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woelen
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@Metallus: Your post is a very good one, thanks for sharing your thoughts with me. I will do more detailed experiments with dichromates and see what
exactly happens with these in combination with peroxodisulfate and with oxone, both in alkaline and acidic solution. The experiments I did last week
were only simple additions of the salt to acidified dichromate, but I did not really care about the pH of the added solution. In both cases, however,
the liquid remained orange, so there was no redox reaction at all.
@AJKOER: The interaction you mention, I almost certainly observed as well. When the two salts are mixed, then with certain compounds (e.g.
nickel(II)), you see a fairly large amount of oxygen being produced and the final result seems to be destruction of the H2O2 and the destruction of
oxone. This, however, only happens at high pH. But, to be really sure, I also have to do these experiments more carefully.
I will come back to this soon (hopefully this week). My time at the moment is VERY limited, I really wish I had more time.
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Metallus
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Quote: Originally posted by woelen  | @Metallus: Your post is a very good one, thanks for sharing your thoughts with me. I will do more detailed experiments with dichromates and see what
exactly happens with these in combination with peroxodisulfate and with oxone, both in alkaline and acidic solution. The experiments I did last week
were only simple additions of the salt to acidified dichromate, but I did not really care about the pH of the added solution. In both cases,
however, the liquid remained orange, so there was no redox reaction at all.
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That is very unusual. I would expect oxone to reduce strong oxidizers like dichromate. Perhaps not form the blue complex, but still at least reduce it
to green Cr (III).
Caro's acid does reduce permanganate to manganese (II), and since in my eyes an acidified solution of oxone would be more or less that, I would expect
the same. Now, permanganates are stronger than dichromates but I still would expect similar behavior.
I don't know what concentration you used, but for the sake of reproducibility it may be useful to prepare 3 solutions.
A) 0.5g K2Cr2O7 in 5 cc H2SO4 2M x2 (enough H+ for the reduction of dichromate and to provide acidic environment)
B) 0.5g oxone in 2 cc H2SO4 2M
C) 1 cc H2O2 (35%) in 3 cc H2SO4 (98%) (or 0.5cc in 1.5cc, whatever you can afford)
1) Add few drops of B to A1
2) Add few drops of C to A2
C and B should be basically the same thing (similar molar ratio of H2O2/KHSO5 to H2SO4), with the exception that B has some neutral salts in it,
therefore I would expect both to give the same reactions with dichromate.
Since B and C are very strong oxidizers themselves, nothing may happen if it's all too concentrated. In that case, try diluting them with 5-10 cc of
water. Oxone contains a lot of sulfates which may hinder dissolution in concentrated H2SO4 solutions. You may even think of heating it up a bit, even
though it shouldn't be necessary.
If all else fails, it might be interesting to try out KMnO4.
As far as Na2S2O8 is concerned, I would treat it as a completely different animal. The -O-O- bridge that connects the two SO3-
doesn't seem to be "that" reactive and it actually needs a catalyst; perhaps once formed it is quite stable (which is what I experienced in practice).
All SDSs report it as a very dangerous and reactive chemical, but common oxidizers have proven to be way more "explosively" dangerous. I still have to
understand it.
[Edited on 23-4-2018 by Metallus]
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woelen
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From experiments I have done now, I get the impression that at very high concentration of H2SO4, oxone and H2O2 become more similar in their behavior.
I did the following sets of experiments:
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1) Prepare a solution of oxone in 98% H2SO4. Some mild heating was required to get all of it dissolved.
2) Add a few drops of 20%-ish H2O2 to 98% H2SO4, then swirl and put under cold running tap and then add a few drops of 20% oleum to neutralize the
water. The resulting liquid is not fuming. This trick compensates for the fact that I cannot use really concentrated H2O2 (this is forbidden in the EU
without an EP-license, the 20%-ish I made myself by freezing a small amount of 12% H2O2).
Both solutions are kept at room temperature for several minutes.
Next, prepare a solution of K2Cr2O7 in appr. 3M H2SO4.
To part of this solution, I added a few drops of the H2SO4/H2O2 solution. This results in marginal reduction of the dichromate. There only is marginal
production of the dark blue peroxo complex. It is very short-lived and has no intense color.
To another part of the solution of K2Cr2O7 in 3M H2SO4 I added a few drops of the H2SO4/oxone solution. The result was similar. There is marginal
reduction of dichromate and a faint and fast transient appearance of a dark complex (probably the peroxo complex, but it is very short-lived and not
intense).
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Prepare a solution of K2Cr2O7 in 3M H2SO4.
Dissolve some oxone in 3M H2SO4.
Add a few drops of 20%-ish H2O2 to 3M H2SO4.
Allow all solutions to stand for a a few minutes.
Add the solution of oxone to the solution of K2Cr2O7: No visible reaction occurs, the liquid remains orange.
Add the solution of H2O2 to the solution of K2Cr2O7: A very dark blue peroxo complex is formed. This complex quickly decomposes (within 10 seconds)
and a large part of the dichromate is reduced. Finally, the liquid is blue/green if sufficient H2O2 is added, otherwise it becomes olive-green (mix of
dichromate and chromium(III)).
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Prepare solutions of oxone in conc. 98% H2SO4 and of H2O2 in conc. 98% H2SO4 with a few drops of 20% oleum added for compensating the presence of
water.
To both of these solutions, add a small piece of solid CrO3.
