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exodia
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[*] posted on 3-10-2016 at 17:04
MnO2 Madness


Hello there, long time lurker first time poster here, (I did look for a presentation post for newcomers but I couldn't find anything so I jumped in)

I know the MnO2 purification it's been discussed really extensively in the forum, but after reading everything I'm stuck.

years ago when I was in my teens I tried battery crud extraction with supermarket grade HCl, and just the sun in the backyard to cristallice the McCl2 and after 1 recristalization I had a really nice pink cotton candy (without any fancy glassware or equipment)

Now I've been trying and I'm getting frustrated as I have proper apparatus, hotplate, reagents and a tiny bit more knowledge and it's seems impossible, to take that iron off...

Tried HCl and evaporating
Tried Dr Klavonn bisulphate + bisulphite method
Tried oxalic acid method
Tried bubbling SO2 in solution
Tried common ion precipitation with hydroxides (as nurdrage suggested)
Tried washing with acetone
and titrating pH for re-dissolution of Fe ions

And nothing its working, I still get a brownish to strong yellow powder not even close to a high enough purity for a proper recrystallization.

I am running out of ideas, and of ways of doing this.

Couple of things to note:

While I do have battery crud I'm using pottery grade MnO2 of which I have the certificate of analysis (stating 95% MnO2, 4% Fe2O3 and other minor impurities)
I'm not too keen to believe that analysis of 95% purity seeing the results I'm getting.

when washed with acetone, I got a really yellow pale looking product not very useful but the acetone turned emerald green, like nickel green (I know it's highly unlikely that it's nickel, but it sure looks like it)

my three main ways here now are:

1: thermite, but I really don't want to get a 60-70% loss in yield.
2: convert it to other salts I haven't seen discussed in the forum like acetates or oxalates or whatever gives a clear distinction in the two salts properties.
I have seen that iron oxalate solubility it's in the order of mg/100ml but I cannot find anything about Mn (all I found it's "slightly" soluble, which is not a very accurate insight)

3: electrolytic Mn: I'm up for this one I think it'll give really pure Mn, and not a lot of yield loss, and an "easy way" of separating, but I'm completely lost in electrochemistry.

Thank you for the effort of reading the longest post ever (it's the excitement :D)

PD: I'll try to keep posting stuff with my questions and problems(many I'm sure), my progresses, and my answers (if I can provide help in other posts)

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battoussai114
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[*] posted on 3-10-2016 at 17:22


Manganese oxalate has o solubility of around 2*10⁻² g/100ml according to the solubility table I checked.



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[*] posted on 3-10-2016 at 22:47


The analysis certificate can very well be right. If you have 4% Fe2O3 in your material, then you certainly will get yellow/brown material instead of pink material. The color of iron(III) in combination with chloride is MUCH stronger than the color of manganese(II) and even a percent or so of this impurity shadows the color of the manganese(II).

In some sense, battery crud is better than pottery grade MnO2. The impurities of battery crud may have a higher percentage, but the material is more reactive and the impurities are easier to separate. I personally dislike pottery grade MnO2 very much, it always contains a lot of iron. It only is good for pyrotechnic experiments, aqueous chemistry is frustrating with this material.




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[*] posted on 4-10-2016 at 09:22


If you are someone like myself that does a lot of organic chemistry, especially oxidations involving permanganate, you always have a healthy supply of active, pure MnO2 as a byproduct.

One way to purify pottery-grade MnO2 as I have found is to fuse it with molten sodium or potassium hydroxide (you'll obtain a green color characteristic of manganates) and then simply leaching the fused material in water, filtering, and acidifying with baking soda to form permanganate. The permanganate formed in this way is very fragile and will break down in solution to form a precipitate of pure MnO2. Unfortunately this isn't good on anything but a small scale.

Finally, many pottery stores stock manganese(II) carbonate in addition to manganese dioxide, and this can simply be dissolved in the acid of your choosing to get the desired salt.

[Edited on 10-4-2016 by Amos]




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[*] posted on 4-10-2016 at 13:19


Removing iron from Mn salts has been discussed here time and time again.

Here's one method that's been tried and tested by several here, including me.

