Sciencemadness Discussion Board - Stuborn MnO2

Stuborn MnO2

YT2095 - 7-1-2007 at 09:30

with the exception of HCl, non of the usual acids will attack it in any appreciable way.

HCl at best will dissolve a little bit of it leaving a brownish soln with loads of unreacted HCl and MnO2.

my goal is make a sulphate, a nitrate and a chloride of this metal.

at the moment the only thing that springs to mind is a Thremit reaction to obtain the metal and start from there.

but it`s quite crude and not very elegant, I`ve no objection to doing this if it`s the only way, but I`de rather not.

any ideas?

kaviaari - 7-1-2007 at 10:00

You can find some info from this thread http://www.sciencemadness.org/talk/viewthread.php?tid=6777#p...

YT2095 - 7-1-2007 at 10:34

Excellent thanks, I must have skipped right past that in the search.

the direct SO2 reaction looks kinda sexy

I think for Now, I`ll put some in the crucible and roast it down to MnO and try from there, maybe give the H2O2 and H2SO4 a shot too in the morning.

woelen - 7-1-2007 at 13:51

MnO2 comes in different forms. I think you have the crystalline calcined stuff. I also have some of it and this stuff SUCKS

. It is rather inert and very hard to get in solution. I could get a small amount in solution, by boiling it with concentrated HCl for 10 minutes or so, but this is horrible. Lots of HCl-fumes.

There also is a very fine amorphous powder, and that material is much more reactive. If added to HCl, it quickly dissolves, giving a very dark brown/green solution and you quickly smell the chlorine coming from the solution.

12AX7 - 7-1-2007 at 14:19

Heh, if you wanted to be really roundabout about it you could fuse it with soda and decompose it in water (as sodium manganate is rather unstable in solution).

HCl tends to reduce and complex(?) with MnO2 so does a fair job, a far sight greater than sulfate, which hardly dissolves Mn(III) anyway (very low pH solutions, rapid hydrolysis). Whatever you do, you need a reducing agent in there.

Tim

YT2095 - 8-1-2007 at 09:38

Woelen, I think it is exactly what you`ve described unfortunately

in addition, according to mt book MnO2 decomposes at 808k I roasted this stuff in the crucuble to a Much higher temp (forced air Propane) and no Green color at all. not only that but with the crucible lid taken off there was even the odd spark/flash taking place inside it!

I think this may be a Bad Product

on microscope examination it is indeed crystaline too for the most part, the Thermit reaction is becoming More and More tempting as time passes.
if I have to go this route, does anyone know how "Violent" an Alu MnO2 reaction is?
I don`t want some exotic type sh!t that`s going ruin the product or create annoyance/disturbance.

the H2SO4+H2O2+MnO2 reaction Does work exceptionaly well, there is little MnO2 left undisolved, there is no need for Conc reagents either (40% sulphuric to MnO2 then add 9% H2O2 dropwise).
seems a real waste of good O2 though, I think I`ll get the gas syringe out when I scale it up and collect it for a rainy day

and that direct SO2 conversion is a Must Try! when I`m free for a few hours

edit: Whoa! a surprise (unexpected) is that not only does the H2O2 trick work with Sulphuric , but also worked in the HCl batch that wasn`t going anywhere either, it`s all dissolved!????

[Edited on 8-1-2007 by YT2095]

kaviaari - 8-1-2007 at 10:06

I am dealing with the same problem. I recieved about 200g of very old MnO2 from school. I assume that the reagent and bottle are from 1950's. I was thinking about converting it to manganese carbonate. 15ml of cons. hydrochlorid acid was poured on one gram of MnO2. The mixture was heated in water bath. Chlorine was released and I left the beaker sit for a week. Even after that there still was some unreacted MnO2, so I added more acid an heated it again. Didn't help at all, so I just filterated off the excess MnO2.

I poured saturated solution of sodium bicarb to acid-MnCl2-mixture. Carbon dioxide was released big time so I did the adding very slowly. Finally brown manganese carbonate started precipitating. Here is the picture of the crude product (it's wet too

).

