Sciencemadness Discussion Board

Manganese Carbonate

mericad193724 - 8-7-2006 at 05:40

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

enhzflep - 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]

mericad193724 - 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

enhzflep - 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.

Mr. Wizard - 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. ;)

DeAdFX - 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

12AX7 - 8-7-2006 at 08:04

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

Tim

YT2095 - 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 :)

Eclectic - 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.

chloric1 - 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.

guy - 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]

The_Davster - 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

guy - 8-7-2006 at 13:11

Yeah but there is another reaction for it.

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

The_Davster - 8-7-2006 at 14:55

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

guy - 8-7-2006 at 15:38

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

woelen - 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.

mericad193724 - 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
Mericad

guy - 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?

12AX7 - 10-7-2006 at 12:55

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

Tim

guy - 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]

mericad193724 - 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

guy - 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.

mericad193724 - 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???

Mericad

12AX7 - 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

Nicodem - 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.

woelen - 15-7-2006 at 11:58

There is more to this than stoichiometry. Even if you add 10 times as much hydrochloric acid and you use pure MnO2, then the liquid still remains dark green for a long time. This dark green stuff is a chloro complex of manganese in a higher than 2 oxidation state (+3, +4 or a mixed +3/+4 oxidation state complex). Adding a few drops of H2O2, however, immediately destroys this complex and makes the solution colorless.

Mn(2+) ion is said to be pink, but even at high concentration it still looks colorless. The color of this ion is so weak that in practice I would call it colorless.

chemoleo - 15-7-2006 at 18:19

Woelen, we had this conversation before....
MnCl2*x(H2O) is not practically colourless I'm afraid. My proof is analytical grade pink MnCl2 crystals... which dissolve in H2O at supremely high concentrations, and of which remarkably beautiful pink crystals can be grown. I used MnCl2 * 11(IIRC) H2O for crystal trials, it works nicely but needs crazy levels of saturation. The crystals are very beautiful, and melt easily. Sadly I redissolved the crystals to get better ones, but currently I am not growing any because the ambient temp varies way too widely here atm to make decent crystals. Maybe in winter.
I need a dedicated temperature controlled room for these things!
Anyway.... MnCl2 is decidely pink, and not just at insane concentrations! If you really really don't believe me I'll upload a pic of my leftover crystals... (to show I can back this up :))

12AX7 - 15-7-2006 at 19:16

I grew some crystals of MnSO4.H2O myself, they are slightly pink, very nice. Seems to be hexagonal. Hard, but cleaves easily, like FeSO4.7H2O crystals do.

Unfortunately, in recrystallizing them, I dissolved them and the first crop removed was bluish. Now the solution is turning orange, and producing a pale off-yellow scum in the bottom of the jar.

I ought to make MnCl2 some time, as well.

Tim

chemoleo - 15-7-2006 at 19:21

MnSO4 crystals suck, it doesn't really make nice crystals. Commerical MnSO4 is shipped as the monohydrate, and is totally a-crystalline (pink powder). Not surprisingly, it dries on the heater as a powder.
The double MnSO4 salt of NH3 or K is ok, but altogether I wasn't impressed at all, compared to MnCl2, or the other transition metal (Ni, Fe, Cu) NH3/K double salts. I'll post a thread on all these trials one day.

Anyway lets get this back to MnCO3.

[Edited on 16-7-2006 by chemoleo]

guy - 15-7-2006 at 21:58

I used MnCl2 at my school once and it was really pink. This is because the chloride is part of the complex, Mn(H2O)4Cl2 unlike the sulfate salt.

Rosco Bodine - 16-7-2006 at 08:36

Eclectic has stated the easiest method in reacting a
hot solution of manganese sulfate from garden supply
with a hot solution of sodium bicarbonate or sodium carbonate to precipitate the manganese carbonate from
a supernatant solution of sodium sulfate .

If managnese chloride was desired , the manganese carbonate could be treated with HCl .

Alternately , hot manganese sulfate solution could be
treated with hot CaCl2 solution to precipitate CaSO4
from the supernatant manganese chloride solution which would form , the plaster filtered out and the residual solution evaporated to obtain the manganese chloride ,
but this would not be as pure as the product from the carbonate intermediate .

Manganese could have usefulness in energetic compositions .

Filtered and rinsed manganese carbonate possibly could be boiled with ammonium nitrate solution until evolution of CO2 and ammonia ceased , and the residual solution would be manganese nitrate , which could have usefulness as a metal ion detonation catalyst in ammonium nitrate compositions . Strontium , Iron , chromium , and copper nitrates could also be useful , and lead nitrate has similar usefulness but these would likely require a different method for synthesis .

DeAdFX - 16-7-2006 at 09:04

Manganese nitrate... In my high school chem storage I saw a solution of Mn(NO3)2 [If I remember the chemical formula correctly] which also contained a small amount of nitric acid. Is the nitric acid used as a stabalizer to prevent hydrolysis?

