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Author: Subject: MnO2 -> MnCl
thioph
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sad.gif posted on 31-1-2008 at 02:32
MnO2 -> MnCl


I recently tried to make some MnCl by heating 1 mole of MnO2 with 4 moles of HCL with magnetic stirring.

On addition of the HCL to the MnO2 it foamed up in a black mess and started giving off alot of chlorine gas. I continued stirring and heating for an hour or so.

Im left with what looks like MnO2(just as much as I had put in) sitting on the bottom of water.

The MnO2 was from alkaline batteries, and the HCL was hardware grade. Im unsure what I have missed or have done wrong. Is it possible that the black powder I'm starting with isn't MnO2 at all? Or have I made some other mistake?

Any suggestions or tips would be great.
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[*] posted on 31-1-2008 at 02:36


try adding a little H2O2 into the mix, also you must remember that the MnO2 in batteries is mixed with Carbon so you will still have black stuff in there.



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[*] posted on 31-1-2008 at 02:42


There are other threads on this, but it will take some searching to find them, so here's a thumbnail.

Battery MnO2 also contains carbon for conductivity, and will remain after the MnO2 has reacted. Filtration will remove the carbon, but ...

There are other cations in the battery MnO2 mixture besides Mn - thin on how a battery works. You need to wash the dioxide by stirring it with water, then letting it settle and decanting; repeat several times. With a bit of practice, testing a bit of the solids with hot HCl, you can learn to pour off some of the carbon still in suspension from the MnO2; this besides removing the cations in solution.
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[*] posted on 31-1-2008 at 03:19


I want to add a final remark. You don't get MnCl, but MnCl2. The reaction is as follows:

MnO2 + 4HCl --> MnCl2 + 2H2O + Cl2 (highly simplified)

In reality, it also is surprising to see how that even with pure MnO2 from a chemical supplier, this reaction does not proceed as smoothly as theory suggests.

I of course, in my starting years, also tried this experiment, and I used lab grade MnO2 and a decent grade hydrochloric acid (so-called 'rein', meaning pure). With this, a dark green solution is produced, which contains a complex of manganese(IV) and chloride, or manganese(III) and chloride, not sure about that.

The final solution then contains MnCl2, remains of excess HCl and this dark green compound, which only VERY slowly disappears. In order to get a colorless solution of MnCl2 in excess HCl, you have to filter the solution, removing any solid remains of the MnO2, and then just a few drops of H2O2 removes the dark green complex (it reacts with the H2O2 to Mn(2+) ions and oxygen).

This is an interesting experiment and there is much more to be said on this than high school chemistry books are suggesting.




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[*] posted on 31-1-2008 at 07:54


MnO2 is always a tough nut to crack.

But just a thought. If MnO2 removed from Battary is mixed with carbon then that 'C' could be profitably utilized for reducing MnO2 to MnO which is much more docile thing to handle.

Just try heating the mix strongly in absence of O2 (but NOT in a completely closed or sealed vessel. Allow water vapours and CO2 formed to escape). The mix should get ignited ca. 400-450 Deg C.

After the reaction is over quench the mass in water and you have freshly prepaired MnO which will readily react with even dilute HCl (and with out giving our noxious Cl2)

2MnO2 + C ----------------> 2MnO + CO2

MnO + 2 HCl ----------------> MnCl2 + H2O

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[*] posted on 31-1-2008 at 08:08


It seems to me that I've reduced MnO2 to a soluble product using ascorbic acid (vitamin C). If I remember, I'll give it a try again -- I've just picked up some fine MnO2 at the pottery supplier, and I've got vitamin C tablets to spare.

Once it's reduced and dissolved, it should be easy enough to precipitate as the carbonate, then redissolve with HCl.
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[*] posted on 31-1-2008 at 08:20


Quote:
Originally posted by gsd
...
After the reaction is over quench the mass in water and you have freshly prepaired MnO which will readily react with even dilute HCl (and with out giving our noxious Cl2)

...
gsd


Only if there is sufficient carbon, otherwise you'll have Mn2O3 and Mn3O4.

Also note that battery MnO2 is fairly reactive, while the ignited oxides are more crystalline and inert.