[Edited on 8-1-2007 by kaviaari]

YT2095 - 8-1-2007 at 10:32

interesting notes, you were actualy lucky to even get Cl2 gas, no such thing happened here at all with this stuff, certainly not at any noticable level (and I did check for it), pity you don`t still have this experiment on the go, I`de recommend adding a little H2O2 to it, it`s cleared mine up just great, although unlike the H2SO4, the HCl is a more foamy reaction I found.
so dropwise and stir and seems to be the best way for HCl.
although I really cannot comment as to the color, since it`s no longer Brown in soln but closer to clear yellow, albeit a Dull clear yellow, not at all Vivid.

[Edited on 8-1-2007 by YT2095]

kaviaari - 8-1-2007 at 11:20

Quote:

Originally posted by YT2095
interesting notes, you were actualy lucky to even get Cl2 gas, no such thing happened here at all with this stuff, certainly not at any noticable level (and I did check for it), pity you don`t still have this experiment on the go, I`de recommend adding a little H2O2 to it, it`s cleared mine up just great, although unlike the H2SO4, the HCl is a more foamy reaction I found.
so dropwise and stir and seems to be the best way for HCl.
although I really cannot comment as to the color, since it`s no longer Brown in soln but closer to clear yellow, albeit a Dull clear yellow, not at all Vivid.
[Edited on 8-1-2007 by YT2095]

I walked to lab and took out my hydrogen peroxide, hydrochloric acid and manganese dioxide and walked straight out to the night. I mixed about a half of teespoon of manganese dioxide with about 10ml of 38% hydrochloric acid. I heated up the reaction and large ammounts of chlorine was released. The smell was quite unpleasant even though I was outside. I added slowly one ml of 30% hydrogen peroxideI and saw a bit of a yellow foam forming. After 10 minutes the solution is clear and has no traces of MnO2 in it, but it is yellow, as yours. Might be due to the impurities like iron?

I am now going to add some more MnO2 to the mixture tomorrow and probably try to convert it to carbonate by adding sodium bicarb.

[Edited on 8-1-2007 by kaviaari]

YT2095 - 8-1-2007 at 12:47

if there`s impurities, it`s going to be much easier to shift them when all is dissolved. at least now you have your Mn compound in soln.
from there you may work

I`m pleased that the H2O2 think worked for you also in HCl, I can`t be 100% certain yet as to WHY it works like it does, but I have a feeling it may be catalytic and catching the free ions as they react with the peroxide and locking them to the acid whilst fee.

I`ll leave that to the experts here to establish for certain, do NOT take my hypothesis as truth!

kaviaari - 8-1-2007 at 12:59

Quote:

Originally posted by YT2095
I`m pleased that the H2O2 think worked for you also in HCl, I can`t be 100% certain yet as to WHY it works like it does, but I have a feeling it may be catalytic and catching the free ions as they react with the peroxide and locking them to the acid whilst fee.

MnO2+4H+2e-->Mn2+ +2H2O

H2O2-->2H+ +O2+2e

And this happens in highly acidic solutions. So H2O2 helps converting MnO2(s) to Mn2+-ions and therefore the manganese reacts way faster.

Waffles - 8-1-2007 at 18:21

Quote:

Originally posted by YT2095
on microscope examination it is indeed crystaline too for the most part, the Thermit reaction is becoming More and More tempting as time passes.
if I have to go this route, does anyone know how "Violent" an Alu MnO2 reaction is?
I don`t want some exotic type sh!t that`s going ruin the product or create annoyance/disturbance.

Too violent. Example:
http://img.photobucket.com/albums/v97/iamthewaffler/science/...
This picture was taken at about 3 in the afternoon, quite sunny. The brightness of the reaction just washed out all the sunlight.
If you want to think about an aluminothermic reduction, you need to use Mn3O4

guy - 8-1-2007 at 18:29

Quote:

Originally posted by iamthewaffler

Quote:

WOW!

chemoleo - 8-1-2007 at 18:39

A way round this is to use large-grain Al, or, as waffler says, Mn3O4 which can be obtained from MnO2 by roasting. I can dig out the procedure if you wish.

