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Saerynide
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[*] posted on 10-1-2004 at 09:33
KMnO4


I bought 100 g of KMnO4 crystals (just because theyre hard to come by where I live) and now I am wondering what to do with them.

Ive scoured the forum and internet, but everything that has anything to do with potassium permanganate is related to explosives will likely get me killed.

So, is there anything interesting that can be made requiring KMnO4 thats not related to organic chemistry and where I wont be blown to bits?
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[*] posted on 10-1-2004 at 10:14


I was going to suggest benzoic acid, but since you ruled out organics...:(

KMnO4 doesn't have much use in inorganic chemistry, except for liberating chlorine gas or liberating oxygen gas from H2O2.

KMnO4 pyrotechnic mixtures are quit safe if you keep them away from moisture and acids and don't store them for too long.




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[*] posted on 10-1-2004 at 10:28


Hmm... What can you do with benzoic acid? More over, what is it? (Now you see why I preferred something non-organic chem related cause I know next to nothing about it :( )
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[*] posted on 10-1-2004 at 10:39


Benzoic acid is C6H5-COOH, tolueneacid so to say. It can be used to make sodiumbenzoate, a component of whistle mix.

It's also an interesting reagent to make benzene and other aromatic compounds.




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[*] posted on 10-1-2004 at 10:42


Try the old glycerol / KMnO4 trick and hope that it ignites :) - and it wont blow you to bits dont worry... unless you use a few tons ;)
Mix it with sulphur/iron, it burns nicely but not particularly vigorously.
Else, are you interested in inorg. chemistry? You could play with the different oxidation states of manganese, and the different colours thereof.
You could use it to produce MnO2, which is mixed with Al, to make lovely elemental manganese!
Going slightly organic, you could make acetic acid from ethanol.... there is loads to do!




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Saerynide
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[*] posted on 11-1-2004 at 01:42


If I oxidzise ethanol to acetic acid, wont I need to use H2SO4 for that? Mixing KMnO4 with H2SO4 is just suicidal...

Is there another readily available acid that works instead?
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[*] posted on 11-1-2004 at 04:43


It's not dangerous if you don't use concentrated acid. Dilute it first and drop the etanol in slowly.
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[*] posted on 11-1-2004 at 14:20
permanganate


I have used KMnO4 in 30 to 40% H2SO4 to generate Br2 from NaBr. I also have mixed 99% concentrated H2SO4 with solid permanganate(in tiny amounts) and got some neat looking mixtures of Mn2O7 with the dioxide. it gave off some red vapors that smelled like a mixture of ozone and chlorine dioxide.



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[*] posted on 11-1-2004 at 16:14


Hmmmm, ozone I understand. Its that Chlorine Dioxide I don't understand. There's no chlorine to combine with. Was there anything else in the flask, or something else you added?
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chloric1
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[*] posted on 11-1-2004 at 19:07
No actual Chlorine dioxide


Was not referring to Chlorine dioxide in a literal sense but only making a reference to the odor. Although the permanganic acid leans further towards an ozone smell than a chlorine dioxide smell. Ozone is pungent-sweet while chlorine dioxide is more like musty- pungent. All of these are VERY dangerous so be cautious when smelling!:D

[Edited on 1/12/2004 by chloric1]




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[*] posted on 11-1-2004 at 20:44


Regarding the old permanganate/glycerol trick, I find it fairly reliable if you add a small amount of water to the glycerol. Once I tried to produce a colored flame by adding PVC filings and CuSO4 to the permanganate. Thing didn't ignite properly.
Note that acrolein is emitted by hot glycerol, so the smoke from ignition is somewhat unpleasant.
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[*] posted on 13-1-2004 at 04:22


Regarding the pyrotechnic mixtures with KMnO4, I found some of them to be somewhat hygroscopic and prone to self-ignition. The well-known example with glycerine, or glycol that works just as fine, can prove that point. If You'd mix KMnO4 with starch (or worse yet, starch/sugar), the mixture would pull moisture from the air and would gently heat as KMnO4 starts to oxidize the sugars. However, if it pulls enough moisture, the reaction picks up speed in much the same way as with glycerine/glycol, and the mixture may start burning. This can even take a few days.

Regarding the inorganic reactions, obtaining green alkaline manganates K2MnO4 or Na2MnO4 can be interesting. Manganese compounds, when melted with alkaline hydroxides or carbonates in presence of oxidizers such as KNO3 or KClO3, yield manganates. Melting solid NaOH should be no problem with even a simple alcohol burner and a test tube. As far as I can remember, it should be enough to melt KMnO4 with NaOH to obtain the reaction. After adding some water to dissolve the product, the water is slowly evaporated and green, rhombic manganate crystals appear. The manganates are stable only in alkaline solutions, whereas neutral or acid solutions oxidize them back into permanganates, which is indicated by a change in color.

Once KMnO4 is converted into some other manganese compound, it may be interesting to attempt to convert it back and observe the color change. Say, Pb-dioxide, on boiling with conc. HNO3, oxidizes all manganese compounds into permanganates. You put a tiny amount of a manganese compound in a test tube, add 1-2 ml of conc. HNO3 and a minute amount of lead dioxide. Then the mixture is well boiled, and water is added down the walls of a test tube so the mixture does not get agitated. As the Pb-dioxide precipitates, a well-known purple color is observed. This reaction can be used to identify manganese in its compounds.
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[*] posted on 13-1-2004 at 08:22
MORE KMnO4 FUN!!


