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Author: Subject: Manganese Carbonate
mericad193724
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[*] posted on 8-7-2006 at 05:40
Manganese Carbonate


Hello,

Is it possible to synthesize Manganese Carbonate from fairly common chemicals.

The reason is I want to make Manganese Oxide, one way is to use HCl and Manganese Carbonate...the other is MnO2 with HCl, which should be heated. (Wikipedia) How much heat? I would prefer method one so I would like so make manganese carbonate.

Thanks

Mericad
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enhzflep
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[*] posted on 8-7-2006 at 06:02


Perhaps the simplest route to Manganese Carbonate would be via a ceramics supplier. It apparently finds uses in glazes. MnCO3

Manganese(IV) oxide - MnO2 is of course found in both alkaline and carbon-zinc cells.

If in fact MnO2 is what you're after and do not need a vast quantity, then perhaps new (dirt-cheap crappy brand) cells could be your best bet. Albeit a somewhat dirty job.

The synthesis of Pottasium Permanganate detailed on Wiki
here --> http://en.wikipedia.org/wiki/Manganese_dioxide

What did you plan to do with it? Catalyst?

[EDIT: removed - can't get sub-scripting to happen]

[Edited on 8-7-2006 by enhzflep]
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mericad193724
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[*] posted on 8-7-2006 at 06:44


WOW...sorry ...big mistake

I would like to make Manganese Chloride, not Manganese Dioxide. I have 500g of MnO2 and though it would probably be a reactant.

I would like to make MnCl2, NOT MnO2. Sorry about that.

Is there a way to easily synthesize this, Wikipedia says MnO2 + HCl with heat. How much heat? Couldn't find any other info

thanks
Mericad
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enhzflep
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[*] posted on 8-7-2006 at 06:57


Hmm. I guess that since Cl2 is a by-product and found on the right hand side of the equation, it will be emmitted. In this case, I would expect that a gentle heating until it is evolved should suffice.

Having a B.P of some 1225 deg C, suggests (as does the fact that it's ionic) that you're not going to decompose it any-time soon. I would be inclined to heat as strongly as your nose can tolerate.

Not forgetting that this could be a reaction that bubbles somewhat. I'd say that your 3 main points to watch would be (a) a boiling over (b) excessive chlorine produced (i.e more than can be dissipated safely) and (c) your HCl is all evaporated and potentially breathed.

Just try a super small sample and keep going until it turns pink, I suppose.

regards.
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Mr. Wizard
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[*] posted on 8-7-2006 at 07:05


Mixing MnO2 with HCl will generate Manganese Chloride, but will also generate plenty of Chlorine gas. It was used by Scheele to isolate Chlorine.
MnO2 + 4HCl ? MnCl2 + Cl2 + 2H2O
Be ready for it. ;)
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[*] posted on 8-7-2006 at 07:39


WHen I was making a chlorine generator for an experiment I used the MnO2/HCl method which didn't work for me. I was using reagent grade materials. There was a chlorine like smell though
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[*] posted on 8-7-2006 at 08:04


Odd, when I add HCl to pottery-grade MnO2, I get plenty of choking fumes.

Tim




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[*] posted on 8-7-2006 at 09:26


couldn`t you do a displacement reaction with another metal Chloride instead? you would at least avoid much of the noxious gas that way.

just a thought :)




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[*] posted on 8-7-2006 at 11:28


Agricultural manganese sulfate, sodium bicarbonate, boiling water.

Dissolve and filter the manganese sulfate. Disolve sodium bicarbonate in boiling water to drive off the excess CO2, slowly add filtered manganese sulfate solution, have plenty of extra volume in container to allow for foaming.
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chloric1
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[*] posted on 8-7-2006 at 11:44


If your MnO2/Hcl smells of chlorine but no large amount of gas evolution ensues then you might need a gentle heat. I noticed that KMnO4 and conc HCl create a brown/green solution which is most likely MnCl3 and is thermally unstable.



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[*] posted on 8-7-2006 at 12:31


Hold on. Way safer way to make it without chlorine. Mix and acid and H2O2 and MnO2. H2O2 will be oxidized and MnO2 will go to Mn2+.

I have done this before and it works, but looking at the redox potentials it looks nonspontaneous. I think it works because H2O2 ---> O2 + 2H+ supplies more H+ as it decomposes making it spontaneous, along with the effect of overvoltage.

[Edited on 7/8/2006 by guy]




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[*] posted on 8-7-2006 at 13:08


The redox potentials indicate it is spontaneous.
MnO2 + 4H+ +2e- --> Mn2+ +2H2O E=1.22
O2 +2H+ +2e- -->H2O2 E=0.70

Net
MnO2 +2H+ +H2O2 --> Mn2+ +2H2O +O2




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[*] posted on 8-7-2006 at 13:11


Yeah but there is another reaction for it.