PS I did this thermite too, and it was fairly vigorous. But using the same Al powder, and Fe2O3, it was equally vigorous. And we all know Fe-thermite works for preparing Fe. Thus its reaction is adjusted by grain size. CuO/Al still tops it all, it is essentially behaves like flash, and is *extremely* sensitive to flame, sparks etc. Any confinement invariably produces a loud bang and a copper cloud.

kaviaari - 9-1-2007 at 03:49

Some more notes about my experiment. Now I have added a teespoonfull of MnO2 to the mixture. I also added two milliliters of 30% hydrogen peroxide. Solution is still quite yellow. I think this is not due the impurities but due to the Mn2+-ion. Some kind of odd yellow precipitate is also formed. I added a millilter of HCl, but it is not dissolving. Well I can filterate it off. The method is very efficent and I am going to scale it up for converting all of the impure manganese dioxide to manganese chloride.

[Edited on 9-1-2007 by kaviaari]

YT2095 - 9-1-2007 at 03:55

wow! is right, that reaction is way out of my league for comfort and I`ve a feeling that picture doesn`t do it justice either, esp as you mention if WAS daylight and it looks like Night time by comparison, perhaps something scaled down to a tenth of that size as a last resort, the only Alu powder I have is 10 micron, far too fine for a "Tame" reaction.

I`ve finished roasting 5g of the MnO2, and the strange part is that the powder is now Brown, and my book lists all oxides of Mn as Black with the exception of MnO and it say Green for that. brown is not mentioned anywhere?

I was intrigued as to why there were several flashes going on in the crucible also, so I put a magnet in the MnO2 I have, and a percentage of it is magnetic? so that`s either the Mn metal itself not fully converted to the Oxide, OR there is an oxide of Mn that is magnetic like there is with Iron hematite.

I`m beginning to question the purity of this material.

woelen - 9-1-2007 at 07:21

Mn(III) is brown in most of its compounds. You roasted the oxide and part of the Mn is converted to +3 oxidation state.

How easy can this Al/MnO2 mix be ignited? Is it simply a matter of holding the mix above a flame? Sounds interesting. I now have some German Dark Al-powder (1.5 micron particle size), so that should give a neat flash. I already tested that German Dark with KNO3/P/Al in Cl2 gas and the result is awesome (much better even than with the 40 micron Al I mentioned some time ago)! So, this definitely is something I want to try with MnO2. I have an almost unlimited supply of MnO2 from a local supplier at a price of just a few euro's per kilo, so that would be a good source for such flash experiments.

Hilski - 9-1-2007 at 08:58

Quote:

I added slowly one ml of 30% hydrogen peroxide I and saw a bit of a yellow foam forming. After 10 minutes the solution is clear and has no traces of MnO2 in it, but it is yellow, as yours. Might be due to the impurities like iron?

Last time I made MnCl2 from MnO2, I added 3% H2O2 to the mix after it appeared to have stopped reacting. Huge amounts of Cl2 came puffing out of the beaker almost instantly, all the rest of the MnO2 went into solution. I have to hand it to guy for recommending the H2O2 method back when I asked about it last spring. It definitely works. But watch out for the chlorine if you use HCL!

[Edited on 9-1-2007 by Hilski]

woelen - 9-1-2007 at 13:51

I did the experiment with MnO2 and I'm really amazed by the powerful flash reaction. I never could imagine that this chemical would be capable of such an extreme reaction. I always thought that was something for the more well-known pyro oxidizers like KClO4, KNO3, etc.

[img]http://woelen.scheikunde.net/science/chem/exps/MnO2+Al+P/flash2.jpg[/img]

I found out that the mix gives a brilliant white flash, but it requires lightling with a hot flame. Adding just a percent or 2 of red P makes the mix much easier to ignite, while it hardly becomes less bright.

I put the result of the experiment on the web in a temporary web page. It is not yet a completed web page, I just added it to share it over here, but don't regard it as a final web page. Watch the bright flash, you can obtain from such a small amount.

http://woelen.scheikunde.net/science/chem/exps/MnO2+Al+P/ind...

EDIT: Why don't I see the image in the post?

[Edited on 9-1-07 by woelen]

YT2095 - 10-1-2007 at 02:52

Woelen, Nice!

I am curious about all these reports of Cl2 being liberated here, I`ve not even had a HINT of of such gas produced with my stuff, plenty O2, but not Cl?

9% H2O2
30% HCl
MnO2 (technical grade IMO)

woelen - 10-1-2007 at 03:04

With just my MnO2 and HCl I also don't get Cl2, so we are on the same line of observations.