If you start from impure MnCO3, dissolve it in as little Fe-free HCl(aq) as possible. Neutralise any excess HCl with NH3(aq) as best as you can.

Split the solution into 1/3 (part A) and 2/3 (part B). Set part B aside for now.

Neutralise part A with NH3(aq), precipitating all Fe and Mn as hydroxides. Filter and wash filter cake repeatedly with water to get rid of soluble salts.

Add washed filter cake to part B and allow to stand overnight. Stir from time to time if possible.

Filter the slurry and the filtrate will be be free of Fe.

It works because the pH of the slurry is too high to keep Fe in solution (it joins the precipitate as Fe(OH)3) but low enough to keep Mn2+(aq) in solution, because of the huge difference in solubility products between Fe(OH)3 and Mn(OH)2.

Then precipitate the manganese as MnCO3(s). It should be a nice, light pink.




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[*] posted on 4-10-2016 at 14:33


blogfast, I believe what you've described is what was meant by:
Quote: Originally posted by exodia  
precipitation with hydroxides (as nurdrage suggested)
...as it's the same procedure shown in NurdRage's manganous salts purification video.

exodia: Did you follow the procedure exactly as describe in the video (or by blogfast, above)? If so, then I'd start to wonder if your starting material really is manganese dioxide. Have you tested your starting material? To do so, you could add it to hydrogen peroxide to see if it catalyzes decomposition, and/or fuse it with potassium hydroxide and see if the result is dark green. Your results suggest you may have been given black iron oxide (or something else) instead.
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[*] posted on 4-10-2016 at 15:45


Quote: Originally posted by zwt  
blogfast, I believe what you've described is what was meant by:
Quote: Originally posted by exodia  
precipitation with hydroxides (as nurdrage suggested)
...as it's the same procedure shown in NurdRage's manganous salts purification video.



Correct.

https://www.youtube.com/watch?v=BLJgBSrhZI8




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[*] posted on 4-10-2016 at 19:02


Hello everyone and thank you for the responses.
I did test it with hydrogen peroxide for O2 decomposition and it proved quite effective (it´s quite finely powdered) and I took it as a positive enough test for MnO2 (and I am taking for granted that the analisys sheet its telling the truth about compounds and percentages).

in relation to nurdrage video I followed all steps (mine was the chloride not the sulphate) but ion precipitation should be the same regardless of the salt i think.

my pottery grade had some unsoluble grey stuff wich looked like mud and didnt dissolve in HCl (I suppose thats some fine sand to blend in if you use it for actual pottery) that got separated easily with filtration.

I did it with NaOH as nurdrage´s and not with NH3 as blogfast sugested I´ll give that a try (it shouldn´t change the outcome though) and I will read pH to make sure it´s not still acidic after hydroxides are added (I didn´t pH test it before because I just assumed it would´t be ok as I was dumping a lot of hydroxides in it.

I do want to try the fused KOH as well (that will have to wait until a get a proper torch or bunsen)
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[*] posted on 4-10-2016 at 22:11


I will have to look up my notes (from ages ago). But, as I recall, the process is far more successful with H2SO4 than with HCl. Bloggers did give me a reason but I forget it.

I quite like the SO2 method - and I used the sulfur candle and funnel method that Nurdrage adopted.
But then I found an OTC source of high purity manganese sulfate (as a fertiliser additive) and lost a lot of my interest in playing with battery gunk. I do have some pottery grade MnO2 the same as you and will probably revisit purification at some stage.

J.




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[*] posted on 5-10-2016 at 03:09


Quote: Originally posted by j_sum1  
I will have to look up my notes (from ages ago). But, as I recall, the process is far more successful with H2SO4 than with HCl. Bloggers did give me a reason but I forget it.

J.


Which process are you referring to?




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[*] posted on 5-10-2016 at 03:55


Quote: Originally posted by blogfast25  
Quote: Originally posted by j_sum1  
I will have to look up my notes (from ages ago). But, as I recall, the process is far more successful with H2SO4 than with HCl. Bloggers did give me a reason but I forget it.

J.


Which process are you referring to?


