomposes near its melting point of 270C. I'm sure the case is similar for sodium permanganate. I've been using at hot plate which has a max temperature of 370C. The melting point of NaOH is 330C so I am just barely getting it hot enough. Any heat source should work I suppose although I'm not sure if I really follow your train of thought? How exactly do you propose to forming NaMnO4/KMnO4 directly from the fused melt?

12AX7 - 18-8-2006 at 14:29

Oops, that should read Na2 or K2, since that's what's made in the fusion.

Tim

Cesium Fluoride - 18-8-2006 at 14:35

Although I don't have those values on hand, I don't think that the melting point of Na2MnO4 or K2MnO4 is of much importance. I've found that as long as your NaOH (330C) or KOH (406C) is molten, the reaction to form the manganate will proceed.

12AX7 - 18-8-2006 at 14:37

Well, the point is, if you can get the reaction Na2MnO4 + 2KCl = K2MnO4 + 2NaCl in the melt, you won't have to deal with unstable sodium manganate.

Tim

Cesium Fluoride - 18-8-2006 at 15:02

I think that the "unstable" sodium manganate is exactly the issue

. Why for instance do we hear so much of potassium permanganate and not of sodium permanganate? Assuming their oxidizing powers are equivalent, the sodium salt should be cheaper to manufacture. The only info I can find about NaMnO4 is

Quote:

RemOx™ L ISCO Reagent sodium permanganate (NaMnO4) is an inorganic oxidant that performs chemically the same way as KMnO4, only in a more concentrated form. The significant advantage to RemOx™ L ISCO Reagent is its high solubility in water, allowing it to be a more convenient and concentrated form of permanganate when used for In-Situ Chemical Oxidation (ISCO).

I know virtually nothing about Na2MnO4 aside from the knwoledge I have gained through experimentation.

Adding a potassium salt to the melt to form K2MnO4 is something I have thought about. I haven't tried KCl, but I've done several trials with KNO3 and KClO3, both of which also supposedly help along the oxidation process. These mixes still are "unstable" when they are hydrated.

So, either a) Na2MnO4 is inherently unstable in which case KOH is ultimately needed or b) there is some impurity (graphite?) which is favoring reduction to MnO2.

I guess I should resort to synthesizing/finding KOH now. What a pain.

[Edited on 18-8-2006 by Cesium Fluoride]

Nerro - 18-8-2006 at 15:24

You might be able to get rid of the graphite in your MnO2 by heating the MnO2/C in a flame untill all the C has burnt up. Alternatively you might try washing it with a lot of water or acetone. In my experience graphite floats on water, I don't know if MnO2 will float as well but if it doesn't that seems like a pretty straight forward way of removing the carbon.

Cesium Fluoride - 18-8-2006 at 15:31

Welll, I have purified MnO2 a long time ago sucessfully. It was annoying though. It involved reacting the MnO2 with excess HCl to yield MnCl2. The soluble MnCl2 was then filtered from the graphite and then reoxidized back to MnO2 with bleach. This process could probably be made less annoying by reacting the MnO2 with sulfuric acid and hydrogen peroxide to yield the MnSO4. I might try the flotation method you suggest, but I reckon it would be difficult to distinguish the MnO2 from the graphite.

Right now, I really want to try the procedure with KOH- thinking about making some now

hodges - 18-8-2006 at 16:45

Cesium, I think you have some sort of impurity (perhaps organic matter somehow) that is causing the KMnO4 to be reduced to MnO2. What you described about MnO2 slowly forming is exactly what happens if you add an organic substance such as an alcohol to a KMnO4 solution. What's more, this reaction takes place much more rapidly in acidic solutions.

Another thing - if you can get an oxidizing agent, even KNO3, you should add it. You may be only getting a small amount of KMnO4 because oxygen in going to be limited unless this is done in the open or air is blown through the mixture. If you end up with only a small amount of KMnO4 then the slightest impurity could react and use it all up.

Hodges

Cesium Fluoride - 18-8-2006 at 18:04

Hodges,

Your explanation may be plausible. The only impurities I know of are graphite and zinc, but I guess anything could be in those batteries. Tomorrow, I will work on properly purifying my MnO2 and will post results then. I plan on reacting the MnO2 with dilute H2SO4/H2O2 and the reprecipitating MnO2 with bleach. This is much "cleaner" than reacting it with hydrochloric acid because no chlorine is formed. I am a little concerned, however, that the MnO2 will catalytically decompose the H2O2 before it has time to be converted to MnSO4.

The reaction definently does speed up under acidic conditions. However, it also does occur under basic conditons when no MnO4- is present. I guess this could still be rationalized- MnO4-2 is probably still a strong enough oxidizer to oxidize whatever impurity may be present.

I have done several trials with this KNO3 and KClO3. KClO3 is suggested in most books- see Walton's Inorganic IIRC. I cannot see any visible difference but perhaps the reaction with KClO3 is occuring faster. Nevertheless, the same problem of the MnO4-/MnO4-2 being reduced still occurs!

not_important - 19-8-2006 at 00:30

Simply washing with water several times should remove most impurities other that the carbon/graphite.

I believe KNO3 melts a bit lower than KClO3, but either should oxidise then carbon away. The nitrate will contribute alkali to the mix, chlorate just chloride. But enough of either can be used to oxidise the carbon.

It is important to remember that the chemistry in molten hydroxide at >300C is much different than happens in aqueous solutions at 100 C. For many D block elements higher temperatures, to a point, increase the relative stability of the higher oxidation states. You see this in ceramic glazes, MnCO3 gives dark red Mn(III), black MnO2, and under some conditions what may be even higher oxidation states.

Good exposure to air is needed, except in the thinnest layers some stirring is needed. In the similar process for producing cromates from Cr2O3, a mixture of alkali hydroxide or carbonate and lime - Ca(OH)2 - was used; the lime mostly serving to give a porous mass that increased the surface exposed to air.

One reason you see KMnO4 used so much more than NaMnO4 is that the sodium salt is hard to crystallise, because of its high solubility and maybe because it's just contrary. For many applications you want to control impurities, easily meter the amount of oxidiser by simply weighing it, or just to be able to keep a jar of it onthe shelf. All MnO4- salts are unstable in water, the decomposition to MnO2 is slow for KMnO4 but it does happen. So you want the dry salt for long term storage.