But when I mix in H2O2, then I get copious amounts of Cl2, and I definitely MUST do the experiment outside. Even just a spatula full of MnO2 and a few ml of the liquids produces a lot of Cl2, one 100 ml erlenmeyer, full of green gas.

YT2095 - 10-1-2007 at 03:12

something`s going really wrong here then, I have all 3 mixed and not even a slight hint of chlorine?

this is in a 250ml flask filled to the 100ml line so far with a mushroom colored muddy liquid that is yellow when the muddy ppt settles, there is also metalic particles in the bottom that fiz quite a bit and are magnetic.
I add the H2O2 1 ml at a time as I don`t want it to foam over the flask.
I took a sample of the yellow liquid and added it to sodium carbonate and got my light pink MnCO3 gel like ppt, so I know Something`s working, just not the Cl2 bit?

JohnWW - 10-1-2007 at 04:39

I remember reading somewhere that Mn(SO4)2 can be obtained from MnO2 by reaction with either vapor-phase SO3 or with fuming sulfuric acid (H2S2O7), but it is very sensitive to hydrolysis (there must be no net water of reaction produced as byproduct).

woelen - 10-1-2007 at 04:42

What concerns me are the metallic magnetic particles and the yellow color. I think you have a lot of impurity in one of your reagents, otherwise I cannot explain what you get. My solutions of MnO2 in H2O2/acid are colorless, Mn(2+) ion is VERY pale pink, almost colorless.

This impurity could inhibit formation of chlorine, the oxidation may go through another pathway, such that hardly any Cl2 is formed.

YT2095 - 10-1-2007 at 04:58

the only listed KNOWN impurities are in the H2O2, it containd H3PO4 as a stabiliser and Phenacetin (no idea what that`s for?).
now those two have never messed up a reaction for me yet.

I think it`s the MnO2 at fault, there`s metalic parts in it, black magnetic parts in it also, IMO I don`t even think it would make Technical grade! :mad:

the brown muddy ppt say a different oxidation state to me also, the MnO2 that roasted in the crucible would also make this color if I added water.

which reminds me, I just finished roasting some with powdered activated charcoal, that stuff is now REALLY Brown almost orange around the edges, and green/gray underneath the surface at the bottom of the pot, I think this stuff will probably be a little safer to use in a thermit reaction.

EDIT: as an afterthought, the H2O2 + HCl rxn is done at room temp here, no heating used at all, maybe that plays a part?

[Edited on 10-1-2007 by YT2095]

[Edited on 10-1-2007 by YT2095]

woelen - 10-1-2007 at 06:00

I do not use any heating in these reactions.
I think that the MnO2 is full of iron impurity. Magnetite, Fe3O4 is black and on heating that may form Fe2O3 (orange/brown), with uptake of oxygen from the MnO2, surrounding it.

I get the impression that your MnO2 is not the chemically pure compound, but some finely powdered ore. I know of an ore (called "mangaanknollen" in dutch, I don't know the english word for it), which is mostly MnO2, but also contains a considerable amount of ironoxides and manganese in its +3 oxidation state.

YT2095 - 10-1-2007 at 06:25

Im beginning to think you`re exactly right, here we call it Pyrolusite.
an ore would make Perfect sense!

in the PM I sent you, I said that of this 250g, probably 100g is actualy good material.
an interesting note though, when roasted with the carbon, the resulting material will react directly with 50% sulphuric acid, giving off a little heat and leaving a pinkish soln with brown mud at the bottom.
it also Hisses when you add it to dil HNO3 but there`s little color change except to a yellow tinge and it becomes quite turbid and a little thicker (viscous).
I can personaly vouch for the Carbon purity (I did it myself with ABA treatment TWICE)

manimal - 14-6-2008 at 13:33

Does anyone know if MnO2 reacts with dilute acetic acid? I have about 350 grams of MnO2 that I extracted from alkaline batteries, and I washed the crude material in vinegar to remove residual Zn and ZnO, and it immediately evolved gas. Is it just H2 from the residual zinc, or is it MnO2 reacting to form MnAc2?

