Sciencemadness Discussion Board - Manganese Dioxide from Manganese Sulfate

Manganese Dioxide from Manganese Sulfate

toothpick93 - 10-5-2014 at 00:33

Last night I watched NurdRage's video on making MnO₂ via MnSO₄ and KHSO₅ which you can substitute for hydrogen peroxide to get roughly the same reaction. Problem is I dont feel like buying 1 kg of Potassium Peroxymonosulfate just to use for this so I was thinking "is there any other chemical with the extra oxygen that could get the same results when I remembered Sodium Percarbonate used for laundry (much easier to get).

I dissolved the 2 chemicals in separate beakers and then mixed the 2 together thinking not much would happen. Straight away a brown-ish precipitate formed with some bubbling also. After about an hour of letting it settle the precipitate is a light brown which looks more like Manganese Carbonate.

Does anyone have the equation to this reaction?

gdflp - 10-5-2014 at 04:40

Your precipitate is most likely manganese dioxide. All manganese(II) salts are a white to light pink color, google manganese carbonate. The equation for your reaction is probably going to roughly be Na₂CO₃ + H₂O₂ + MnSO₄ -> Na₂SO₄ + H₂O + CO₂ + MnO₂

[Edited on 2-19-2018 by gdflp]

toothpick93 - 10-5-2014 at 05:00

But Manganese(II) Dioxide is black. Would the colour change when it gets filtered and dried?

macckone - 10-5-2014 at 05:24

This is going to most likely be a classic precipitation reaction
with the MnCO₃ partially decomposing ie:

3Na₂CO₃*3H₂O₂ + 3MnSO₄ -> 3MnCO₃ + 3Na₂SO₄ + 9H₂O₂

Then the hydrogen peroxide and the manganese carbonate will
react to some extent yielding oxygen, carbon dioxide and
manganese oxide.

This reaction will not yield a pure oxide or a pure carbonate.

Edit by moderator: Removed erroneous quote

[Edited on 2-19-2018 by gdflp]

blogfast25 - 10-5-2014 at 05:42

Quote: Originally posted by toothpick93

But Manganese(II) Dioxide is black. Would the colour change when it gets filtered and dried?

The colour of the end product does indeed depend somewhat on specific preparation conditions and ranges from dark chocolate brown to almost black.

Mn(II) salts are so easily oxidised to MnO2 that wasting an expensive oxidiser on it is dumb.

Dissolve the MnSO4 in water and precipitate it with an alkali as Mn(OH)2, an off white gel. On filtering and drying the product will be air oxidised acc.:

Mn(OH)2 + 1/2 O2 === > MnO2 + H2O

If you are going to use an oxidiser, use bleach (sodium hypochlorite solution).

[Edited on 10-5-2014 by blogfast25]

[Edited on 10-5-2014 by blogfast25]

Texium - 10-5-2014 at 05:44

Just wondering, where did you get your manganese sulfate?

toothpick93 - 10-5-2014 at 07:12

I picked up 500 grams of manganese sulfate as a water soluble fertilizer
http://www.manutec.com.au/default.asp?pageid=10&template...

Texium - 10-5-2014 at 07:56

Oh, nice! It looks like they sell quite a few useful chemicals as fertilizer. Too bad they're only in Australia. It seems like all of the fertilizers for sale in the US are mixtures of a ton of different things, not of much use for chemistry.

toothpick93 - 10-5-2014 at 08:05

Here in australia its hard to get your hands on useful chemicals, Looking for the is the key factor about finding things that may come in useful. I buy my Sulfur, Copper Sulfate, manganese sulfate, Magnesium nitrate from these guys for a cheap price. About $10 for 500 grams

Texium - 10-5-2014 at 08:23

That's nice, I bought copper sulfate at the hardware store, I think it was $12 for 2 pounds as root killer, but they don't have manganese sulfate, magnesium nitrate, or even sulfur. They did have potassium nitrate though, which I know can be a bit difficult to come by in other countries. I guess different places have different chemicals that are easier to come by than they are in other places.

toothpick93 - 10-5-2014 at 09:36

Yeah I've never came across potassium nitrate as a fertilizer but i know in USA they have it. Only nitrates ive came across OTC is Ammonium nitrate in ice packs, Magnesium Nitrate in fertilizers. Its very hard finding a good quality fertilizer with little to no contaminates

Texium - 10-5-2014 at 09:41

Yep, it's usually not worth going through the trouble to separate them out.

toothpick93 - 10-5-2014 at 09:49

Yeah, too much hassle, you dont always get a good yield so you end up waisting a lot of product and money

Jesse Pinkman - 10-5-2014 at 10:01

Much easier way to obtain manganese dioxide is from alkaline/heavy duty batteries, preferable D size.