For some applications you just want the strong oxidiser, the sodium salt may be better because its high solubility means you don't have to move so much solution. So long as you sell the stuff pretty quickly, the stability shouldn't be an issue. (OK, you're having problems, I talking about the industrial manufacturers)

There are several manufacturers listed here

http://www.powersourcing.com/se/sodiumpermanganate.htm

Interestingly the US DEA wants to put NaMnO4 onto the Class II list. The USA sure seems to be interesting in slitting its own throat.

Cesium Fluoride - 19-8-2006 at 08:52

Thanks for that information not_important.

I'm not sure, but I think KClO3 decomposes into oxygen before it melts. Also, I am working at just above 330C (well below red heat). Does KClO3/KNO3 readily oxidize carbon at this low of a temperature?

My layers have been thin (next time I'll measure actually how thin) and I do stirr occasionally. I will try another trial in a couple of hours. This time I will use 10g MnO2/graphite, 10g NaOH, and 5g KClO3. I will thoroughly wash the MnO2/graphite with water as suggested and I'll use my slightly more powerful hot plate and keep the reaction going for 8 hours instead of 2.

I found an industrial process that uses NaOH/MnO2 and no mention of problems so this should be possible! I guess my problems are a) impurities and b) a low yield.

Can somebody give a valid explanation of why I get a red color instead of purple when I increase the acidity of my MnO4-2 solution? This occurs even when there are K+ ions floating around in solution.

Thanks guys for all your help. You have inspired me to keep on trying!

12AX7 - 19-8-2006 at 09:53

The only red Mn I know is somewhere around Mn(III), which can be seen in a strong acid solution. I have a solution sitting in my lab which is deep burgundy red. It hydrolyzes very readily.

KClO3 melts at a rather high temperature, but I could see if KClO3 + charcoal (note: not graphite) ignites on top of freezing lead.

I don't know how fast graphite oxidizes. It should be completely stable in solution (against MnO4 I don't know), and it doesn't usually oxidize appreciably in air until a good glowing redness. I don't know what an oxidizing melt would do to it, especially if it has capacity to absorb CO2.

Tim

Cesium Fluoride - 19-8-2006 at 14:44

Ok, I aborted my trial today after 30 minutes of strong heating. I washed my MnO2/graphite and noticed that there was some silvery impurity (zinc I suppose) which floated around. I removed this and since my <10g MnO2/C was very wet I dissolved my KClO3 (5g) and NaOH (10g) along with it. I boiled my solution down to dryness. I figured this would ensure an intimate mix of the reactants. As my mix was boiling the last bits of water away, a light green color appeared on the top of the mix. Some manganate seemed to form even before the NaOH was molten, which is odd. I then transfered my container to a more powerful hot plate (which I previously had not used), and within minutes part of the mix started to burn! I guess this was the KClO3 oxidizing the C? Pieces of the mix would glow red and then smoke a bit. At this point, the top of the mix was grey and was no longer green. After 30 minutes of heating, it was still grey and I thought that the trial was a failure. I cooled the mix down and added water for disposal, but then I noticed the water turned dark green! I did not even extract with boiling water because at this point the mix was half way down the drain, but I added vinegar and I got the familiar red color. This time the solution seemed a little more stable. I separated the solution in two containers and boiled one. It mangaed to stay red for longer than any other similiar solution I've worked with, but it soon turned brown. To the other half, I added some KCl in an effort to increase K+ concentration, but upon doing so the mix immediately turned brown. I then remembered that my KCl is only about 96% pure (or so I had calculated from the nutrition facts of a salt subsitute). I looked at the other ingredients and one of the impurities is fumaric acid! My MnO4- surely had oxidized across fumaric acid's double bond. I make my KClO3 from this KCl source so any mix that I use KClO3 with is going to be contaminated with fumaric acid. Hence, the instability of my solution. This, however, still does not explain the instability of my solution that is formed with only NaOH and MnO2.

I guess I will try another trial- this time with washed MnO2, NaOH, and KNO3. I only have 3g of KClO3 left and I don't really feel like purifying it/making some more pure KClO3 just yet. I think that using water to ensure an intimate mix of the reactants and the more powerful heat source helped. It is odd, however, that I did not see the mix turn green on the surface.

I still do not understand why my solution is red. The burgundy color you describe Tim is precisely what I have when I add an acid. This is a guess of what may be happening.

MnO4-2 (green) + H+ --> MnO2 + MnO4- (purple)

but immediately some impurity (in this trial, fumaric acid) partially reduces the MnO4-

MnO4- --> Mn(III) (red) + CO2?

Could this explain the gas formation?

[Edited on 19-8-2006 by Cesium Fluoride]

Edit: I have read that acidified MnO4- oxidizes and cleaves alkenes to a carboxylic acid and a ketone. But in the case of fumaric acid (2-butenedioic acid), you have two carboxylic groups as well. How would this affect oxidation?

[Edited on 19-8-2006 by Cesium Fluoride]

12AX7 - 19-8-2006 at 15:55

Better question- why is what little fumaric acid present causing the whole thing to reduce?

Is there any likely catalyst that causes MnO4 to break down to O2? H2O2 aside.

Tim

Cesium Fluoride - 19-8-2006 at 16:38

My thinking is that my yields are so low that the fumaric acid is able to cause the whole thing to reduce. I am firing up another reaction now and expect better yields. Will post results in 3-4 hours.

Did some more research and found some more stuff about the Na2MnO4 decomposing. Still don't understand it though; if I have a potassium salt in the melt I should have K2MnO4. Furthermore, there should just be MnO4-2 ions in solution that aren't particularly attached to any one cation.

“Sodium manganate, Na2MnO4, is formed when a mixture of equal parts of manganese dioxide and soda-saltpetre is heated for sixteen hours; the mass is then lixiviate with a small quantity of water and the solution cooled down, when the salt separates out in small crystals isomorphous with Glauber’s salt, and having the composition Na2MnO4 + 10H2O. These dissolve in water with partial decomposition, yielding a green solution.”

A similiar description of K2MnO4 is given, but nothing is mention of partial decomposition.

Edit: I admit the fumaric acid explanation is not entirely adequate, but it does explain what happen when I added KCl to solution.