2e + H2O2 + 2H+ ----> 2H2O E=1.77




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[*] posted on 8-7-2006 at 14:55


But that is when peroxide acts as the oxidizing agent, in this reaction it is the reducing agent.



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[*] posted on 8-7-2006 at 15:38


How does it choose? It is supposed to be strong enough to oxidize MnO2 back to MnO4-.



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[*] posted on 9-7-2006 at 22:36


MnO2 definitely is not transformed to MnO4(-) with H2O2. If the solution is sufficient acidic, then the reaction proceeds smoothly to form Mn(2+) and O2. But it is important to have a very acidic solution (e.g. 1 M H2SO4), otherwise the MnO2 simply catalytically decomposes the H2O2.

I also noticed, that some forms of MnO2 are very inert. I have two grades. The one is an almost black very fine powder, the other also is a fine powder, but it consists of many small glittering crystals. This latter MnO2 hardly dissolves in any solvent, only with difficulty it dissolves in conc. HCl. I think that it is calcined and made more inert that way. This inertness is common with many other calcined metal oxides.




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mericad193724
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[*] posted on 10-7-2006 at 02:29


thanks for the help guys.


I will try MnO2 I got off eBay with HCl as 31.45% acid used to treat pools. If it doesn't work I will slightly heat it.
I will wear a respirator so no worries about the chlorine gas.
--------------------------------------------------------------------------------------------------------------------
I don't want to start a new topic for this...

Is it possible to convert Ferric Chloride to Ferrous Chloride. I have 1 gallon of Ferric Chloride used as an etchant in electronics.

thanks
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guy
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[*] posted on 10-7-2006 at 12:31


Quote:
Originally posted by woelen
MnO2 definitely is not transformed to MnO4(-) with H2O2.


Yes I know this, but why?




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[*] posted on 10-7-2006 at 12:55


Mn catalytically decomposes H2O2 for one. I don't know if that is meaningful electrochemically...

Tim




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[*] posted on 10-7-2006 at 17:56


6e- + 3H2O2 + 6H+ -------> 6H2O
2MnO2 + 4H2O ----------> 2MnO4- + 8H+ + 6e-
_______________________________________
2MnO2 + 3H2O2 ----------> 2H2O + 2MnO4- + 2H+

Therefore in acidic conditions this reaction will not happen.





[Edited on 7/11/2006 by guy]




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mericad193724
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[*] posted on 10-7-2006 at 18:09


I just tried MnCl2 synthesis. 5g MnO2 was put in a test tube and 31.45% HCl was added. Slight bubbling was observed, nothing excessive. When heated slightly, it bubbled much more. I still have a black liquid, not a light pink liquid.

It didn't really work with my MnO2:(

Mericad
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[*] posted on 10-7-2006 at 18:15


Quote:
Originally posted by mericad193724
I just tried MnCl2 synthesis. 5g MnO2 was put in a test tube and 31.45% HCl was added. Slight bubbling was observed, nothing excessive. When heated slightly, it bubbled much more. I still have a black liquid, not a light pink liquid.

It didn't really work with my MnO2:(

Mericad


Some chlorine was produced, try adding more acid. Or try the H2O2 method, it works well.




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[*] posted on 15-7-2006 at 06:38


I got a small sample of MnCl2 from my teacher. When mixed with unstabilized peroxide 3%...is bubbled up and turned black. Why did this turn black???

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[*] posted on 15-7-2006 at 06:54


It was oxidized to the insoluble MnO2 (and whatever Mn2O3 and other half-assed oxides may be produced).

Tim




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[*] posted on 15-7-2006 at 08:18


Quote:

I just tried MnCl2 synthesis. 5g MnO2 was put in a test tube and 31.45% HCl was added. Slight bubbling was observed, nothing excessive. When heated slightly, it bubbled much more. I still have a black liquid, not a light pink liquid.


There is something called stoichiometry that you have to take in account when making any reaction whatsoever. It seams pretty obvious that much MnO2 remains in the reaction mixture unless there is enough HCl(aq) used. You need at least 4 mol of HCl for every mol of MnO2. But in practice you need an excess of HCl to drive the redox reaction to the end in less than infinite time, as well as some heating. Likely there will still be some undissolved solids, not necessarily MnO2, but some impurities perhaps, so good laboratory practice calls for filtration before stripping of the solvent and HCl and Cl2 remains.

PS: How do you put 5 g of MnO2 in a normal test tube and the minimum 25 ml of HCl(aq) if 36% concentration is used? You can't, so use proper glassware for your reactions.




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