Sciencemadness Discussion Board

Soluble Manganese (II) salts from battery MnO2 w/ metabisulfite

Random - 11-8-2013 at 05:16

I have read somewhere a way to reduce MnO2 using metabisulfite, but it also included NaHSO4 which I don't have, neither I do have sulfuric acid. I guess acidic conditions must be present for the reaction to proceed?

I have tried mixing just metabisulfite with washed MnO2 and obtained a small amount of white powder on mixing with NaHCO3 solution. I am not sure if it was so small amount because of small amount of MnO2 used or I haven't met the reaction conditions.

Anyone knows what is NaHSO4 used in the reaction along with metabisulfite? Any substitutes?

plante1999 - 11-8-2013 at 05:25

NaHSO4 is used to dissolve the reduced manganese, any acid that doesn't react with manganese dioxide directly should work.

Random - 11-8-2013 at 05:34

Do you mean reduced manganese as in manganese (II) sulfite? I heard transition metal sulfites are quite insoluble. So adding acetic acid to this would work? Or would it be too weak?

Should I first add metabisulfite then after a while add the acid?

bfesser - 11-8-2013 at 06:54

Quote: Originally posted by Random  
I have read somewhere a way to reduce MnO2 using metabisulfite, but...
Yet you don't provide a reference? Come on, you're not new here, you should know better than that.

[edit] <strong><a href="forumdisplay.php?fid=12">Beginnings</a></strong> is not meant to be a forum to default to when you're feeling too <a href="http://lmgtfy.com/?q=manganese+metabisulfite&l=1" target="_blank">lazy</a> <img src="../scipics/_ext.png" /> to cite references or do your own research. Please don't use it in this way.

[Edited on 11.8.13 by bfesser]

Random - 13-8-2013 at 18:59

I would gladly provide the reference, but all I remember that I have seen it on versuchschemie.net. I can't find it no matter how much I have been searching though. Not speaking german fluently doesn't help either.

On topic, I have washed MnO2 from batteries and dissolved metabisulfite in vinegar and another test tube was filled with water. To both test tubes I added MnO2.

Test tubes cooled after the solvent addition and then became warm again. I guess it was just metabisulfite dissolving.

I'll try to filter both solutions then compare results with NaHCO3(aq). Btw, I have noticed that there is still I assume most of the MnO2 on the bottom of the test tubes.

bfesser - 13-8-2013 at 20:06

<strong>Random</strong>, I've been unable to find it on versuchschemie.de (my German is rusty, and never was great to begin with). I'll keep looking around, and will report back if I find anything of interest.
<a href="http://lanthanumkchemistry.over-blog.com/article-manganese-dioxide-dissolution-103233590.html">
Quote:
Manganese dioxide from an alkaline battery or a carbon-zinc battery is more problematic. In these, large quantities of iron impurities are generally present, which can be difficult to extract. Occasionally, you may find a battery that happens to be free of iron impurities. If so, a colorless solution will be formed upon heating the manganese dioxide with hydrochloric acid. Otherwise, a dirt brown solution will form. This may be remedied by dissolving the manganese dioxide in sodium metabisulfite solution. The sulfite will reduce the manganese(IV) to soluble manganese(II), while the iron(III) remains untouched. <img src="../scipics/_ext.png" />
</a>

[Edited on 14.8.13 by bfesser]

woelen - 14-8-2013 at 11:10

Quote: Originally posted by Random  
I have read somewhere a way to reduce MnO2 using metabisulfite, but it also included NaHSO4 which I don't have, neither I do have sulfuric acid. I guess acidic conditions must be present for the reaction to proceed?

I have tried mixing just metabisulfite with washed MnO2 and obtained a small amount of white powder on mixing with NaHCO3 solution. I am not sure if it was so small amount because of small amount of MnO2 used or I haven't met the reaction conditions.

Anyone knows what is NaHSO4 used in the reaction along with metabisulfite? Any substitutes?

Have a look at the half reactions and then you understand why acid is needed:

MnO2 + 2e + 4H(+) ---> Mn(2+) + 2H2O

The half reaction for metabisulfite is:

S2O5(2-) + 3H2O --> 2SO4(2-) + 6H(+) + 4e

Combine both reactions, such that the number of electrons is the same on both sides. This can be done by taking the first reaction twice and the second reaction once:

2MnO2 + 4e + 8H(+) + S2O5(2-) + 3H2O --> 2Mn(2+) + 4H2O + 2SO4(2-) + 6H(+) + 4e

Simplifying:

2MnO2 + 2H(+) + S2O5(2-) --> 2Mn(2+) + H2O + 2SO4(2-)

As you can see, the net redox reaction consumes H(+) ions. That is the reason why you need to add acid to have this reaction running smoothly. Any reasonably strong acid will do, you can use H2SO4, NaHSO4, HCl, HNO3.

bob800 - 15-8-2013 at 10:20

Quote: Originally posted by Random  
I would gladly provide the reference, but all I remember that I have seen it on versuchschemie.net. I can't find it no matter how much I have been searching though. Not speaking german fluently doesn't help either.


Is this what you're referring to?: https://www.sciencemadness.org/whisper/viewthread.php?tid=14... I believe I saw the same or a similar post by Doktor Klawonn on versuchschemie.

Fantasma4500 - 15-8-2013 at 13:04

ive seen this done by an asian (sounding) guy, using NaHSO3 and NaHSO4 on MnO2, letting it stand overnight then react with NaHCO3 if i recall it properly, i tried it myself..

35g MnO2
x0.625
25g NaHSO3
81.25 NaHSO4
96.875 Na2CO3

these numbers i still have in my 'chemistry' note for random stuff, that im spending time on

anyhow i had all of this stuff mixed up and SOMEHOW SUDDENLY THIS FUCKER decided to decompose into SO2
thats alot, like.. ALOT, before i got it into the bottle to bubble the SO2 through NaHCO3
my respect for SO2 was upped that day, but it will never reach the heights of SO3, although
by what i remember i did not find any results in the solution, no pink colour what so ever
i later tried H2SO4 98% and what not, described in NurdRage's video

talked to my teacher today about it, and he think we should do it in school someday with HCl, and said with HCl 30% might be too low conc.
im really having some immensive rage over MnO2 (:

but hey it must be beyond brilliant for anodes, if it can withstand 98% H2SO4 just like that

a very small scale test with 37% HCl and MnO2 could be done, whereafter neutralized with NaHCO3 to ppt. the MnCO3, as MnCO3 is very insoluble (0.0003somethingsomethingsomething, http://en.wikipedia.org/wiki/Solubility_table)

i believe if you search MnO2 on youtube, find NurdRage's video on the H2SO4 method, the video with the asian sounding dude will be in responses, something with MnCO3

The Volatile Chemist - 18-4-2014 at 12:37

I hate to dredge this post up, but I wanted to ask if NaHSO4 and MnO2 will react without SO2 or other gas production? Side note, MnO2 LOVES to rip oxygen from chemicals. I was able to decompose a dye with it and NaOCl to make benzene, among other things.
So is the reaction I mentioned safe?

bismuthate - 18-4-2014 at 15:04

Generaly MnO2 and NaHSO4 will react in the presence of a reducing agent but they will not react with each other alone.

blogfast25 - 19-4-2014 at 04:43

Quote: Originally posted by bfesser  
<strong>Random</strong>, I've been unable to find it on versuchschemie.de (my German is rusty, and never was great to begin with). I'll keep looking around, and will report back if I find anything of interest.


There are more threads on MnO2/Mn(II) from battery gunge on SM than on anything else. Every conceivable reducing agent has already been discussed. And yet more threads are opened up...