[Edited on 20-8-2006 by Cesium Fluoride]

not_important - 19-8-2006 at 19:44

Consider that you
1) did a more thorough washing, and saw some crud separate
2) saw a distinct oxidation during the heating, suggesting carbon being oxidised.
3) use a potassium source that included a decent reducing agent.

If what you removed in 1 and 2 were busy reducing the manganese in higher oxidation status in your fusion mix, you might have been having very low yields of manganate; low enough that small amounts of reducing substances could overwhelm your desired product.

It sounds as if you are doing everything correctly, so it likely is a problem with unwanted reducing agents. Long ago I tried to make KMnO4, and had better luck than you. But I was using MnO2 made from the remains of chlorine generation, and fertiliser KCl.

Adding a potassium based oxidiser may cause some KMnO4 to crystalise out, but you need enough K(+) and MnO4(-) to exceed the solubility products for KMnO4 else it's just ions, as you said.

Permanganate will at first convert one of the double bonds into a pair of hydroxyl groups, giving a diol. That is then further oxidised with breaking of the remaining C-C bond, to give a pair of carbonyl oxygens; depending on the structure around the original double bond those will be ketone or aldehyde oxygens, aldehydes generally get quickly oxidised to carboxylic acids.

So in the case of fumaric acid you create two carboxylic acids groups, each being one half of an oxalic acid molecule; one fumaric => two oxalic and the oxalic is further oxidised to H2O and CO2

C4H4O4 + 6[O] => 4CO2 + 2H2O

plug KMnO4 => K(+) + MnO2 + 2[O] into the above and you can see why that 4% might be causing problems for you.

Cesium Fluoride - 19-8-2006 at 20:35

not_important, with your help and that of others I think I may be close to figuring this whole thing out. I suppose it is safe to assume that the gas formed was indeed CO2. Thanks for that explanation.

Quote:

If what you removed in 1 and 2 were busy reducing the manganese in higher oxidation status in your fusion mix, you might have been having very low yields of manganate; low enough that small amounts of reducing substances could overwhelm your desired product.

Precisely, what I'm thinking.

Quote:

Long ago I tried to make KMnO4, and had better luck than you. But I was using MnO2 made from the remains of chlorine generation, and fertiliser KCl.

I'm curious. Were you using NaOH and did you get a purple solution upon acidification? And you used the KCl to make KClO3 or added it to the NaMnO4 solution?

I have another brand of KCl salt substitute here in front of me. Its impurities are <1% cream of tartar (potassium bitartrate), silicon dioxide, and natural flavor. Silicon dioxide shouldn't be too problematic as it'll just react to form silicates. It looks to me as if the hydroxyl groups on the bitartrate ion wold be easily oxidized by MnO4-, however. I suppose I could easily purify both brands of KCl by fractionally crystallizing the fumaric acid/potassium bitartrate out as they are both almost insoluble in water.

I think if I properly purify my reagents I should be able to get the desired purple solution. I would have never guessed that impurities would be so problematic

12AX7 - 19-8-2006 at 20:49

I reproduced the reactions this evening.

I put eyeballed amounts of NaOH and MnO2 (pottery grade) into a steel crucible (with some zinc still brazed in that I was too lazy to dissolve out). I heated the crucible with my induction heater, so open air, no CO2 or H2O besides what's in the basement. The NaOH melted uneventfully (if a bit sticky at first, from the humidity). Adding MnO2 immediately produced bubbling as it was wetted (CO2? H2O?), and the sludge turned dark forest green and reached a thick consistency. No further change was observed, though bubbling continued, especially when heated further (I have only guesses at the temperature).

The second attempt was identical (give or take), but I added KCl to the mixture. It didn't dissolve well, and the final (molten) product had lumps at the bottom.

Both materials dissolve in water to give a green solution which hydrolyzes easily, ultimately producing abundant brown sludge (MnOOH or whatever). A sample of the green solution, acidified with H2SO4 (~20%) immediately turns to pink or red, and hydrolyzes. The color is indeed very similar to my strong acid solution of Mn(III). HCl added to this solution causes the color to turn yellow, then mostly colorless (it was a weak solution), and brings the odor of chlorine.

Both solutions are still sitting, I'll check them tomorrow.

It is worthy of note that I have the same potential impurity: zinc. In addition, my stir stick was copper plated graphite, and I think some of the copper dissolved, causing the melt to occasionally produce blue coloration on creep surfaces (e.g., over the crucible's rim). Also, when the fusion was cooled, breathing on the product (while still reasonably hot) produced a chalky blue surface in spots. I'm guessing this is from sodium cuprate of some description.

I would like to know if these reactions can be approached from the other side; does a KMnO4 solution decompose when Na ions are present in a basic solution of MnO4(2-)? Does zinc catalyze decomposition? Graphite powder? Copper? Iron?

Tim

[Edited on 8-20-2006 by 12AX7]

Cesium Fluoride - 19-8-2006 at 21:08

Tim, thank you for proving I am not crazy- I appreciate it! All of what you describe sounds identicle to what I observed. I once quickly filtered some of the red solution and stuck it in the freezer. I checked two hours later and it was completely brown.

Aside from perhaps focussing on ridding of the zinc impurity, I don't really know where to go from here. Do you think that there is something inherently problematic about using NaOH instead of KOH? This logic makes no sense to me though.

Also, I actually may have seen some of that blue stuff in one of my trials and I don't think I had any copper around. It was a very small amount and was only concentrated in one small section of the melt atop of intensely colored green area.

In the library here, I found this info about NaMnO4 (Condy's liquid) being produced industrially.

"The substance [sodium permanganate] is obtained by mixing the caustic soda obtained from 1,500 kilos of soda-ash with 350 kilos of finely-divided manganese dioxide in a flat vessel, and heating this mixture for forty-eight hours to dull redness. The product is then lixiviated with water, and the solution either boiled to the requisite degree of strength or evaporated to dryness. If the manganate is to be completely converted into permanganate it is neutralized with sulfuric acid, the solution concentrated until Glauber’s salt separates out, and these crystals are then removed and the liquid further evaporated."

Edit:

Quote:

I would like to know if these reactions can be approached from the other side; does a KMnO4 solution decompose when Na ions are present in a basic solution of MnO4(2-)? Does zinc catalyze decomposition? Graphite powder? Copper? Iron?