Sorry if I was unclear.
This process of purifying manganese sulfate (that was previously made from battery gunk reacted with H2SO4) is, in my experience, more effective than the similar process on manganese chloride (made from battery gunk and HCl.)

I seem to recall you giving me something of an explanation as to why. But my memory might be playing tricks on me. My lab notes are locked away in a box at present.

Edit: Fixed URL

[Edited on 10-9-2016 by zts16]




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[*] posted on 5-10-2016 at 07:56


Quote: Originally posted by j_sum1  

Sorry if I was unclear.
[url=https://www.youtube.com/watch?v=BLJgBSrhZI8]This process[/url of purifying manganese sulfate (that was previously made from battery gunk reacted with H2SO4) is, in my experience, more effective than the similar process on manganese chloride (made from battery gunk and HCl.)

I seem to recall you giving me something of an explanation as to why. But my memory might be playing tricks on me.


I think it is. There's no good a priori reason to believe sulphate would work better than chloride. The anions are merely spectator ions here, they should make very little difference in the overall outcome. Neither Mn nor Fe form strong complexes with Cl-.




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[*] posted on 9-10-2016 at 07:50


thank you for the responses, I will get on with it shortly (I´m away from the lab at the moment).
I was wondering I maybe separation of Fe and Mn would yield better efficiency if I try to separate the oxides and not the carbonates/sulphates/chlorides ....

maybe some H3PO4 to dissolve the iron oxides before converting into anything else (as far as i know MnO2 doest react with phosporic)
what do you think about this idea (I couldn´t find anything related), I´m just contemplating different approaches that the long discussed here to compare results.
thank you :)
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[*] posted on 9-10-2016 at 08:02


Quote: Originally posted by exodia  
maybe some H3PO4 to dissolve the iron oxides before converting into anything else (as far as i know MnO2 doest react with phosporic)
what do you think about this idea

I think iron phosphates are insoluble.

[Edited on 10-10-2016 by zwt]
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[*] posted on 9-10-2016 at 11:09


that's what I though too, I was thinking that a mixture of iron phosphates and the MnO2 will be easier to react with a solvent or acid that only touches de Fe leaving the MnO2 in situ.
(as far as I see in the web, iron phosphates are insoluble in water and organic solvents) my first choice would be an acid that doesn't react with the MnO2 (or at a really slow rate, and really quick with the Fe)
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[*] posted on 9-10-2016 at 18:27


You said: "some H3PO4 to dissolve the iron oxides," but you seem to know that phosphoric acid may react with, but will not "dissolve", the iron oxides. Strange.
I don't think converting one insoluble iron salt to another insoluble iron salt will help at all.

It should be possible to leach out some of the iron with dilute sulfuric acid, then treat the less-contaminated manganese dioxide with one of the standard methods, like reaction with sulfur dioxide, followed by the hydroxide purification method (if needed).
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[*] posted on 10-10-2016 at 09:29


yeah, my bad, I did mean react.
what I thought as a new approach (Im currently doing all the methods that I posted in my first post to see wich one yields better purity, and annoyingly enough I´m not getting too far, so I´m just exploring different methods)

my idea with the iron phosphates was, that even being insoluble they may react with an acid (I´m not finding too much info with that) quicker and more vigorously than the manganese oxides)

as in for example, having a mixture of iron carbonate (or any carbonate) and manganese oxides with HCl, in an ice bath will minimise the reaction to form MnCl2 and quickly get rid of any carbonates turning them to chlorides)

what I found is that most of the methods discussed here work with leeching Fe out of solution, using the different solubility constants or pH sensitivity to precipitation and so on (wich is not massively acurate due the impossibility for really acurate reading, and more in my case with only chinese pH paper to test :D)
I was trying to find a way of complete reaction due being 2 chemicals with totally diferent properties, as in a phosphate and an oxide.