The reverse reaction MnO2 -> MnSO4 is more troublesome.

gsd - 10-5-2014 at 10:07

Quote: Originally posted by toothpick93

But Manganese(II) Dioxide is black. Would the colour change when it gets filtered and dried?

What is Manganese(II) Dioxide ?

It is either Manganese(II) Oxide = MnO
or
Manganese(IV) Oxide = MnO2

gsd

Texium - 10-5-2014 at 10:08

That is what I use. I'm making some manganese chloride with it right now. It takes about 24 hours to complete the reaction and fumigates my back yard with chlorine.

bismuthate - 10-5-2014 at 10:11

Well, I believe that bleach would react with Mn (II) salts to form MnO2. Here is a link to an experiment where this occurs.
http://woelen.homescience.net/science/chem/exps/bleach/index...
So, that is a pretty cheap way to make some MnO2.
Also the MnO2 in batteries is really impure so it might not work for some things.

Jesse Pinkman - 10-5-2014 at 10:29

If Very pure MnO2 is desired, it can be obtained from MnSO4 by the following method:

"Preparation of gamma-Manganese Dioxide
Manganese (II) sulphate monohydrate (140 g) was dissolved in 2.66 liters of water and heated to 60°C. Potassium permanganate (97.3 g) in 1.85 liter of water was added over a period of 15 minutes and stirred at 60°C for 1 hour, until manganese dioxide precipitated out. The reaction mixture was filtered and the residue was washed with deionised water until no sulphate ion was present. The solid was dried under suction for 2 hours followed by drying at 70°C under vacuum to a constant weight (about 8 days) to give 115 g of a dark brown powder."

Brain&Force - 10-5-2014 at 10:55

Quote: Originally posted by bismuthate

Well, I believe that bleach would react with Mn (II) salts to form MnO2. Here is a link to an experiment where this occurs.
http://woelen.homescience.net/science/chem/exps/bleach/index...
So, that is a pretty cheap way to make some MnO2.
Also the MnO2 in batteries is really impure so it might not work for some things.

Adding some hydrogen peroxide will convert the permanganate back to manganese dioxide in a neutral or basic solution.

blogfast25 - 10-5-2014 at 12:33

Quote: Originally posted by Jesse Pinkman

Much easier way to obtain manganese dioxide is from alkaline/heavy duty batteries, preferable D size.

The reverse reaction MnO2 -> MnSO4 is more troublesome.

You obviously haven't ever done EITHER.

MnO2 from batteries is seriously messy, in part because the battery dry electrolyte contains about 30 % very fine graphite, not mention quite a few other ingredients. Without serious purification that stuff is good for NOTHING.

[Edited on 10-5-2014 by blogfast25]

blogfast25 - 10-5-2014 at 12:37

Quote: Originally posted by Brain&Force

Adding some hydrogen peroxide will convert the permanganate back to manganese dioxide in a neutral or basic solution.

Woelen created conditions favourable to the formation of permanganate but trust me: if you go easy on the bleach you won't produce any permanganate at all. The colour is deceptive because of permanganate's insanely high colour intensity: in reality there's almost no MnO4-, perhaps 0.0001 M or so.

[Edited on 10-5-2014 by blogfast25]

Jesse Pinkman - 11-5-2014 at 04:22

I have disassembled many times heavy duty batteries. You are right the procedure is really messy, but as seen from NurdRage's videos, he has obtained reasonably pure MnO2 from heavy duty batteries ( http://www.youtube.com/watch?v=knc1lSupAwQ ) . The carbon is mostly in form of a graphite rod. The MnO2 powder has impurities like NH4Cl and soluble iron salts, which can be washed with water (many videos on YouTube).

From this battery obtained MnO2 he has later produced MnSO4 (with nice pink color) and KMnO4.

Simple further oxidation of MnO2 with bleach will not yield much KMnO4, only in trace quantities. Better way to permanganate is to produce first KClO3 from bleach/bleaching powder then to follow the directions from VersuchsChemie : http://www.versuchschemie.de/topic,10934,-Kaliumpermanganat.... .

blogfast25 - 11-5-2014 at 05:05

Quote: Originally posted by Jesse Pinkman

I have disassembled many times heavy duty batteries. You are right the procedure is really messy, but as seen from NurdRage's videos, he has obtained reasonably pure MnO2 from heavy duty batteries

From this battery obtained MnO2 he has later produced MnSO4 (with nice pink color) and KMnO4.