This is an ingenious idea! I hope somebody can try this out, but unfortunately I don't have KMnO4 (that's why I'm trying to make it) Four different Sears in my area and the local hardware stores don't have it.

[Edited on 20-8-2006 by Cesium Fluoride]

not_important - 19-8-2006 at 21:22

Or just roast either KCl source in the open air, up to where it is just starting to glow dark red if you can reach that; stirring a bit. Should break up any organic and the air will take it to CO2 and H2O.

I just used KCl as a source of potassium, plain NaOH + MnO2 for the roast. I rigged a small pan with a 'propeller' stirring bar, using a battered boys' construction kit, so the mix was constantly being churned. I think it took about 15 seconds to do a revolution, which is much faster than needed.

Dissolved in hot water, added KCl, then bubbled CO2 from dry ice through it for some time. It was sitting on a warming tray, designed to keep cooked foods warm, used to pick them up for next to nothing at second hand stores. But it wasn't hot, around 40 to 45 C, and got that nice permanganate purple when a bit was squirted into a test tube and let stand so the MnO2 would settle. Finally filtered, stoppered the filter flask and put it on the warming tray, kept the vacuum going to quick concentrate until I thought I saw crystals, poured into a small beaker and let cool.

It was a lot of work, needed some scale up design. I did one more run to verify that using crude MnO(OH) from mixing hot MnCl2 solution with Ca(OH)2 and blowing air through the hot mix, would work in place of MnO2. Used NaOH + Na2CO3, didn't get hot enough to decompose CaCO3 so it came out with the MnO2.

About the I ran into a description of the electrolytic process to replace the acidification stage, you get a lot less MnO2 formed. That put things on hold until I could find out more about the electrolytic process, and then I found a fairly cheap source of the commercial product - water treatment - so there wasn't a need to make my own.

If I was going to try it again, I would go for making a rotary klin out of a longish piece of iron pipe and a portable cooking grill, some extra piping to preheat air being blown into the pipe. A motor and gearing/belts to slowly rotate the pipe to stir the mix. Big enough to run at least a kg at a time. And the electrolytic route from manganate to permanganate.

Consider how difficult making hydrazine can be, impurities can be a real problem. Or put a tiny bit of cobalt oxide into H2O2.

OoooooooOOOoooooooo...

12AX7 - 20-8-2006 at 19:49

Don't ask me how exactly I got this... I just did. LOL

This was from the second solution, the one with (at least some) potassium.

The first solution, which I idly added a little spare acid to last night, is fully hydrolyzed. The liquor is a little yellowish, and I think very basic. Lots of brown sludge.

The solution with potassium is still dark green to blue, though containing ample sludge.

Both solutions sat out overnight, so have absorbed CO2, and who knows what else.

I took some of the potassium solution and added a mild acid, carbonic acid. Hey, we have club soda in the house! w00t! Not much effect besides dilution (even with a large amount of soda, a small solution remained greeney/bluey), so I then added drops of diluted H2SO4. After enough addition, it turned pink and became efforvescent -- probably CO2. (In the RBF, I tested the gas with a match and it extiguished it; most likely CO2, not O2 loss.) The strongly basic solution ought to have absorbed CO2, after all. The smell was sort of musty, like wet rock.

I added acid until a floc precipitated. The solution looked depressingly red, but tantalizingly it also looked purple around the meniscus. So I set the flask down, and as the floc settled I discovered what is shown in the above photo.

So, I don't really know -- dilution perhaps? CO2? Something in the club soda? (Shouldn't be anything other than H2O and CO2, eh?)

I'll try this somewhat more controlled...like measuring things for instance.

Tim

Cesium Fluoride - 20-8-2006 at 20:08

Wow, I am SOOO jealous! This second batch that worked- you waited a day before you added the acid to what was a green mixture? Odd that the first mixture completely reduced and the second one did not.

woelen - 20-8-2006 at 22:38

I don't want to disappoint you guys, but how much starting material did you have? If I look at that RBF, then I can only say that this is a VERY VERY dilute solution of KMnO4, most likely less than 0.01% by weight. KMnO4 has an increadibly strong color, and even a tiny crystal of 1 mm size makes a whole liter of water almost opaque, so the yield of your experiments is very low I expect. But anyways, you made some KMnO4 and you might be able to optimize the procedure to get a better yield.

Cesium Fluoride - 21-8-2006 at 17:52

I was doing some more experimenting today and though what the heck I am just going to try adding some oxidizers to this Na2MnO4 solution. I first added some 3% H2O2 to the green solution and I got a slightly yellow liquid with MnO2 precipitate at the bottom and lots of gas given off. I added some vinegar to another batch and got the old red solution. To this I added H2O2, and it turned colorless (I couldn't see any pink but presumably these were Mn+2 ions).

Then I tried bleach. I added lots (didn't measure) of bleach to another green solution of Na2MnO4 and no color change was observed. I then added some vinegar to this and it turned purple!

purple

Woelen, a little more concentrated, no?

12AX7 - 21-8-2006 at 18:28

I've got better results now, by fusing the reaction:
2NaOH + 2KCl + MnO2 (+ O) = 2NaCl + K2MnO4 + H2O
-- Presumably. I don't know if potassium or sodium manganate is preferred. XRD of the glom could probably figure it out...

The stoichiometric mixture (17.5g KCl, 20g NaOH, 22.5g MnO2) reacted at about the melting point of NaOH, forming a solid mass. I continued heating and it fully melted (to about half the solid volume) around the melting point of KCl (also around the curie temp of the steel crucible, about 1400-1500°F) to a black, slightly greenish mess. This dissolved to a green solution, with much hydrolysis.

Interestingly, the green solution hydrolyzes in large volumes of water to a weak purple color (which is what, ppm MnO4?).

I took some of the solution, at any concentration, and acidified it until purple. After settling the MnOx out, I put the solutions together and evaporated it. The solution is currently cooling in the 'fridge, a deep purple color but I'm not sure just how deep it is. For sure, I've lost a *lot* of manganese from hydrolysis. Damn its strong color!

I may try weighing the fused crud and adding it to a stoichiometric acid solution. Not sure what the excess of acid is going to do, though... probably oxidize the chloride present...