Last night I decided to take it upon myself to dissolve some KMnO4 in some concentrated Hydrochloric since i have not done this in a while. I used about 5 or so grams and just less than 200 ml of HCl but I want to at least have the 2 moles KMnO4 to 16 Moles HCl for complete reduction to the Manganese(II) oxidation state. It was late and I was tired so I let the reaction simmer over night. It was close to zero degrees Centrigrade or less and I awoke this morning to find a curious result. I had a very dark brown liquid suggesting of Manganese Dioxide but there where no solids deposited. I know it could not have been MnCl4 or MnCl3 because i have never seen any literature that would suggest there existence. I have noticed in the past that MnO2 dissolves and decomposes HCl usually only when heated. So somehow I have a susupension of MnO2 in HCl or maybe a complex mixture of products. I offer this photo for your viewing pleasure I know it is not great but I aint a photograher either:P. I am now heating this to drive the reaction further in hopes of obtaining MnCl2. Later!:cool:

[Edited on 1/13/2004 by chloric1]




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[*] posted on 13-1-2004 at 08:27


Did my photo post???

KMnO4+HCl.JPG - 63kB




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[*] posted on 13-1-2004 at 08:30


I think if you dilute it you will get a nicely dark green colour (MnCl2). Did you try that? It can't be MnO2 or something, because that is reduced too to form MnCl2, in the presence of HCl.



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[*] posted on 13-1-2004 at 08:48


What can you do with the MnCl2 besides admiring its green color? :D
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[*] posted on 13-1-2004 at 09:12


You could oxidise it back to MnO2 with H2O2, which stains your glassware nicely (once you got rid of free HCl). With MnO2, you could do a little thermite reaction, and make lovely elemental manganese. Or, you could make K2MnO4 (not KMnO4). YOu could add complexing agents to MnCl2, and watch if something happens (and isolate any potential products). Such as hydrazine, ammonia, etc... or embark on making double salts (if they exist, in both cases). I can get you details if you actually ARE going to try some of that!



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[*] posted on 13-1-2004 at 18:02
Of coarse!!!


After a little heating it became olive drab and then I diluted with 400 ml water to obtain a clear solution with a pinkish tint. I realized that, as with so many other transition elements, the chloride solutions are capable of being intensely colored because of complexes. :o Man I need to get it together. Actually, my goal is to use manganese(II) ions to reduce permanganate to Manganese dioxide. I have read that there is an alpha and beta form of this oxide and I want to play around with many variations of MnO2. Some can even be vigorously stirred with active organics(terpenes etc) for about 3 or 4 days and convert them to aldahydes! This has said to work with Benzyl alcohol also! I can get you the specifics if you ask nicely:)

[Edited on 1/14/2004 by chloric1]

[Edited on 1/14/2004 by chloric1]

[Edited on 1/14/2004 by chloric1]




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[*] posted on 14-1-2004 at 13:13


Manganese (II) chloride is pink in solution and the hydrated crystals are pink too.
The green colour is due to chlorides/ chloro complexes of higher oxidation states. These are not stable.
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[*] posted on 15-1-2004 at 00:50


So do the unstable chlorides break down into MnCl2?
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[*] posted on 16-1-2004 at 12:53


Yes, they give MnCl2

Is the group of chemicals you work with the non-esistent ones like those 2?

[Edited on 16-1-2004 by unionised]
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[*] posted on 14-8-2004 at 07:48


chloric1 i have produced a similar looking solution to yours by adding excess HCl to MnO2 and heating. After the reaction subsided the solution looked very similar to the one in your post. I filtered it, producing a clear solution - can't remember exactly what colour. I then boiled it down - forming a hydrated yellow salt which upon further heating formed and apparently anyhydrous yellow-orange-possibly red salt (i can't quite remember, this was some time ago). I left the product out in the open for some time and it appeared to gain it's water of hydration back again forming a yellow salt. Could this have possibly been a Mn(III) or Mn(IV) salt?

I might perform this experiment again due to my sketchy recolection - this was some 2 years ago! And as the colour i recall seeing might've been due to contamination of my reactants :o.
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[*] posted on 14-8-2004 at 07:52


if youre looking for something to do, add the KMnO4 to conc H2SO4 slowly. you'll get conc permanganic acid and potassium sulfate. be very careful with the acid as it will explode on contact with organics and metals.



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[*] posted on 14-8-2004 at 13:14


I have read somewhere that Mn2O7 can be obtained, as a dark oily liquid, by adding supercooled H2SO4 slowly to supercooled KMnO4 solution, the mixture being kept cold in the process. On dilution of the mixture with water, it forms a strongly acid solution of HMnO4.

The only other stable compound of Mn(VII), other than MnO4-, Mn2O7, and HMnO4, except possibly for dangerously explosive covalent permanganate organic esters, is permanganyl fluoride, MnO3F. I think it is formed when HF is added in the above reaction.

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[*] posted on 15-8-2004 at 01:12


Chemoleo, when you add H2O2 to MnCl2, you said that MnO2 was regenerated, then, where did the chloride go?

Regarding the thermite reaction that gave manganese, do you use aluminium as well?




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