Gently boiling the solution from reacting MnO2 with HCl will drive of chlorine and convert any higher oxidation state Mn to Mn(II)
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[*] posted on 31-1-2008 at 08:50


Quote:
Originally posted by -jeffB
It seems to me that I've reduced MnO2 to a soluble product using ascorbic acid (vitamin C). If I remember, I'll give it a try again -- I've just picked up some fine MnO2 at the pottery supplier, and I've got vitamin C tablets to spare.


I think this reduction should work with Oxalic acid, glucose, plain sugar, Jaggary etc.
why waste vitamin C if these mundane compounds can do its job ;)

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[*] posted on 31-1-2008 at 09:05


Quote:
Originally posted by not_important

Only if there is sufficient carbon, otherwise you'll have Mn2O3 and Mn3O4.

Also note that battery MnO2 is fairly reactive, while the ignited oxides are more crystalline and inert.



Just consider this. MnO2 : MW = 87, therefore 2MnO2 = 174; C : AW=12

Which means C required for the reaction is only about 6.5 % by wt in the mix of C & MnO2 to get a good conversion to MnO

Also it is very important to quench the reaction mass in water while it is hot. The slow cooling in air results in re-oxidation of MnO to higher oxides and also promotes the crystallization of the matrix rendering it inert.

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[*] posted on 31-1-2008 at 10:53


I've performed this reduction before, using ground hardwood charcoal; 2MnO2 + C = Mn2O3 + CO begins about when carbon starts burning (500Cish) and, in a heated crucible, proceeds to a peak temperature of about 800C. Gas production throws the reaction about; a screen or filter is helpful (a pad of ceramic fiber held firmly on the crucible suffices). When the reaction is done, a brown product is observed. On further heating with excess carbon, the white MnO can be reached, well away from air, over 1000C.

The brown product gives Cl2 when added to HCl, but it dissolves easier than MnO2, giving less noxious chlorine.

Tim




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[*] posted on 31-1-2008 at 12:11


A few comments.

Spent batteries do not contain a great deal of MnO2. Efficient alkaline cells will convert it to such things as Mn2O3, MnO(OH). The amount of carbon used is variously quoted as 10-15% by weight to 50% by volume, the latter being the amount for maximum conductivity.

This carbon cannot be easily removed. Igniting to bright red heat will burn it off if plenty of air is allowed. However, the MnO2 will also be decomposed by the heat and also reduced by the carbon. You will get Mn2O3. At around 1000C (AFAIK) this oxide also decomposes to Mn3O4.

You will never get MnO if there is any oxygen around. It oxidizes to Mn2O3 in air rapidly. MnO is pyrophoric, IIRC.

All the oxides bar MnO produce chlorine with HCl, in vaying amounts, MnO2 the greatest.. All produce (ultimately) MnCl2. MnO2 "dissolves" in cold Conc. HCl to produce a brown liquid which may transiently contain MnCl4 but disengages Cl2 to give MnCl3, which is highly unstable, and when the brown liquid is heated, gives MnCl2 and more Cl2.

All divalent Mn salts are somewhat unstable and tend to oxidize to an oxide or hydroxide and become brown. If a small amount of the parent acid is kept with them, this can be avoided.

Mn3O4 is reddish; Mn2O3 is brown/black . MnO2 varies - it is black in the non hydrated form and brownish when hydrated. And MnO is green or grayish.

MnO can only be made by heating and reducing one of the other oxides, with hydrogen for example.

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[*] posted on 3-2-2008 at 09:07


I've extracted MnCl2 from spent (non-alkaline) batteries several times using a procedure posted on this forum by DerAlte. It works well.

My purpose was always to obtain the actual MnO2, so I always re-oxidise the MnCl2 back to MnO2 (for use in thermite reactions).

I've found the content of MnO2 in my spent battery cruds (from a "no-frills grade" of battery) to be just over 30 w%, the rest is mainly graphite.

Now I make up some battery crud, then calculate the amount of HCl (32 w%, "drug store grade") needed, add a little excess HCl and leave the crud/HCl slurry to stand overnight, at RT. It generates chlorine of course and foams quite a bit.The next day I'll give the slurry a quick boil to complete the reaction and flush out most chlorine.