Bleach, huh? That's a possibility

Tim

Cesium Fluoride - 21-8-2006 at 18:44

With those high temperatures, I expect you should get some pretty nice yields. I left this last reaction going for 6 hours, but at only 370C, I can't hope for much. I have read that if you add enough water to Na2MnO4 it dispropotionates because in such dilute concentrations the pH is lowered enough for the conversion to occur. I made a larger purple solution now and I added some KNO3 and am currently boiling it down. I have so many ions in solution- acetate, sulfate, hypochlorite/chlorate, chloride, nitrate, etc. that I'm not sure if I'll be able to get KMnO4 to crystallize out. The color is a deep purple but I have no idea of its concentration. There isn't too much MnO2 crud on the bottom (I filtered it), but probably more will form during boiling.

The_Davster - 21-8-2006 at 23:01

I figured I would join the permanganate party

8g KOH and 3g KClO3 were melted in a metal crucible(either SS or nickel, not sure, acquired by dumpster diving) in an electric furnace. When it melted I slowly spooned in 5g MnO2 while stirring with a graphite rod. As more and more was spooned in the mix solidifed to a green mass. The heating was continued for 15 min. I thought the crucible was reacting at one point because the mix was a faint red and no longer green, but then I relaized I just had the heat too high and the crucible was glowing red. It was around this time I realized the crucible was stuck in the furnace hole. I did not bother to wait for the mix to cool, I just dumped 150mL H2O into the crucible after turning off the furnace, not ideal, but it was getting late. Made a minor mess, very little lost. I then boiled the deep green solution while adding some dry ice lumps to it. Two lumps of dry ice added and the solution was completly turned to very very dark purple. Filtered through a glass filter, MnO2 was collected.

Procedure based off Walton's Inorganic Preparations, I was too lazy to bother filtering once before adding dry ice though. Might isolate it tomorrow, but boiling a concentrated highly coloured staining solution is not my idea of a good time.

[Edited on 22-8-2006 by rogue chemist]

guy - 21-8-2006 at 23:08

Using CO2 to acidify may result in purer solutions because bicarbonate is relatively insoluble in high concentrations and in cold water.

The_Davster - 21-8-2006 at 23:21

Forgot the picture

permanaganate.JPG - 17kB

woelen - 22-8-2006 at 02:06

Quote:

Woelen, a little more concentrated, no?

It is more concentrated, butI still can look through it. Solutions of KMnO4 must be totally opaque, even a few mm of it, otherwise you will have a solution of no more than a few hundreds of percent only. The color of KMnO4 is REALLY strong.

Rogue chemist, your picture looks more promising. It is really dark and also along the glass, a thin layer still looks quite purple. It is sad that your prep requires more exotic stuff. Dry ice is not something most people have at their home and the use of KClO3 also is a drawback, but the latter can be made by means of an electrolysis cell, so that is less of a concern.

The_Davster - 22-8-2006 at 09:09

The use of dry ice can be replaced by bubbling in CO2 from the standard baking soda and vinigar reaction.

not_important - 22-8-2006 at 09:38

In North America, dry ice is available at many supermarkets and elsewhere for much less cost than making the same amount of CO2 from a cheap carbonate and acid. In other areas this site http://www.dryicedirectory.com/ might help you find a retailer by asking a listed wholesaler.

CO2 has the advantage that the mother liquer after crystallisation of KMnO4 may be reused as part of the alkali for the next batch. At the low red heat of the roasting process, the Na/K carbonate will react as the hydroxide does with MnO4 or MnO(OH).

The use of chlorate is simply a speed-up device. When I ran this, I leave it roasting for 4 or 5 hours with mechanical stirring going. Many of the industrial process documents roast for much longer than that, but they are using thicker layers.

Cesium Fluoride - 22-8-2006 at 10:26

Nice job, rogue chemist!

Woelen: I have a feeling that the temperature of the reaction helps improve yields more than the means of acidification.

The_Davster - 22-8-2006 at 22:06

Actually, as I have been messing around with this reaction a bit tonight without oxidizer in the mix, it seems time is more important than temperature.
Do not add all the MnO2 at once, at it little by little, stirring the melt for quite some time until it is pure green after each addition. I have noticed that after adding a lump of MnO2 the size of a BB to the molten KOH that the melt darkens, and then after a time of stirring, then lightens to the origional green. Adding all the MnO2 at once produces a paste that is grey and almost solid. If this putty is blown on, it darkens on the surface to green. Kneeding this putty with the stirrod causes it to reveal gery areas of unreacted MnO2 which slowly react with oxygen in the air turning green, unfortunatly just on the surface.

As it took me over an hour to add 1g(out of 5g) to the molten KOH while assuring it had all reacted before adding more, doing this reaction without oxidizer present seems a waste of time. After an hour and a half I simply was out of patience and dumped the MnO2 in, causing a mess of solidifing putty, which in my opinion can be ground and reheated as many times as you want, but which will never equell a good dissolved oxidizer or slow addition.

The literature says to add the MnO2 to the molten KOH and chlorate right after it melts for good reason, as this is when the chlorate will be decomposing into oxygen, so it is beneficial to make sure you have the MnO2 and oxygen in the same place to get the manganate.

If you are wanting to pursue this reaction without extraneous oxidizer I cannot stress enough to add the MnO2 slowly making sure it is all reacted before adding the next portion.

12AX7 - 22-8-2006 at 23:11

Ah, I've got some solutions that dark.

My primary batch of liquor must have a lot of other stuff in it, though.. the meniscus is still see-through-able, and the color is deep purple, but not much deeper in color than, say, a concentrated CuCl2 solution is green (considering purple is a darker color than green).

Tim

mericad193724 - 18-11-2006 at 13:16

I am working on some KMnO4 now too...

I fused 20g NaOH and 20g MnO2 for 1/2 hour and added the green sodium manganate to 300ml of water and then added 25g KCl. I heated up the mix and bubbled CO2 gas through the mix while it was very hot. I let it sit for 1 week so the auto oxidation happens and bubbled CO2 through it again...ITS REALLY PURPLE!!!!

My question is can I filter the solution through paper filter paper, or will the KMnO4 oxidize the paper and I will loose a lot of my product?

thanks

Mericad

12AX7 - 18-11-2006 at 18:32

It reacts with paper. Matter of fact you can dab some MnO4 solution with paper and watch the color change as the purple neutralizes.