The video says NOTHING about the purity of the product, nor is it guaranteed he used THAT SAME product for the production of MnSO4. But even if he did, the reduction of the MnO2 with sulphuric acid and oxalic acid would leave behind the carbon, which is there as a conductor and does not react. I've made MnSO4 from battery crud that way and I can assure you there was plenty graphite in the starting product.

I worked on a commercial project (as a consultant/analyst) for the recycling of battery electrolyte and on one large batch found the Mn content to be about 27 %. That is much, much lower than the Mn content of MnO2. And that there's a lot of graphite is shown by just how staining this material is.

Alkali batteries may contain less graphite, as carbon reacts with strong alkali. All non alkaline batteries I've taken apart contained much graphite.

He states quite clearly it's NOT an alkaline battery.

Take some of your own battery 'MnO2' and dissolve it in strong HCl and see what happens.

From Wiki on alkaline batteries:

"A cylindrical cell is contained in a drawn steel can, which is the cathode connection. The positive electrode mixture is a compressed paste of manganese dioxide with carbon powder added for increased conductivity."

So even those contain graphite.

[Edited on 11-5-2014 by blogfast25]

Zyklon-A - 11-5-2014 at 05:32

Quote: Originally posted by blogfast25

Well yeah, but purification is rather easy. The first time I did it, I simply washed the black crap with water, because the salts in there (zinc chloride, ammonium chloride ect.) are very soluble. After it dried it, I heated the stuff with a torch in air to burn off most of the carbon. I don't remember exactly what the yields were, but I weighed it before and after, and it certainly lost mass as the carbon was oxidized by air. I'm not sure how pure it it, and I'm sure it still contained iron contaminates. Either way, I think it's pure enough for several applications.

[Edited on 11-5-2014 by Zyklonb]

blogfast25 - 11-5-2014 at 06:08

Zb:

Considering you can buy decent MnO2 for chips, and even decent MnSO4/MnCl2, it's more a 'rite of passage' for backyard chemists, if you ask me.

Often overlooked is that in spent batteries the Mn is mostly present as Mn2O3, not MnO2. Otherwise the battery (or cell) could not have delivered its chemical energy.

Also, there's almost always iron present too: and removing that fully is NOT so easy. It can't be done without dissolving (thus reducing) the Mn oxide.

A full purification to 99+ % purity is quite a kerfuffle.

[Edited on 11-5-2014 by blogfast25]

Texium - 11-5-2014 at 06:13

Well, if you only use unused batteries, then you get only MnO₂, and from there it's simple to clean. I didn't even bother burning off the graphite for what I was doing, because it remained unreacted in the reaction and was easily filterable afterwards.

Afterwards I removed the iron using the hydroxide method mentioned in the Manganese Chloride Crystals thread and used by NurdRage. My end goal was to isolate manganese, and it seemed to work very well. There is no doubt that there was still some iron left, but that doesn't really bother me a whole lot.

[Edited on 5-11-2014 by zts16]

Arthur Dent - 11-5-2014 at 08:13

It is indeed pretty hard to purify and clean-up MnO2 from batteries, it's a mess from beginning to end, but my favorite trick (as mentioned above) is that after extracting Zinc-Carbon battery material, thoroughly rinse the black paste to remove soluble Zinc salts, then mix with hydrochloric acid to make some MnCl2. Lots of nasty gases will be emitted, beware!

Filter the stuff to remove the remaining insoluble impurities, you'll get a dark tea-colored solution, that you can reduce a bit with gentle heating (it will actually get lighter in color), and and when it starts crystallizing, you'll have a fairly clean Mn salt, from which you can make whatever you need.

Even then there are probably still trace impurities of zinc and iron, but it's okay for most purposes...

Robert

blogfast25 - 11-5-2014 at 09:10

Quote: Originally posted by zts16

Well, if you only use unused batteries, then you get only MnO₂, and from there it's simple to clean. I didn't even bother burning off the graphite for what I was doing, because it remained unreacted in the reaction and was easily filterable afterwards.

Afterwards I removed the iron using the hydroxide method mentioned in the Manganese Chloride Crystals thread and used by NurdRage. My end goal was to isolate manganese, and it seemed to work very well. There is no doubt that there was still some iron left, but that doesn't really bother me a whole lot.

[Edited on 5-11-2014 by zts16]

If you're going to dissolve the Mn oxide (to remove the iron), then it doesn't matter one iota whether you start from MnO2 or Mn2O3 because the latter is reduced to Mn(II) anyway, just as MnO2. Might as well use spent batteries, in that case...

