| Pages:
1
2
3 |
MephistosMinion
Harmless
Posts: 24
Registered: 16-1-2005
Member Is Offline
|
|
KMnO4 synthesis
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
International Hazard
Posts: 3227
Registered: 13-7-2003
Location: Wisconsin
Member Is Offline
Mood: Rock n' Roll
|
|
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
Hazard to Others
Posts: 491
Registered: 7-2-2004
Location: Mediterranean
Member Is Offline
Mood: Quantized
|
|
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]
Theory guides, experiment decides.
|
|
|
Blackout
Harmless
Posts: 25
Registered: 26-11-2004
Location: Canada
Member Is Offline
Mood: No Mood
|
|
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.
\"Si vis pacem, para bellum.\"
\"If you wish for peace, prepare for war.\"
|
|
|
neutrino
International Hazard
Posts: 1583
Registered: 20-8-2004
Location: USA
Member Is Offline
Mood: oscillating
|
|
HMnO<sub>4</sub> will dehydrate to an oily liquid, Mn<sub>2</sub>O<sub>7</sub>. This won’t react with KCl.
|
|
|
The_Davster
A pnictogen
Posts: 2861
Registered: 18-11-2003
Member Is Offline
Mood: .
|
|
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
International Hazard
Posts: 1583
Registered: 20-8-2004
Location: USA
Member Is Offline
Mood: oscillating
|
|
On second thought, it might not dehydrate so easily without the dehydrating power of H<sub>2</sub>SO<sub>4</sub>, as it does
in the preparation of Mn<sub>2</sub>O<sub>7</sub>. Still, I doubt you'd get a permanganate precipitate, as the protons
liberated when the permanganate precipitates would probably react back to form HMnO<sub>4</sub>. Remember that weak acid + salt doesn’t
give you a strong acid (HCl) and a salt.
|
|
|
The_Davster
A pnictogen
Posts: 2861
Registered: 18-11-2003
Member Is Offline
Mood: .
|
|
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
International Hazard
Posts: 5102
Registered: 1-11-2003
Location: UK
Member Is Offline
Mood: No Mood
|
|
You might want to look into the properties of Mn2O7 before you think about making it.
|
|
|
Cesium Fluoride
Harmless
Posts: 48
Registered: 26-1-2005
Member Is Offline
Mood: Cupric
|
|
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
International Hazard
Posts: 3873
Registered: 21-7-2006
Member Is Offline
Mood: No Mood
|
|
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
Harmless
Posts: 48
Registered: 26-1-2005
Member Is Offline
Mood: Cupric
|
|
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
International Hazard
Posts: 3873
Registered: 21-7-2006
Member Is Offline
Mood: No Mood
|
|
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
Harmless
Posts: 48
Registered: 26-1-2005
Member Is Offline
Mood: Cupric
|
|
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
Harmless
Posts: 48
Registered: 26-1-2005
Member Is Offline
Mood: Cupric
|
|
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
Post Harlot
Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline
Mood: informative
|
|
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
Harmless
Posts: 48
Registered: 26-1-2005
Member Is Offline
Mood: Cupric
|
|
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
Post Harlot
Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline
Mood: informative
|
|
Oops, that should read Na2 or K2, since that's what's made in the fusion.
Tim
|
|
|
Cesium Fluoride
Harmless
Posts: 48
Registered: 26-1-2005
Member Is Offline
Mood: Cupric
|
|
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
Post Harlot
Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline
Mood: informative
|
|
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
Harmless
Posts: 48
Registered: 26-1-2005
Member Is Offline
Mood: Cupric
|
|
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
National Hazard
Posts: 596
Registered: 29-9-2004
Location: Netherlands
Member Is Offline
Mood: Whatever...
|
|
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.
#261501 +(11351)- [X]
the \"bishop\" came to our church today
he was a fucken impostor
never once moved diagonally
courtesy of bash
|
|
|
Cesium Fluoride
Harmless
Posts: 48
Registered: 26-1-2005
Member Is Offline
Mood: Cupric
|
|
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
National Hazard
Posts: 525
Registered: 17-12-2003
Location: Midwest
Member Is Offline
|
|
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
Harmless
Posts: 48
Registered: 26-1-2005
Member Is Offline
Mood: Cupric
|
|
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!
|
|
|
| Pages:
1
2
3 |
![[*]](images/xpblue/default_icon.gif){kind=icon}