Rosco Bodine - 16-7-2006 at 09:19

Generally those nitrates which are subject to hydrolysis to formation of basic nitrates are simply crystallized from
slightly acidic solutions and then kept dry . But a few drops of acid in the storage bottle can be used as an
added insurance against hydrolysis in long storage.

woelen - 16-7-2006 at 10:26

Quote:
Originally posted by chemoleo
Woelen, we had this conversation before....

Yep, now you write this, I remember.

Quote:
[...] Anyway.... MnCl2 is decidely pink, and not just at insane concentrations! If you really really don't believe me I'll upload a pic of my leftover crystals... (to show I can back this up :))

I certainly believe you, so for that reason you do not need to provide pictures. Of course, pictures always are welcome for other reasons. They make chemistry really alive!
I'll try to see that more intense color myself. I have the monohydrated MnSO4.H2O (I have a picture of that VERY light pink powder on my website). I'll dissolve some of this in conc. HCl and see the more intense pink color of the chloro-complex. I'll make pictures of the solution and post them here (also my website will be updated if I succeed in making this pink solution ;)).

woelen - 16-7-2006 at 13:47

I did the experiment of dissolving some MnSO4.H2O in as little as possible reagent grade HCl (30%). The MsSO4.H2O also is a good grade chemical, and not something from a pottery supplier.

The solution, however, does not show up as pink, it at first remains somewhat milky (the powder only dissolves slowly), but when all is dissolved, it is only VERY pale pink, almost colorless, while the concentration is quite high.

Next, I decided to heat the liquid. Sometimes that aids in formation of complexes. When I do that, then I obtain fumes of HCl (driven out of the liquid), and the liquid becomes pale yellow/brown/green. This can only be explained by assuming that a minor part has been oxidized by oxygen from the air (or oxygen dissolved in the 30% HCl).

I added a pinch of Na2S2O8 to the liquid, and quickly the liquid turns dark green/brown, and some chlorine is formed. This shows that the manganese (II) indeed is oxidized to a higher oxidation state.

So, from this I conclude that manganese (II) is VERY pale pink, almost colorless, at least under the conditions I had. Another interesting observation is that the presence of a tiny amount of oxidizing species makes the liquid dark green/brown, exactly the same color, when MnO2 or KMnO4 is reduced by hydrochloric acid. Apparently the reaction, which often is written as MnO2 + 4HCl --> Mn(2+) + 2Cl(-) + 2H2O + Cl2 is much more complex and in reality there is an equilibrium, probably with some chloro-complex involved. For me it was very enlightening to see that the small amount of oxidizer makes the liquid dark green/brown, even in this concentrated acid (hence large excess of HCl).

So, now I'm wondering how I could obtain that deeper pink color. Chemoleo, did you make that manganese chloride, or is it a commercial sample? Of course, I still believe that you have such a deeper pink solution. I only wonder, what is the difference with my situation and why I don't get this pink solution. This is a new chemical riddle, apparently there is more to say about this ;).


[Edited on 16-7-06 by woelen]

Rosco Bodine - 2-10-2006 at 20:33

Just thinking here further concerning the conversion of manganese carbonate to manganese nitrate , and wondering if this could have useful catalytic effect in
lowering the decomposition temperature and accelerating the burn rate of other nitrate oxidizer
plus fuel compositions such as are used for rocket propellants . Ammonium nitrate has been experimented with as an oxidizer , but generally requires something
really hot in the way of a fuel , like powdered magnesium
in order to get a decent burn rate .

What I am thinking is that * if * manganese nitrate
as part of the mixed nitrates was catalytic ...then
perhaps manganese nitrate could be added in small percentage to a tertiary eutectic nitrates mixture , like
the tertiary eutectic in Urbanski 3 , page 257 :

66.5% NH4NO3
21% NaNO3
12.5% KNO3 f.p. 118.5 C

and something like sorbitol could be used as a fuel ,
perhaps with some airfloat charcoal added to the melt
as a further combustion catalyst .

This would provide a very intimate mixture melt castable fuel , having a high percentage of ammonium nitrate
as the oxidizer . The combination of components
might just do the trick of getting the ammonium nitrate
to burn at a reasonably fast rate without the need for
magnesium . If it works , the fuel should have a lot more energy than the sorbitol / KNO3 mixture .

Potassium permanganate is another possible additive
but I think this one could be a dangerous addition in
a melt . Ammonium dichromate is another .

[Edited on 3-10-2006 by Rosco Bodine]

sulfuric acid is the king - 18-8-2018 at 12:02

Can it be produced by mixing aq sol of manganese chloride with sodium carbonate,or manganese sulfate with sodium carbonate?

fusso - 18-8-2018 at 23:45

Quote: Originally posted by sulfuric acid is the king  
Can it be produced by mixing aq sol of manganese chloride with sodium carbonate,or manganese sulfate with sodium carbonate?
Yes