Oh, an addendum to this thread -- I decomposed the permanganate solution and crystallized it out. I got about 50g total NaSO4, NaCl and KCl from the solution and a pile of guess what... maybe a few grams of brown crud.

Tim

SOoo CLOSE!!!

mericad193724 - 22-11-2006 at 17:32

I have been continuing my project of making KMnO4 and I was really close today, but failed!

I decantated my KMnO4 solution that was really dark and boiled it from 250ml to 50ml. I added 200ml 100% nail polish acetone to this and heated to dissolve all the KMnO4; this was followed by a decantation.

This acetone solution was pretty dark and I boiled it down expecting to crystallize KMnO4 crystals at the end... I left and returned when the liquid was down to about 75ml, but it was clear!!!! With a lot of black precipitate (MnO2).

I must have all decomposed into MnO2!!!!

What caused this???? I realized that the Acetone was denatured with Denatonium Benzoate, could the KMnO4 have oxidized it? The other culprit could have been the earthenware piece of pottery I put in as a boiling stone. I don't think the KMnO4 changes the acetone itself.

Here are some pictures...

The KMnO4 solution after standing for one week but before being decantated from all the MnO2

The KMnO4 solution after bubbling CO2 into it twice.

KMnO4 in acetone, slightly lighter color (should be ONLY KMnO4).

The end result

a clear solution with a good amount of black precipitate.

What caused the decomposition, was it the Denatonium Benzoate? Boiling Stone Earthenware Pottery? Or acetone?

We are so close to making KMnO4 a practical synthesis!

I will retry the experiment if I can figure out my mistake.

Mericad

[Edited on 24-11-2006 by mericad193724]

Aqua_Fortis_100% - 4-6-2007 at 07:47

Really Sorry by disturbing this old thread again, but i think , i've noted a few things..

some days ago, i'm traveled the Frogfot page about dichromate synthesis and see the
reaction:

Quote:

If first step of reaction is carried out without a base, dichromate will form directly, however, this will produce lots of toxic NOx. I believe reaction will go as follows:

Cr2O3 + 2KNO3 --> K2Cr2O7 + 2NO

so, the follow can exist (???):

MnO2 + KNO3 --melting--> KMnO4 + NO

when i see the possibility of such reaction, i'm turned very happy..
Although the very bad properties of the NOx to the health, many byproducts can be made easily, if anyone have the right stuff on hand to control these unwanted toxic gases.

And If the person goes out throgh the NO (to oxidize it, all of the gas or partially to NO2

the O2 supplied by, e.g. an OTC and cheap aquarium pump or another device),theoretically the amateur chemist can (i think)obtain:
- nitric acid (extra O2 supplied and condensation, following by the freedom of the waste gas through a outside place or absorbing it with some proper solution);
- nitrites (unfortunately, i have no idea how the person can adjust the O2 inlet , to form equimolar amounts of NO and NO2 (NO + NO2 <---> N2O3

);
- "lead chamber" process, to catalysis the SO2 oxidation and to obtain H2SO4 or maybe , hopefully , even oleum ;
- anything more interesting (???)

but after, thinking about, i see only a problem: i think which without proper equipment, the yield of KMnO4, originally desired, can be quite low, since the manual stirring should be forbidden ( NOx = death

).. so, i'm planning also create a mini improvised eletric stirrer , such as mini-engines of toy car... IIRC should exist a thread lying here, so as soon as possible, i will search.

what about this reaction, is possible ???
thanks

EDIT#2: the gas , just regenerating (big??) parts of the original nitrate decomposed by the reaction, just absorbing in an alcaline solution... (the only problem can be the "suckback" desgraceful problem)
i can not see if is feasible scales up this.. MAYBE..

[Edited on 4-6-2007 by Aqua_Fortis_100%]

12AX7 - 4-6-2007 at 14:00

Similar thing happens with chlorate -- 2KClO3 + Cr2O3 = K2Cr2O7 + O2 + Cl2. I tried running it with O2-neutral stoichiometry (using KCl as a potassium source) and got green remainder, even with the help of some additional KClO3.

Back on topic, MnO2 of course decomposes chlorate, so that's no good. One could try adding KClO3 or KNO3 to a molten KOH / MnO2 slurry, but that's not far from the usual synthesis.

Tim

Aqua_Fortis_100% - 5-6-2007 at 03:35

thanks for the share, 12AXT.
I'm really interested in this and will try as soon as possible (still in the "MnO2 recovery from battery").. the fact is which here KMnO4 is very expensive (10 100mg tablets (1g) = $$$$$ ) and the method using a base seems to me to be a little bitc* to purify and clean the materials... and i will like to experiment new things..(but i will try also the "conventional" method, to see what good will be the product. (using NaOH - KOH is something rare here :mad:

))

another thing i probably will like in melting these stuffs without a base ,as i say ,is the possibility in make others good products at same time.

although ,another possible BIG problem i see is ,maybe, the KNO3/MnO2 coming too hard to be stirred.. i'm remembering now when i tried a sodium nitrite syntesis with NaNO3/Ca(OH)2/graphite .. the stuff (NaNO3

beginning to melt, but when i progressively added lime/graphite become to hardening and the stirring impossible to do.

( 2 NaNO3 + Ca(OH)2 + C ---> 2 NaNO2 + CaCO3 + H2O )

so what device you used to stirr the thing releasing toxic Cl2 (the reaction using chlorate you mentioned)???
what you sugest to begin?
Thanks again.

12AX7 - 5-6-2007 at 15:03

Well, the dichromate synthesis starts off as fluidized powder, so enough gas is given off that it remains liquid. At a critical point, it stops gassing and "combusts", turning orange. The potassium dichromate melts around red heat.

Tim

Chemophiliac - 5-6-2007 at 19:23

So, what are the properties of Na2MnO4 or K2MnO4? I assume it decomposes in warm water. Also, does it have perculiar magnetic properties? Additionally, could someone tell me if my K2MnO4 synthesis method will work well? I would like someone with more knowledge then me to evaluate it.

Essentially, my synthesis of K2MnO4 could be written as following:

4 KOH + 2 MnO2 + O2 --> 2 K2MnO4 + 2 H2O

How could I make KMnO4 from this without having soluable components like KCl or K2SO4 left over?

