I saw the creation of the dioxide from the sulfate using potassium peroxymonosulfate from NurdRage and I am thinking. Could it be possible to just use
sodium hydroxide to make the hydroxide and then oxidize to the 4 state that with bleach?
Sorry if the question is stupid. It is not so easy to get dioxide in pure quantities here and the battery extraction seems like a pain.
My ultimate goal is to get activated manganese dioxide, but all of those require, ironically enough, to make the permanganate first, all of those
always ask for the dioxide.
Is there a shorter way to go from the sulfate to activated manganese dioxide in any case?
Otherwise first I would need to make the dioxide, then the permanganate to then the activated dioxide finally.helzblack - 27-9-2024 at 09:15
Wikipedia does say that using the sulfate and hydroxide and an oxidizers would work but I am not sure...Random - 27-9-2024 at 10:55
So you dissolved manganese dioxide in potassium metabisulfite and acetic acid? I’ve seen a method using sulfur dioxide as a reducing agent, using
your method would generate it in situ.clearly_not_atara - 27-9-2024 at 12:30
Quote:
My ultimate goal is to get activated manganese dioxide, but all of those require, ironically enough, to make the permanganate first, all of those
always ask for the dioxide.
MnO2 is not always activated, and generally the method of preparation is important. If you're trying to make activated MnO2, you should worry about
making MnO4- first and then use the established method.
In order to make permanganate (Mn7+, purple) you basically make manganate (Mn6+, green) by oxidizing in strongly basic conditions and then bring the
pH to neutral, which causes manganate to disproportionate into permanganate and MnO2 (s). The pKa for the reaction:
MnO4(2-) + H+ << >> MnO4H-
is 7.4, so below pH 7.4, manganate rapidly disproportionates.helzblack - 27-9-2024 at 13:00
Ok, so it is important to keep alkaline conditions to allow the formation of the manganate while we are oxidizing to then disproportionate to the
permanganate.
Correct?
But while this works for the MnIV state can the oxidization be carried from the MnII of the sulfate all the way to the sixth?
I wanna use chlorate as the oxidizer, is that even viable? Do I need something else? I am just worrying about the other ions interfering.
Sorry for my ignorance.Random - 27-9-2024 at 13:23
So you dissolved manganese dioxide in potassium metabisulfite and acetic acid? I’ve seen a method using sulfur dioxide as a reducing agent, using
your method would generate it in situ.
I have not completely read this Lab NoteBook today which I wrote inside 20i3 as I uploaded the images. That was more than 10 years ago in Summer of
20i3.
...
I have perfected the procedure. I still have Manganese... is it Carbonate? Made from the procedure.
Edit:
I just found a plastic storage container with Manganese II Carbonate written on it. I checked inside and there is just a little bit of this brownish
powder.
According to my memory what I made was pink-white-yellow.
...
You see what happens in 10 years.
I was 16 in 20i3.
[Edited on 27-9-2024 by Random]
[Edited on 27-9-2024 by Random]
Edit:
I really worked on the procedure. I still remember as I was doing it.
...
I was really skilled for 16 years old in Chemistry. Unfortunately I did not have most of the Equipment.
[Edited on 27-9-2024 by Random]
Edit:
I am not sure if this is egzact Manganese II Carbonate I was now looking for. According to my personal memory it was as I said pink-white-yellow. And
there was more of it, at least according to my personal memory.
[Edited on 27-9-2024 by Random]Sir_Gawain - 27-9-2024 at 16:29
Manganese carbonate turns brown over time due to oxide formation. Random - 27-9-2024 at 21:20
Ok, so it is important to keep alkaline conditions to allow the formation of the manganate while we are oxidizing to then disproportionate to the
permanganate.
Correct?
But while this works for the MnIV state can the oxidization be carried from the MnII of the sulfate all the way to the sixth?
Read the permanganate thread, particularly the early pages. A couple of people did experiments with air and/or nitrate oxidation of manganese with
alkali. IIRC, aerobic oxidation can take you all the way to Mn6+ (manganate) in the presence of alkali. The manganate should then disproportionate
when the pH is lowered and KMnO4 precipitates with excess K+ at low temperature.
