Sciencemadness Discussion Board - Reduction of permanganate by ascorbic acid → surprise!

Reduction of permanganate by ascorbic acid → surprise!

Keras - 2-10-2022 at 04:24

Folks,

I was planning to make a few µg (yes, µg) of activated MnO₂ to test its oxidising properties with xylene. I was pretty pleased by what I obtained using sodium thiosulphate: a crystal of sodium thiosulphate added to a 25 mL beaker containing water into which a few µg of KMnO₄ have been dissolved turns it immediately to a brownish Mn²⁺ solution, to which a few beads of NaOH are added to make it precipitate. By the way: manganese dioxide is reputed insoluble. So what Mn²⁺ compound makes the water brownish?

The fact that the precipitate is very finely divided is obvious, given the volume it takes in the water. One could easily guess there are several milligrammes held in suspension, but when the water is evaporated, nothing is left apart from very light traces of powder on the beaker, so the real amount is indeed in the order of a few µg.

Mixing one drop of benzyl alcohol with the (manganese dioxide??) in water thus produced gives off a very faint scent of almond, so there must indeed be some form of oxidation.

Now, I remembered that reducing Cu²⁺ ions with ascorbic acid leads to the formation of copper nanoparticles, so I wondered if the same held true for manganese. But… it turned out that a solution of permanganate, on addition of ascorbic acid, becomes totally… colourless (see attached video). Just as if all ions suddenly evaporate. I suppose there must be some chelating action at work, because the solution assumes a transient brown tint just before clearing up.

I’m going to try and see if that solution can oxidise benzyl alcohol into benzaldehyde, too. It turns out there are certainly available Mn²⁺ ions, because an addition of a few beads of sodium percarbonate results in a very clear bubbling off of oxygen from the beads.

If anyone has a clue about what really happens to those manganese ions, I’m all ears

Attachment: phpn1XQqR (7MB)
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[EDIT: Indeed, when I add a smaller amount of ascorbic acid, the solution turns brown and manganese dioxide precipitates out. On addition of more ascorbic acid, the precipitate dissolves and the solution is left totally clear.]

[Edited on 2-10-2022 by Keras]

j_sum1 - 2-10-2022 at 05:59

The brown is very fine suspended particles of MnO2 or Mn2O3.
It is still able to reduce to Mn2+ which it does on addition of asvorbic acid. This also explains its ability to oxidise tge benzyl alcohol. Mn2+ is a light pink, almost colourless.

I think that explains all your observations.

Keras - 2-10-2022 at 06:28

So, the upshot here is that ascorbic acid shouldn't be used to reduce Mg(VII) compounds, unless you want to remove spots or stains, because it goes all the way down to Mg(II), which is then more or less inert (and in any case not able to oxidise anything by turning into Mg(0)).

On the other hand, sodium thiosulphate or sodium sulphite gracefully end up their reducing action at Mg (IV).

Frankly, I would've expected sodium sulphite to be a more powerful reducing agent than ascorbic acid.

[EDIT] I remade manganese dioxide using sodium thiosulphate as the reducing agent. Adding a drop of BzOH or even ethanol doesn’t seem to alter the thin precipitate in anyway. But as soon as I add ascorbic acid, it literally vanish into nothingness, almost magically. I’ll try using CH₂Cl₂ as the solvent instead.

[Edited on 2-10-2022 by Keras]

Boffis - 2-10-2022 at 08:20

I use ascorbic acid to remove black Mn oxides from mineral specimens and samples for analysis. It is amazingly efficient an as you have observed, the product quite colourless. Once all of the ascorbic acid has been consumed the solution is slightly unstable and slowly oxidises again and turns a mauvish brown colour. I presume that the ascorbic acid is oxidized to some acid that then combines with the reduced Mn2+.

So yes, ascorbic acid can reduce the MnO2 you are trying to prepare but it could probably still be used if you are careful about the amount you add. It sounds like you used an excess. Try using excess KMnO4.

Bedlasky - 2-10-2022 at 10:04

Permanganate is reduced in neutral solution to MnO2. If you want to reduce it in to Mn2+, you need acidic environment. But it is little bit pointless to make Mn2+ solution and than add base and H2O2 to make MnO2. Just reduce KMnO4 directly in neutral solution and filter.

Keras - 2-10-2022 at 11:11

Quote: Originally posted by Bedlasky

Permanganate is reduced in neutral solution to MnO2. If you want to reduce it in to Mn2+, you need acidic environment.

So you mean the fact that ascorbic acid reduces permanganate to manganous ion does not stem from a particularly efficient reducing property, but just because besides being a reductant, it is also acidic?

Bedlasky - 2-10-2022 at 22:05

Quote: Originally posted by Keras

Quote: Originally posted by Bedlasky

Permanganate is reduced in neutral solution to MnO2. If you want to reduce it in to Mn2+, you need acidic environment.

Yes, exactly.

Keras - 2-10-2022 at 22:14

Quote: Originally posted by Bedlasky

Yes, exactly.

Ok, thanks

Texium - 3-10-2022 at 05:44

You can see the same effect by adding sodium metabisulfite to a permanganate solution. You’ll get manganese dioxide precipitate at first. Then add a few drops of sulfuric acid (not reducing in its own right) and you’ll see the oxide disappear as it gets reduced to Mn(II).

Keras - 4-10-2022 at 01:54

Quote: Originally posted by Texium

You can see the same effect by adding sodium metabisulfite to a permanganate solution. You’ll get manganese dioxide precipitate at first. Then add a few drops of sulfuric acid (not reducing in its own right) and you’ll see the oxide disappear as it gets reduced to Mn(II).

So, what's the reaction here?
MnO₂ + 2H⁺ → Mn²⁺ + ????

Bedlasky - 4-10-2022 at 02:48

You forgot metabisulfite.

2MnO2 + (S2O5)2- + 2H+ --> 2Mn2+ + 2(SO4)2- + H2O