The_Davster - 5-6-2007 at 19:46

Yeah, that will work, but not well, and probably only isolatable yield in some sort of industrial furnace. In practice oxygen can only react at the surface of the melt, which is why something like KClO3 is added as it decomposes releasing oxygen throuought the mass oxidizing to manganate. I tried it at one point without oxidizer and was left with only the surface oxidized. See the posts I made last year on the second page of this thread.

12AX7 - 5-6-2007 at 19:59

You can't- think about it, you have excess K, it has to be paired with something.

Na2MnO4, as near as I can tell, is stable (I've had it up to red heat and it remains deep green), but hates water.

Ya know, after my experiences with chromate, I may have to try permanganate again. As I recall, I omitted any oxidizer -- it turns green on its own, but so doesn't chromate. But there's a big difference between the drab yellow sort of color you get on the surface of a chromate fusion, as compared to the throughly deep red you get from a pure potassium dichromate melt. Likewise, the drab, exceedingly dark green color may be so dark due to Mn(IV) (MnO2 or manganites). I'll have to try it again with KClO3 (my only prodigious oxidizer...next to that 5lb jug of KMnO4...er yeah...anyway...) and see what happens.

Ooh, and it's a source of potassium. Lemme see...
3 MnO2 + NaOH (shit, I don't have any more NaOH...I need to find some*) + KClO3 = 3 K2MnO4 + NaCl + 1/2 H2O, hmm that needs about 5 KOH and another 2.5 H2O to balance, doesn't it.

Tim

*Red Devil lye is off the shelves now. A little farther down the shelf I saw a big heaping bucket of Rooto Number 2, which claimed 86% NaOH or something. Any idea what the other 14% are?

Aqua_Fortis_100% - 6-6-2007 at 10:59

Guys, thanks a lot for the great amount of good info.. armed with this knowledge i will save LOTS of money!!! Permanganate is very good stuff for pyro displays , and great fun things can be made.
I will probably try synthesise small amounts in different ways in this weekend if the time gets good.

unfortunatelly , i can't buy or get anything chromium based to try dichromates, because these things were forbidden by local law because of *somewhat* carcinogenic nature... shi* state of burocracy!
any idea?

Quote:

originally posted by 12AXT:
*Red Devil lye is off the shelves now. A little farther down the shelf I saw a big heaping bucket of Rooto Number 2, which claimed 86% NaOH or something. Any idea what the other 14% are?

sorry by this little off topic, but, isn’t the “Rooto” stuff a sulfuric acid based product ?(better yet question: are in the US (which i assume you live) , ALL drain cleaners based on beautiful sulfuric acid ? )(UNFORTUNATELLY :mad:

, here this stuff is forbidden..so the usual source is the infamous acid battery..)
although here this NaOH for drain cleaner purposes is readily avaliable.. some brands are 98% content

maybe the buffers of your rooto can be some usual : chlorides, quicklime and even (i see on the label of a generic brand on the shelves here) sodium chlorate, etc,etc,etc. (The unknow chemicals in some products also up set me very much!!!)

[Edited on 6-6-2007 by Aqua_Fortis_100%]

Chemophiliac - 6-6-2007 at 14:32

Because this whole post is so long and confusing, could someone just sum up the whole reaction from MnO2 and K salts all the way to KMnO4 in one reply? Thanks, much would be appreciated. What is my best bet for getting a high kield of KMnO4 from the original reactants. I want to make KMnO4 so I can use it here for various lab procedures, like as an oxizing agent, for making exotic permanganates. KMnO4 is very effective in redox reactions so I can get other metals up to high, sometimes unusually high (for a particular elelment- like Ag++) oxidation states, as my chem teacher once told me.

The_Davster - 6-6-2007 at 14:43

Walton's Inorganic Preparations
http://www.sciencemadness.org/library/index.html
Full synthesis is in that book.
(my DJVUviewer appears to be broken)

12AX7 - 6-6-2007 at 15:27

Ok-

So, you need some source of manganese, obviously. MnSO4 and MnO2 are common fodder. In the former case, you'll want to precipitate some form of "non-salt" manganese, like MnO2, which would be through a combination of base and oxidizer. You might go this route:
MnSO4 + Na2CO3 --> MnCO3(s) + Na2SO4(aq)
MnCO3 can be (should be, anyway!) good to roast in air to at least Mn2O3. It may oxidize to that or MnO2 with peroxide, but I don't know if the decomposition of peroxide is more probable (does Mn(II) catalyse peroxides?).

Anyway, once you have a manganese oxide, you need to 1. oxidize it and 2. add lots of base. The typical reaction is:
2 KOH + 2 KNO3 + 2 MnO2 = 2 K2MnO4 + 2NO + H2O (Unbalanced: N can go from +5 to 0, while Mn goes from +4 to +7, so the stoichiometry is complicated and I won't write it out.) You'll probably have an excess of base.

Finally, you need to make it into permanganate. This takes a dash of acid and oxidizer. You might use H2SO4, HCl, HOAc, H2CO3, etc., and atmospheric O2 (essentially, ignore it) for the oxidizer. Do it hot so it gets nice and concentrated, then cool it near 0C to precipitate most of the KMnO4 (some 2.5g/100ml solubility!). You may also filter it to remove excess MnO2, though you'll probably need an excellent filter to do so.

Tim

DerAlte - 6-6-2007 at 20:51

A belated comment to "mericad193724" post above - does not acetone from 'nail polish remover' also contain ethyl acetate? KMnO4 oxidizes primary alcohols to acids and secondary alcohols to ketones, if memory serves. (Organic chemists please correct if wrong). If ethyl acetate is present, isn't it likely that KMnO4 will oxidize it to acetic acid? KMnO4 goes for double C=C bonds - or alkynes - with vigor. It doesn't usually attack a single C-C band but can attack a C=O bond but usually leaves most ketones alone, and is a favorite for actually producing them.

Temperature

CRK - 3-11-2014 at 17:58

Is the temperature strictly kept below 10C to prevent a runaway reaction or does a high temperature lead to unwanted byproducts like dichloroacetone?

Edit: Sorry, I had two tabs open for science madness and realized I posted this comment on the wrong thread.