In the H2O2-H2SO4 solution, the piece of CrO3 slowly gives off a bright green color and there is bubbling of oxygen. This reaction is very slow.
In the oxone-H2SO4 solution, the CrO3 does not dissolve, the liquid only becomes very pale green.
To both solutions add a big drop of water and swirl. This results in a hissing noise on contact. In both solutions, now the CrO3 dissolves a little
bit more easily. In both solutions, the liquid becomes bright green with a blue hue. So, the CrO3 is reduced to chromium(III) in both solutions. At
these very high concentrations of H2SO4 it looks like the two solutions behave nearly the same.
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Finally, I tried the other way around.
Prepare a solution of chromium sulfate (or chrome alum) in 3M H2SO4. The solution is bluish grey.
To half of this solution add a solution of 20%-ish H2O2 in 3M H2SO4, to the other half of the solution add a solution of oxone in 3M H2SO4.
Both liquids remain bluish/grey, no visible reaction occurs.
Next, heat both liquids, such that they are just boiling/just not boiling. In both liquids, there is bubbling of oxygen. Slowly, the solution with
oxone in it becomes yellowish/green. A large part of the chromium(III) is oxidized to chromium(VI). The solution with H2O2 in it does not change. It
becomes bright green with a blue hue, the well-known color of chromium(III), coordinated to sulfate ion. There is no redox reaction.
Allow both liquids to cool down and then add the H2O2-containing liquid to the oxone-containing liquid. Immediately a dark blue peroxo complex is
formed and a lot of oxygen is produced. Within 10 seconds or so, the blue peroxo complex decomposes and a green liquid remains behind.
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So, from all these experiments one can conclude (with caution) that in water and dilute H2SO4, oxone and H2O2 really are very distinct ompounds,
exhibiting different types of behavior. Oxone is a stronger oxidizer under such conditions, it oxidizes Cr(III) to Cr(VI), albeit with difficulty. It
does not form a peroxo complex with chromium(VI). H2O2 does form a peroxo complex and then by deomposition, it (indirectly) reduces chromium(VI) to
chromium(III).
I also did a simple experiment with NaCl. If you add solid NaCl to a solution of oxone in 3M H2SO4, then there is vigorous bubbling of Cl2. If you do
the same with a solution of H2O2 in 3M H2SO4, then there is only a faint reaction, hardly visible. This also demonstrates that oxone really is
different from H2O2.
Next weekend I hope to have time for a much larger set of experiments with oxone, peroxodisulfate, H2O2, at very low pH and at very high pH, with
dichromate, manganese(II) and nickel. There is a lot to be investigated and discovered with this!
[Edited on 23-4-18 by woelen]
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Metallus
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@Woelen
Very nice results you got there.
I expected them to behave similarly in conc. H2SO4 (basically they are both piranha solution) but behave differently in diluted acid. Afterall classic
H2O2 (<5%) in sulphuric acid is not such a formidable oxidant, compared to its conc. version; quite the opposite actually. One of the reasons why
you can't consider H2O2 (3%) in H2SO4 as piranha.
There are several reactions in which H2O2 acts as a reductant. Only when you pair conc. H2O2 with conc. H2SO4 you get something that is hardly
oxidized. Indeed you only get very faint and slow reaction with the dichromate when you are in conc. solution (since their reduction potentials are
similar).
I've always been interested in the chemistry of H2O2 because that compound gives very different reactions despite its supposedly skyhigh reduction
potential. In this regard, you might find this paper on H2O2 helpful, since it deals in detail about this controversial compound that changes
behaviour depending on
1) acidic/basic environment
2) conc. of the acid/base
3) conc. of H2O2
https://www.scribd.com/document/377242955/Oxidation-Reductio...
If you don't have a sub or whatever, I can PM you.
Getting to the oxone, it is interesting how in dilute H2SO4 nothing happens. I would have expected a reaction to occur, even if just slightly, but
instead you got nothing at all (even though in conc. acid it gave a faint reaction). I would expect that adding H2SO4 in excess would at some point
displace one of the K+, even if just temporarily
KHSO5 + H2SO4 ⇌ H2SO5 + KHSO4
then I would expect the H2SO5 in excess water to partially go back to H2O2 and H2SO4 which in turn would react, even if only slightly, with the
dichromate, but this does not happen. We have further confirmation of this from the fact that oxone in 3M H2SO4 actually oxidizes Cr (III) to Cr (VI)
when suitably hot, so it is reasonable that nothing happens once the Cr (VI) is there.
However, when you do the same experiment in a different modality (that is, mixing the conc. reagents and then add a big drop of water), the reaction
actually occurs. So what I would think is that 3M H2SO4 is not sufficient to displace the K+ and produce H2SO5 which is unstable. On the
other hand, KHSO5 in conc. H2SO4 does produce H2SO5 and when you add water, it doesn't have the time to go back to its KHSO5 form (where the
persulfate ion is "safe") and it quickly decomposes to H2O2 + H2SO4 which reacts with the dichromate (so a matter of kinetics). But now I'm just
speculating.
Anyways, looking forward to your next experiments with manganese and nickel. You might also want to consider lead nitrate and see if you can oxidize
it to Pb (IV). Back in the days I could partially oxidize Pb (II) to Pb (IV) with 3% H2O2 in nitric acid, but this would also result in the bubbling
of oxygen (so H2O2 was both reduced and oxidized, or it simply decomposed). The end result was a bright orange solid (my guess was minium, Pb3O4).
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