S.C. Wack - 11-5-2014 at 09:55

Why not Mn2O3 from percarbonate?

I don't see any reason to slap down battery MnO2 users...it really isn't a big deal to dissolve and convert to CMD, and the mess is limited when most everything is done under water with gloves...and, I doubt much Fe is found in batteries these days as I've said before, certainly not alkaline batteries. With zinc batteries, you get bonus zinc to melt down, and the rods I guess if anyone has found them useful...
http://www.tronox.com/products/electrolytic-manganese-dioxid...

blogfast25 - 11-5-2014 at 10:03

Mn2O3 from percarbonate? Explain.

There's no 'slapping down' here.

MnO2 and Mn(II) salts are very cheap and not hard to get. Messing with batteries can be fun but hardly economical. Or easy, if you want a decent end product. Many battery peddlers are happy when it's black and tests positive for Mn. But that's too easy.

CMD? What's that?

Under water? Separating the graphite out 100 % is not easy and seriously messy.

Iron: I bet $100 (easy) that I'll find significant amounts of Fe in any randomly selected non-alkaline battery electrolyte. It's always been the outcome in my case.

[Edited on 11-5-2014 by blogfast25]

[Edited on 11-5-2014 by blogfast25]

S.C. Wack - 11-5-2014 at 10:34

Why not Mn2O3? Explain.

Separation of the solution from graphite is achieved by the high-tech method of decantation. Things may be possible that you aren't aware of.

I'd add (again) that the zinc in certain brands of new alkaline batteries will give shiny zinc powder with minimal washing, though others disagree with that too.

What's CMD? Well, the MD is MnO2...
http://www.sciencemadness.org/talk/viewthread.php?tid=26790&...

[Edited on 11-5-2014 by S.C. Wack]

blogfast25 - 11-5-2014 at 10:39

Quote: Originally posted by S.C. Wack

Why not Mn2O3? Explain.

Separation of the solution from graphite is achieved by the high-tech method of decantation. Things may be possible that you aren't aware of.

Stop trying to read my mind. Yes, I'm aware of various methods of floatation/decantation than can be used to separate graphite from something else. Do you think any of the battery cutters here have actually tried them? High tech methods?

Now kindly tell me about the connection between Mn2O3 and percarbonate, because that I don't see.

Yes, MD for manganese dioxide. E for electrolytic. C for ? Aaahh, chemical, got it. But ammonium peroxodisulphate? Don't make me giggle.

[Edited on 11-5-2014 by blogfast25]

bismuthate - 11-5-2014 at 14:53

Ok so this may be a stupid question, but could one remove iron and zinc from the battery mixture by treating it with sulphuric acid and filtering it?

Jesse Pinkman - 11-5-2014 at 14:57

Of course much more easier and economical way is to obtain directly the MnO2/MnSO4 from a chemical supply store.

MnSO4 can be mixed with ammonium sulfate and 60% H2SO4, then electrolysed to (III) oxidation state to produce benzaldehyde from toluene. (Hilski's thread: http://www.sciencemadness.org/talk/viewthread.php?tid=6882 )

Or MnSO4 can be converted to the carbonate and then to manganese(II) acetate, which on oxidation will give the more precious manganese(III) acetate for small scale production of P2P from benzene (Leukemia) and acetone

S.C. Wack - 11-5-2014 at 16:49

Quote: Originally posted by blogfast25

C for ? Aaahh, chemical, got it. But ammonium peroxodisulphate? Don't make me giggle.

You not knowing CMD was giggly...I found their whole process and the claim of a particularly active product interesting. Ammonium persulfate is $26.09 a kilo to your door here. Everyone knows about using KMnO4 instead already.

Those interested in organic oxidations are more interested in CMD than the cheap product sold by pottery suppliers.

blogfast25 - 12-5-2014 at 04:43

Quote: Originally posted by S.C. Wack

Using abbreviations and then giggling at those who don't know them is silly.

There is no a priori reason to suspect that other oxidisers will yield a less active product, for organic oxidations or for whatnot. At that price of oxidiser reclaiming MnO2 from spent batteries can not be economical at all. Form what I understand from a company that produces spent battery electrolyte, finding an economical end use for it is a big problem.

High grade MnO2 ('CMD' or 'EMD' to use your false dichotomy) is inexpensive too.

And I still don't know anything about Mn2O3 and percarbonate... Oh well.

[Edited on 12-5-2014 by blogfast25]