PS: I may be overthinking this, and probably I´ll just buy good purity grade MnO2, or the chloride for straigth away future uses, but for the ammount of pottery grade I have now its a matter of honour :P
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[*] posted on 14-10-2016 at 12:35


First, an exploration of the underlying electrochemistry per Wikipedia on the zinc–carbon battery (link: https://en.m.wikipedia.org/wiki/Zinc–carbon_battery ), to quote:

"Anode (marked -)
Zn(s) → Zn2+(aq) + 2 e− [E° = −0.7626 V]

Cathode (marked +)
2 MnO2(s) + 2 e− + 2 NH4Cl(aq) → Mn2O3(s) + 2 NH3(aq) + H2O(l) + 2 Cl− [E° ≈ +0.5 V]

Other side-reactions are possible, but the overall reaction in a zinc–carbon cell can be represented as:

Zn(s) + 2 MnO2(s) + 2 NH4Cl(aq) → Mn2O3(s) + Zn(NH3)2Cl2 (aq) + H2O(l)

If zinc chloride is substituted for ammonium chloride as the primary electrolyte, the anode reaction remains the same but the cathode reaction is:

MnO2(s) + H2O(l) + e− → MnO(OH)(s) + OH−(aq)

and the overall reaction:

4 Zn(s) + 8 MnO2(s) + ZnCl2(aq) + 9 H2O(l) → 8 MnO(OH)(s) + Zn(OH)Cl(aq) + 5 H2O +4 ZnO "
---------------

Next, a collection of likely important chemical reactions relating to processing Manganese salts, in particular, MnO, Mn2O3, and MnO2 that may be present in old batteries:

2 HCl (dilute) + MnO → MnCl2 + H2O (see http://chemiday.com/en/reaction/3-1-0-1402 )

4 HCl (conc) + MnO2 → Cl2 + MnCl2 + 2H2O (see http://chemiday.com/en/reaction/3-1-0-177 )

6 HCl (conc, heated) + Mn2O3→ 2 MnCl2 + Cl2 + 3 H2O (see http://chemiday.com/en/reaction/3-1-0-7473 )

2 Mn(OH)2 + Ca(ClO)2 → 2 MnO2 + CaCl2 + 2 H2O (see http://chemiday.com/en/reaction/3-1-0-7503 )

Implying equivalently:

Mn(OH)2 + NaClO → MnO2 + NaCl + H2O

or, more generally, in a reaction paralleling the transition metal iron producing a classic Fenton reaction with H2O2, we also have Fenton-type reactions with the likes of copper, manganese,.. which could proceed as follow with, for example, acidified bleach creating hypochlorous acid, HOCl:

Mn(II) + HOCl → Mn(III) + Cl- + ·OH
Mn(III) + HOCl → Mn(IV) + Cl- + ·OH

where the action of such transition metals, in general, has been proposed to be active in biological systems creating problematic hydroxyl radicals which can attack DNA...(as a reference, see for example, "Nutrition and Immunology: Principles and Practice", edited by M. Eric Gershwin, page 100, link: https://books.google.com/books?id=hf8ICAAAQBAJ&pg=PA100&... ).

One should also expect some HCl formation from the decomposition of the unstable HOCl:

2 HOCl → HCl + O2

which in the presence of HOCl or MnO2 will produce Cl2. The latter in water can create once again HCl and HOCl.

More recent research suggests a reaction between the hydrochloric acid and a hydroxyl radical joined with a single water molecule resullting in the monoatomic chlorine radical:

HCl + ·OH·(H2O) → ·Cl + (H2O)2 (Source: see "The exothermic HCl + OH·(H2O) reaction: Removal of the HCl + OH barrier by a single water molecule" by Guoliang Li, et al, J. Chem. Phys. 140, 124316 (2014); link: http://scitation.aip.org/content/aip/journal/jcp/140/12/10.1... )

Then: ·Cl + Cl- = ·Cl2-

The above being possibly confirmed per one source that reported in 1976 in what is now described as a Fenton-type reaction between the transition metal Vanadium salt V(III) and HOCl plus HCl (from aqueous chlorine). Apparently, in the first stage of the Fenton-type reaction, an oxidation of V(III) to V(IV) and the formation of the dichloride radical ion. Source: "Inorganic Chemistry of the Transition Elements", by B. F. G. Johnson, Vol 6, p. 49, at https://books.google.com/books?id=PQMJLNje2icC&pg=PA49&a...

