Sciencemadness Discussion Board - KMnO4

KMnO4

Saerynide - 10-1-2004 at 09:33

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?

vulture - 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.

Saerynide - 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

)

vulture - 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.

chemoleo - 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!

Saerynide - 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?

notagod - 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.

permanganate

chloric1 - 11-1-2004 at 14:20

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.

Mumbles - 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?

No actual Chlorine dioxide

chloric1 - 11-1-2004 at 19:07

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!

[Edited on 1/12/2004 by chloric1]

Geomancer - 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.

overseer - 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.

MORE KMnO4 FUN!!

chloric1 - 13-1-2004 at 08:22

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

. I am now heating this to drive the reaction further in hopes of obtaining MnCl2. Later! :cool:

[Edited on 1/13/2004 by chloric1]

chloric1 - 13-1-2004 at 08:27

Did my photo post???

KMnO4+HCl.JPG - 63kB

chemoleo - 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.

Saerynide - 13-1-2004 at 08:48

What can you do with the MnCl2 besides admiring its green color?

chemoleo - 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!

Of coarse!!!

chloric1 - 13-1-2004 at 18:02

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.

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]

unionised - 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.

Saerynide - 15-1-2004 at 00:50

So do the unstable chlorides break down into MnCl2?

unionised - 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]

Pyrophoric - 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

budullewraagh - 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.

JohnWW - 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.

John W.

darkflame89 - 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?

budullewraagh - 15-8-2004 at 05:32

i could imagine that it was released as either HCl or Cl2

KMnO4

mick - 17-8-2004 at 12:44

About what you can do with KMnO4
This is from a book first published in 1931 and this is from my1961 reprint
In solution it is the most powerful common oxisising agent available.
The substances oxidised include nascent hydrogen, unsaturated hydrocarbons, ammonia (only slowly in solution) and ammonium salts, nitrites, hydrogen sulpide, sulphur dioxide, sulphites, thiosulphates, phosphine, phosphites, hypophosphites, phosphorous acid, hydrochloric acid, chlorides, hydrobromic acid and hydroiodic acid and their salts, arsenites, organic matter of most kind, the lower salts of most metals, such as ferrous, cuprous, manganous, titanous, stannous salts and metals.
It has then got a page and a half on how to work out the relevant equation for what you want to do eg if you are doing it in suphuric acid, write down the oxidising equation for what you are doing and subtract this equation.
The other thing is, if its fairly pure, you can use it to find out the purity of some things that can be oxidised in water because it is self indicating, the colour changes or MnO2 is formed.
Mick

JohnWW - 17-8-2004 at 14:15

More powerful (but mostly more expensive except the last group) oxidizing agents in aqueous solution than MnO4- are:
alkali metal salts of plumbate(IV), PbO3--, and bismuthate(V), BiO3-; acid solutions of Ce(IV) salts e.g. Ce(NO3)4; and alkaline solutions of Cl2, NaOCl and NaClO2. These are capable of oxidizing Mn(II) to MnO4-, and most of them can also oxidize Fe(III) to FeO4--.

John W.

tom haggen - 17-8-2004 at 17:16

Just exactly how safe is it to mix KMn04 with concentrated sulphuric acid?

[Edited on 18-8-2004 by tom haggen]

budullewraagh - 17-8-2004 at 18:28

i'd think it would be not so dangerous as long as you added the KMnO4 slowly

Twospoons - 17-8-2004 at 19:58

If you see green crystals forming you're in trouble - I've been told they're extremely unstable.

(Oh Joy! Primary explosive in a jar of sulphuric acid!)

JohnWW - 18-8-2004 at 02:32

At low temperatures, the product would be Mn2O7, as a dark oily liquid.

John W.

unionised - 21-8-2004 at 04:44

Question "Just exactly how safe is it to mix KMn04 with concentrated sulphuric acid?"
Answer; not safe.

Hang-Man - 21-8-2004 at 05:35

I seem to remember reading somewhere that KMNO4 + H2SO4 produces O3 among other nasty things.

blazter - 21-8-2004 at 07:49

I have a friend who has a well that produces what they call "sulfur water". Based on the smell, I think it contains a small amount of H2S dissolved in it. The odor is pretty strong, like fresh shit sitting in the sun. I was talking with him and he mentioned seeing a treatment system that used a pink or purple liquid that gets injected at the water pump. Knowing that KMnO4 has that characteristic color in solution, and having a vague idea that it might oxidize the H2S I plan on making up a dilute solution and adding a couple drops to a gallon of that water. Titration should be simple if it does react since the pink will become clear.

The only possible problem is that if it isn't H2S or if the KMnO4 doesn't oxidize it. Anyone else ever experience this "sulfur water" or can verify that it is H2S? Better yet, will adding dilute KMnO4 actually do something, or do I just have a wild hair up my ass?

budullewraagh - 21-8-2004 at 09:55

last i checked, conc H2SO4+2KMnO4 -> HMnO4+K2SO4

chemoleo - 21-8-2004 at 10:04

JohnWW, you mentioned organic permanganyl esters.
Do you have more details? I always wondered whether such esters exist, and u seem to have heard about them. Please post details on a well-known/stable one if you have it (partiuclalry if it is easily made - then it probably warrants its own thread anyway) - thanks

[Edited on 21-8-2004 by chemoleo]

S.C. Wack - 21-8-2004 at 12:19

blazter, I don't know about any injection systems, but I get my KMnO4 for "greensand" filter regeneration. This sort of system takes out the H2S by precipitating it as S.

I've heard that an O2/O3 mix can be made from sulfuric and KMnO4, too. Probably not much.

The reaction between KMnO4 and H2SO4 is dependant on conditions. How concentrated, whats added to what, the temp, etc. There are several possibilities and products.

Adding hot sulfuric to the solid will just give sulfates, water, and O.

Adding the solid to conc. H2SO4 can give Mn2O7. This may not be something that you want to make much of.

JohnWW - 21-8-2004 at 13:03

Quote:

Originally posted by chemoleo
JohnWW, you mentioned organic permanganyl esters.
Do you have more details? I always wondered whether such esters exist, and u seem to have heard about them. Please post details on a well-known/stable one if you have it (partiuclalry if it is easily made - then it probably warrants its own thread anyway) - thanks

[Edited on 21-8-2004 by chemoleo]

I think I said that organic permanganate esters may possibly exist, not that they do for a fact. Like chromate, perchlorate and nitrate esters, they would be highly explosive, if they could be made. It would be possible to make them only at low temperatures, by the reaction of alcohols with Mn2O7 or with HMnO4, plus a small amount of H2SO4. Like those other esters, they would depend for their formation on their auto-decomposition into explosion products being kinetically slow at low temperatures, although highly thermodynamically favorable.

John W.