After cooling the slurry filtrates to a light yellow filtrate of dilute MnCl2. Never had any problems with unexpected green complexes of Mn.

For re-oxidising the Mn [+II] to [+IV] with thin bleach (hypochlorite), I recommend neutralising the MnCl2 to MnCO3 (with washing soda), rather than to Mn(OH)2 (with caustic soda). Using the hydroxide route I've found that that inevitably produces some unwanted permanganate (MnO4(1-), deep purple). Presumably the MnCO3 buffers things to approx. the right pH to avoid permanganate formation.

It's widely believed that MnO2 is "black as the ace of spades" but that isn't actually true: pure MnO2 is actually a dark "pure chocolate" brown. Native MnO2 (Pyrolusite) is in Dutch called "Bruinsteen" which literally translates as "brown stone"...
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[*] posted on 4-2-2008 at 00:08


I am Dutch and know the word "bruinsteen", but I still am inclined to think that this name is misleading if it is used for the pure chemical MnO2. I have labgrade pure MnO2, a very fine powder, and it is black like soot. I also have pottery grade MnO2, and that also is black (very dark grey, glittering crystals). The pottery grade stuff is hard to get into solution, it requires a lot of heating.

The brown material is something of composition MnO2-x, with x > 0 (but still much less than 1). So, it can be regarded as MnO2, deficient in oxygen and hence containing manganese(III) and manganese(IV).

The green complex I mentioned, can be destroyed by boiling, but I found it to be remarkably persistent. However, a pinch of sodium sulfite, added to the solution, immediately destroys the complex. This complex can be obtained more easily by adding some KMnO4 to concentrated hydrochloric acid. Just try it, add a tiny amount to a few ml of 30% HCl and see what happens. Only use a tiny amount, otherwise the solution becomes too dark and nothing interesting can be observed in that case.




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[*] posted on 4-2-2008 at 08:18
Black or brown?


Well, well. The only MnO2 I've ever laid eyes on is stuff I've made myself by re-oxidising MnCl2 extracted from battery crud. In some cases I've seen it precipitate (immediately after oxidation from +II to +IV) as a black slurry, only to see its colour lighten upon filtration, washing, drying and calcining. Mostly it precipitates as a brownish slurry. But in the end it always comes out as very dark brown, but easily distinguished from pure black...

I don't have any lab grade or pottery grade MnO2, so cannot compare. But after I saw the first batch of home made MnO2 to be a very dark brown, I was under the impression that most people believed it to be black because when photographed it would be difficult to distinguish the colour from black. Also, there remains this consistent misconception that battery crud = MnO2, but battery crud owes its blackness partly to the majority of black graphite it contains. Wiki describes manganese dioxide as This blackish or brown solid [...].

I wonder what is causing this slight difference in colour. Could it be calcining temperature, for example? I can only calcine up to about (est.) 600 DC, so technically it's not even really calcining, although enough to drive off any remaining moisture (which is my purpose).

Woelen, another riddle? I very much like your pages on the Cu [+I]/Cu [+II] complex riddle...
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[*] posted on 13-3-2008 at 12:27
MnO2 to MnCl2


Today I tried making MnCl2 from MnO2 found in alkaline batteries.
MnO2 + 4HCl -> MnCl2 + Cl2 + 2H2O
I took the black powder from the battery and dumped it in a beaker. It must have been around 3g. I then added 2M HCl until there was no more chlorine gas developing. The mixture was filtrated to remove remaining black powder (quiet a lot, should be carbon?) and the water boiled off until dry. I then poured in some more water, heated until all dissolved, put in fridge and filtered off solids by vacuum.
According to wikipedia the hydrates of MnCl2 are pink in color. However, my product was bright white. It weighted around 500mg... All I could analyze with was a flame test, which showed a yellow, maybe greenish color (which manganese(II) should), but I have not practiced this test many times and I could be wrong.