It should also be noted that in place of inexpensive bleach or acidified bleach (HOCl), one can also employ KMnO4:

3 MnO+ 2 KMnO4+ H2O = 5 MnO2 + 2 KOH

3 Mn2O3 + 2KMnO4 + H2O = 8 MnO2 + 2KOH (See, for example, https://www.google.com/url?sa=t&source=web&rct=j&... )
--------------------------

Next, a look at some published comments on the Zn-MnO2 battery starting with "End-of-life Zn-MnO2 batteries: electrode materials characterization", by Cabral M, et al, published in Environ Technol,2013 May-Jun;34(9-12):1283-95, link: https://www.ncbi.nlm.nih.gov/pubmed/24191461, to quote part of abstract:

"X-Ray powder diffraction allowed for identifying several phases in the electrodes, namely zinc oxide, in the anodes of all the types of saline and alkaline batteries tested, while zinc hydroxide chloride and ammine zinc chloride only appear in some types of saline batteries. The manganese found in the cathode materials is present as two main phases, MnO x Mn2O3 and ZnO x Mn2O3, the latter corroborating that zinc migration from anode to cathode occurs during the batteries lifespan. A unreacted MnO2 phase was also found presenting a low crystalline level. Leaching trials with diluted HCI solutions of alkaline and saline battery samples showed that all zinc species are reactive attaining easily over than 90% leaching yields, and about 30% of manganese, present as Mn(II/III) forms. The MnO2 phase is less reactive and requires higher temperatures to achieve a more efficient solubilization."

Per another source: "Recovery of manganese oxides from spent alkaline and zinc-carbon batteries. An application as catalysts for VOCs elimination", Waste Manag. 2013 Jun;33(6):1483-90. doi: 10.1016/j.wasman.2013.03.006. Epub 2013, to quote from abstract:

"Two different manganese oxides were recovered from the leachate liquor: one of them by electrolysis (EMO) and the other by a chemical precipitation with KMnO4 solution (CMO). The non-leached solid residue was also studied (RMO). The solids were compared with a MnOx synthesized in our laboratory. The characterization by XRD, FTIR and XPS reveal the presence of Mn2O3 in the EMO and the CMO samples, together with some Mn(4+) cations. In the solid not extracted by acidic leaching (RMO) the main phase detected was Mn3O4."

Link: https://www.ncbi.nlm.nih.gov/pubmed/23562448
---------------------------------

Based on the above, one possible method (untested) to obtain a good yield of MnO2 cheaply from old zinc–carbon batteries that are free from Fe contamination would be to first thoroughly rinse to remove soluble zinc salts, ammonia and any of the primary electrolyte. Next, treat with an excess of heated concentrated HCl, in say, an open vessel outdoors, so that liberated chlorine will not form large amounts of HOCl, or in a heated closed system, have the chlorine directed into a cold scrubbing solution of NaOH (creating NaOCl and NaCl). Next, add NaOCl to precipitate MnO2. In the presence of any amphoteric Zn(OH)2 and considering paths to zinc hypochlorite, a possible decomposition leading to the formation of ZnO (or even an insoluble zinc oxychloride, see https://books.google.com/books?id=FnrTAAAAMAAJ&pg=PA671&... ). Finally, rinse the product with dilute HCl as MnO2 requires concentrated HCl to dissolve, but ZnO will dissolve.

In the event of iron contamination, now represented in the presence of Fe(OH)2, the latter dissolves in dilute HCl (see http://chemiday.com/en/reaction/3-1-0-5423 ) and should be addressed in the dilute HCl rinse. Similarly, Fe2O3 dissolves in dilute HCl, albeit slowly (see http://chemiday.com/en/reaction/3-1-0-1486 ) although warming may expedite the process.
-----------------------------------------------------------------------

For those interested in a Fenton-type reaction not based on H2O2 or HOCl, see "Fenton chemistry in biology and medicine*" by Josef Prousek, to quote reaction (15) on page 2330:
"For Fe(II) and Cu(I), this situation can be generally depicted as follows [20,39],

Fe2+/Cu+ + HOX → Fe3+/Cu2+ + .OH + X- (15)
where X = Cl, ONO, and SCN. "

[Edited on 14-10-2016 by AJKOER]
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