[Edited on 4-11-2014 by CRK]

AJKOER - 8-11-2014 at 09:02

Here is an interesting comment I happen to read in Atomistry (see http://oxygen.atomistry.com/chemical_preparation.html ) that appears to suggest a particular favorable temperature. To quote:

"When a mixture of manganese dioxide and sodium hydroxide is heated to dull redness in a current of air, sodium manganate is formed:

4NaOH + 2MnO2 + O2 = 2Na2MnO4 + 2H2O.

The absorption of oxygen begins at 240° C., the rate of absorption increasing with the temperature, the optimum temperature being 600° C. The product, on treatment with steam at 450° C., evolves oxygen, sodium hydroxide and manganese dioxide being regenerated:

2Na2MnO4 + 2H2O = 4NaOH + 2MnO2 + O2.

The foregoing reactions were made the basis of a commercial method for the preparation of oxygen from the air, but, owing to the short life of the solid phase, the process has not proved particularly successful. "
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Some speculation on alternate paths, starting with the cited reaction:

2Na2O2 + 2H2O = 4NaOH + O2

and noting that the products are in the first referenced equation above, upon substituting therein, we could speculate that:

2Na2O2 + 2H2O + 2MnO2 = 2Na2MnO4 + 2H2O

Or, an implied direct action of Sodium peroxide (which can be formed from the action of oxygen at around 300 C on Na2O, which can be created by heating Na2CO3 to 851C) on MnO2, or at least in the presence of water vapor:

Na2O2 + MnO2 --H2O Vapor?--) Na2MnO4

[Edit] Apparently, not too speculative as I just found this cited one pot reaction on SM to quote DerAlte at http://www.sciencemadness.org/talk/viewthread.php?tid=8480&a... :

"With Na2CO3 instead of NaOH, one might get
Na2CO3 + MnO2 + NaNO3 --> Na2MnO4 + CO2 + NaNO2
There is a very large difference between Na2CO3 in solution – which is mildly alkaline – and dissolved in NaNO3. There are no OH- ions. The presence of OH- seems essential to avoiding decomposition of magnates and also nitrates. In general the hydroxide must always be in excess of stoichiometric.
Hence I am very surprised you managed to get oxidation at 800C. which is a temperature very noticeably cherry red in daylight, and also that you managed it with sodium carbonate. Further, you probably had only somewhat impure Mn2O3 and not dioxide."

where the strongly heated Na2CO3 supplies the Na2O and the NaNO3 the required additional oxygen to form the Na2O2. Also, found another reference http://www.allreactions.com/index.php/group-1a/natrium/sodiu... citing my speculated reaction and even the required temperature, to quote:

" Na2O2 + MnO2 = Na2MnO4 (400—500° С) "

The obvious problem with this preparation is the over 800 C required temperature for the thermal decomposition of the Na2CO3 directly to Na2O. However, ball milling the Na2CO3 and MnO2 may provide a path to a much lower required temperature (400-500 C) based on mechanochemical processing (see
https://www.google.com/url?sa=t&source=web&rct=j&... ). The authors claim on page 22:

"In contrast to the carbonate decomposition in the non-milled mixture (Fig. 6d, 0 h, 400–800 °C), occurring in several steps and in a broad temperature range, which is characteristic for a physical mixture of Na2CO3 and Nb2O5 particles (Jenko, 2006), the mixture milled for only 1 hour releases CO2 in a much narrower temperature range, i.e., 400–500 °C (Fig. 6d, 1 h). We attribute this effect to the smaller particle size after 1 hour of milling, which is known to decrease considerably the decomposition temperature of Na2CO3 in the Na2CO3–Nb2O5 mixture due to reduced diffusion paths (Jenko, 2006). In comparison with the 1-hour milled sample, upon milling for 5 hours only small changes are observed in the shape of the EGA(CO2) peak (Fig. 6d, 5 h, 400–500 °C)."

As a sidebar, per Atomistry http://sodium.atomistry.com/sodium_peroxide.html one must avoid MnO2 containing any Carbon (as would be the case in a dry cell battery) as "It [Na2O2] is reduced to sodium [actually Sodium vapor per another source] by charcoal or carbides of the alkaline-earth-metals."
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Also, if one used NaO2 in place NaO, further and seemingly unsupported speculation:

NaO2 + MnO2 --??--) NaMnO4

and, even more interesting would be replacing the Sodium salts with their Potassium counterparts.

Unfortunately, per Wikipedia on Sodium superoxide (see http://en.m.wikipedia.org/wiki/NaO2 ), its preparation may be challenging, to quote:

"NaO2 is prepared by treating sodium peroxide with oxygen at high pressures:[1]

Na2O2 + O2 → 2 NaO2"

However, if one assume the the formation of NaO2 provides a possible pathway, making it in situ implies heating the mix with the addition of oxygen under pressure. This could be performed by adding an extra amount of a dry salt that produces O2 on heating (like KClO3, but certainly not in the presence of any carbon as it will sensitize the chlorate to explode) and sealing (but, with some pressure release mechanism like a cover on which a weight is placed) the reaction chamber to pressurize the liberated oxygen (where I am assuming, this will not cause the heated KClO3, for example, to explode, but I would institute safeguards assuming such a scenario for safety).

[Edited on 9-11-2014 by AJKOER]

AJKOER - 13-11-2014 at 06:01

Ok, I just found something I think is exciting. No need for Na2CO3, just NaNO2 as per this source http://books.google.com/books?id=2BpMo7HpXzIC&pg=PA150&a... pages 150 to 151, it decomposes to Na2O when heated in the open to allow venting of the NOx to Na2O and even Na2O2 in a stream of inert gas (or air) at 330 C, but under 350 C at which point the Sodium peroxide decomposes back to Na2O (see http://www.allreactions.com/index.php/group-1a/natrium/sodiu...), to quote:

" 2 Na2O + O2 = 2Na2O2 (250—350° С, р) "

Over heating would require one to reheat the mix at the indicated lower temperature range in the presence of oxygen.
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Prior referenced success using a mix of Na2CO3 and NaNO3 may actually be largely due to the presence of NaNO3 as at 600 C the equilibrium reaction:

2 NaNO3 = 2 NaNO2 + O2

and the favorable pathways via Sodium nitrite, per above, on cooling to a lower temperature.

[Edited on 14-11-2014 by AJKOER]