Could anyone point out what went wrong? The reaction seemed so simple, but I obviously made a mistake?
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[*] posted on 13-3-2008 at 12:47


Quote:
Originally posted by bluetrain
Could anyone point out what went wrong? The reaction seemed so simple, but I obviously made a mistake?

What makes you believe you made a mistake? Just because what you got was not pink? You might reconsider. Check Woelen's posts above and this thread as well: http://sciencemadness.org/talk/viewthread.php?tid=6181 .

Also, it is highly advisable to use the search engine before posting just in case there already is an ongoing discussion about the same topic. Like for example this thread with which I just merged your post.

[Edited on 13/3/2008 by Nicodem]
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[*] posted on 13-3-2008 at 23:55


Bluetrain, manganese(II) salts can be white (or VERY pale pink). Especially the anhydrous salts can have very pale colors. In solution, the Mn(2+) ion has a very pale color also. I consider it colorless in all practical situations.

IIRC batteries also may contain NH4Cl. It could be that your material contains NH4Cl as well. You can easily check out. Take a minuscule spot of the white solid, add a few pearls of solid NaOH and crush the mix. Even if your material contains a small amount of ammonium chloride, you will notice the smell of ammonia. If there is no smell of ammonia, then the amount of ammonium chloride will be neglectable.




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[*] posted on 15-3-2008 at 15:01


Washed the black stuff from the batteries with water first to remove eventual ammonium chloride and other water solubles. Also, used a higher concentration of hydrochloric acid. Got around 2g of pale pink crystals from 5g "black stuff". Formed an orange solution with water, and some yellow colored impurity had to be removed from raw product by recrystallization.

Chlorine gas is shit! Could almost not smoke the entire night, had such a cough. Will have to do it in a more ventilated area, preferably outside...

[Edited on 16-3-2008 by bluetrain]
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[*] posted on 15-3-2008 at 17:54


I have a couple solutions of manganese crap that are yellow to orange. I'm guessing some equilibrium of O2(g) vs. Cl2(aq), with Mn(II)--Mn(III) somewhere in the middle. I need to put H2O2 in these solutions to reduce the Mn. And that's what you need, H2O2.

Quote:
Originally posted by bluetrain
Chlorine gas is shit! Could almost not smoke the entire night, had such a cough. Will have to do it in a more ventilated area, preferably outside...


Why would you need to smoke? The chlorine damaged your lungs as much as a couple packs of smokes would. Who needs cigarettes, really, when you can kill yourself with chlorine a whole lot faster.

Tim




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[*] posted on 29-3-2008 at 18:32


Recently I mixed (1 mole Mn02 & 4 moles HCL) with magnetic stirring let sit overnight alot of chlorine gas was generated. The next day I boiled it for a while and more gas was driven off. I was left with a brownish-black-green solution to which I added a couple of drops of h2o2 the solution turned a yellow-orange color and after 20mins there was a white salt at the bottom.

The same day I once again mixed (1 mole Mn02 & 4 moles HCL) but this time I immediatly added a couple of drops of h2o2 to the mix. This drove off chlorine gas quite rapidly for around 30mins and what was left looked very similar to the first experiment (yellow-orange solution with white salt looking stuff at the bottom).


My question is are either of these MnCl2 or have i made something else? It doesn't seem to be as soluble as I would of thought.

Thank You.
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[*] posted on 30-3-2008 at 15:47


Preparation: http://www.ucc.ie/academic/chem/dolchem/html/elem/elem017.ht...

The most common laboratory method for preparation of Chlorine is to heat 100 gm. of Manganese Dioxide with 300 ml. of concentrated Hydrochloric Acid.
MnO2 + 4 HCl ==> MnCl2 + 2 H2O + Cl2
The gas is bubbled through water to remove any traces of hydrochloric gas that may be present and then it is dried by bubbling it through concentrated sulphuric acid.
Chlorine may also be prepared by dropping cold concentrated Hydrochloric Acid on crystals of Potassium Permanganate.
2 KMnO4 + 16 HCl ==> 2 MnCl2 + 2 KCl + 8 H2O + 5 Cl2
The gas is bubbled through water to remove any traces of Hydrochloric Acid gas that may be present and then it is dried by bubbling it through concentrated Sulphuric Acid.




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