The_Davster - 14-6-2008 at 13:45

I am not sure if MnO2 will react quickly with acetic acid, but if it did, it would not produce any gases. You are likely right in that it is just fizzing with the residual zinc.

12AX7 - 14-6-2008 at 20:18

Alkalines? Could be CO2 from atmospheric exposure.

Tim

MnO2 thermite problems

blogfast25 - 15-6-2008 at 09:55

Another problem with Mn thermites, especially MnO2 based, and often overlooked is the following. To produce decent metal, the slag and newly formed metal have to separate out from each other. For this to stand a chance, both slag and metal have to be liquid at the end of the reduction reaction. In the case of a thermite reaction, the slag is mainly alumina, with a melting point of 2,054 C. Unfortunately the boiling point of Mn is only scarcely higher: 2,061 C and as a result a successful MnO2 reduction can only lead to part of the metal being boiled off... The end-temperature of such a thermite reaction (in adiabatic conditions) can be estimated quite accurately. My own calculations show clearly that the end-temperature of a fast MnO2 aluminothermic reduction is well above 2,061 C.

An example of such a calculation can be found on this page, here applied to a Titanium thermite reaction.

I've used various grades of MnO2, with fairly high amounts of flux (CaF2 - a slag fluidiser) and have rarely obtained good quality metal and always with quite low yields (< 20 % of theoretical yield).

I believe the key points to obtaining good Mn metal are:

• use the lower oxides or a blend of MnO and MnO2
• use fairly coarse Al to reduce reaction speed and suppress thermite end-temperature: too high and your metal will simply distil off
• use adequate amounts of CaF2 to improve slag fluidity

In one particular reaction, using homemade (from battery crud) but quite pure MnO2, I found much of the metal back in the form hundreds of sub-mm Mn globules around the crucible: frozen Mn rain (I kid you not)!

Industrially, Mn metal is made by thermite reduction but in a semi-continuous way, rather than by batch method. This allows to keep the temperature as low as possible, while still melting the alumina, necessary to allow the Al2O3 and Mn metal to separate out.

I've not found any Internet accounts of successful Mn thermites that mentioned good metal obtained. Batch-wise the heat balance may be too precarious...

[Edited on 15-6-2008 by blogfast25]

12AX7 - 15-6-2008 at 12:05

Probably fine in a bomb. Got to be careful it doesn't melt though, hahah. Probably a large steel pipe nipple lined with refractory (sand or etc.) and the thermite charge would suffice. A good way to make sodium, too!

Tim

nodrog19 - 15-6-2008 at 16:21

my friend was messing around in hos lab, had a MnO2 related accident.
MnO2+2HCl-->Mn(OH)2+Cl2
Mn(OH)2+2HCl-->MnCl2+2HOH
he got a mouth full clorine, but then he (i dont know why)
decides to add H2SO4
MnCl2+H2SO4-->MnSO4+2HCl
squirted in to his eyes and burned his bench.
good thing he had a wash bottle laying around

blogfast25 - 16-6-2008 at 11:44

Regarding possible explosions with MnO2 thermites:

in my notebooks I found the following adiabatic heat balance for MnO2:

MnO2 + 4/3 Al ---> Mn + 2/3 Al2O3

Heat of reaction ΔHr = - 597 kJ (per mol of MnO2)

The enthalpies to heat 1 mol of Mn and 2/3 mol of Al2O3 to 2,740 K are respectively 110.5 kJ and 300.1 kJ or a total of 411 kJ. Since as ΔHr is -597 kJ, there is about 186 kJ to spare. Simply put, in adiabatic conditions a pure MnO2 thermite is energetic enough to heat both the formed metal and the alumina slag to well above 2,740 K (2,470 C) and thus well above the boiling point of Manganese metal. In an enclosed space, pressure build up would be inevitable and an explosion becomes a possibility.

For MnO + 2/3 Al ---> Mn + 1/3 Al2O3 the situation is very different because ΔHr = -176 kJ (per mol of MnO) and this reaction would not reach 2,740 K because:

176 < 110.5 + 150 or 176 < 260 (all kJ)

A blend of MnO and MnO2 at carefully chosen molar ratio of the oxides, might allow the right temperature to be reached without boiling off the metal on the one hand or obtaining very poor slag/metal separation (too low temperature) on the other hand.