Sciencemadness Discussion Board

reaction with Potassium Peroxymonosulfate (oxone)

symboom - 15-12-2016 at 07:20

My new favorite chemical useful in quick isolation of chlorine bromine and iodine just add water :-)

Mixed with sodium chloride made chlorine gas
Mixed with sodium bromide made bromine liquid
Mixed with potassium iodide made solid iodine fast
I guess this makes sence it has an acid and oxidizer

I wonder what would the reactions be with ammonium hydroxide it does make a gas?


Mixed with calcium hypochlorite
I think it makes chlorine dioxide? Am I right
Or chlorine monoxide

If this is true what would form with potassium chlorate
Or perchlorate for that matter


Which is great if you need that was trying to figure a way to mix a salt that gives off oxygen


If mixed with ethanol would it form acetylehyde ethyl acetate or acetic acid how far could it oxidize

Or even acetone would be oxidized to? which I would be afraid of it forming acetone peroxide

[Edited on 15-12-2016 by symboom]

Melgar - 16-12-2016 at 21:08

Quote: Originally posted by symboom  
Or even acetone would be oxidized to? which I would be afraid of it forming acetone peroxide

That seems a very likely route to me, considering the standard synthesis uses sulfuric acid, peroxide, and acetone. Never made it myself, only read about it, but it's definitely something you want to avoid accidentally making.

zed - 17-12-2016 at 16:19

Well, you might be able to form Acetone Peroxides with Oxone, if you went off the beaten path. I dunno.

The combo of Oxone-Acetone is very good for some known reactions. It can perform some very clever oxidations.

https://en.wikipedia.org/wiki/Dimethyldioxirane

MeshPL - 18-12-2016 at 03:08

I guess you could turn chlorate to perchlorate with oxone. Not a very useful reaction, but if you need perchlorates and onlu have chlorates...

Mixing oxone with calcium hypochlorite may produce... singlet oxygen? If the reaction procedes analogously to simple hydrogen peroxide + hypochlorite than yes. If differently, than maybe you will get chlorine dioxide, chlorites, chlorates or something. You need to check this, to check singlet oxygen hypothesis mix reactants in dark, red glow indicates singlet oxygen turning to triplet via chemiluminescent reaction.

symboom - 2-1-2017 at 19:11

Quote: Originally posted by zed  
Well, you might be able to form Acetone Peroxides with Oxone, if you went off the beaten path. I dunno.

The combo of Oxone-Acetone is very good for some known reactions. It can perform some very clever oxidations.

https://en.wikipedia.org/wiki/Dimethyldioxirane


Thanks for that I had no idea
Also reaction with methanol and isopropanol
Has no reaction although the methanol became clowdy
Also added water after I saw no reaction thinking that oxone does not desolve well in alcohols many sulfates have that property

MPS is a versatile oxidant. It oxidizes aldehydes to carboxylic acids; in the presence of alcoholic solvents, the esters may be obtained.



[Edited on 3-1-2017 by symboom]

symboom - 5-1-2017 at 00:11

Oxone with isopropanol forms yellow liquid insouble in the solution
I hope this is not a peroxide
also added sodium bicarbonate and noticed a reaction thinking I made sodium acetate also

[Edited on 5-1-2017 by symboom]

[Edited on 5-1-2017 by symboom]

JJay - 5-1-2017 at 02:48

Is oxone available OTC, or do I have to get it from a chem supplier?

symboom - 5-1-2017 at 04:11

Great question oxone is avaliable as non chlorine shock for pools you probally notice ive been experimenting alot with it seeing what its capable of.


Also sodium hydroxide decomposes oxone to oxygen gas
Now this is interesting

Oxone reaction with sodium carbonate oxygen gas produced and noticable heat


Oxone with sodium bicarbonate CO2 produced solution turned colder didn't seem to allow for a brighter burn may have produced oxygen too but not as much

Oxone and calcium hypochlorite makes oxygen at first because alkalinity decomposes oxone then I think it makes
chlorine dioxide correct me if im wrong.

[Edited on 5-1-2017 by symboom

[Edited on 5-1-2017 by symboom]

JJay - 5-1-2017 at 05:00

Quote: Originally posted by symboom  
Great question oxone is avaliable as non chlorine shock for pools you probally notice ive been experimenting alot with it seeing what its capable of.




Cool. I usually try to pick up at least one new chemical and a new piece of equipment every month. It looks like this month it will be a burette and some oxone.

symboom - 5-1-2017 at 05:46

Quote: Originally posted by MeshPL  
I guess you could turn chlorate to perchlorate with oxone. Not a very useful reaction, but if you need perchlorates and onlu have chlorates...

Mixing oxone with calcium hypochlorite may produce... singlet oxygen?
to answer your question at first it does produce oxygen but not in the excited state no it does not just a face full of chlorine oxides/chlorine still not sure which.

In the dark I tried your suggestion
Oxone with
sodium hydroxide
Sodium carbonate
Sodium bicarbonate
All decompose oxone with no light produced
Percarbonate with oxone no light produced
Also tried percarbonate oxyclean with calcium hypochlorite no singlet oxygen

I will definatly try to see if it will make perchlorate from chlorate
Thanks for the suggestion.

Update potassium perchlorate from chlorate

symboom - 8-1-2017 at 09:37

It seems that it sodium does form potassium perchlorate I have a percipitate formed after putting it in the freezer anyone that can double check that this is indead a route to form potassium perchlorate.

Also a beach like smell arises from the mixture and I added
Sodium bicarbonate to the solution after filtering and bubbles form to check if it is acidic then I remembered it could be the persulfate decomposing

Fantasma4500 - 5-4-2018 at 09:28

this compound is formed through H2O2 and conc H2SO4 which is then neutralized with potassium carbonate, this means it potentially has properties of piranha acid, i suppose peroxysulfuric acid could be baited out from its hiding place with some sulfuric acid?

im interested mostly in turning mine into potassium sulfate to use for double displacement reactions, it says on wikipedia that it decomposes in water, supposedly it loses its oxygen in water?
i poured a bit of 20% NaOH with a scoop of KHSO5 salt and a gas was formed, mysteriously also a bleach like smell (it could have soaked up some HCl from being stored for so many years next to 37% fuming HCl bottles, this could maybe hint at ClO from HCl? didnt smell very pure chlorine like)
anyhow a gas was formed and a wad of cotton was stuffed in the flask, lit on fire, it lit back up once it got down in the flask making it appear likely that it does form oxygen upon neutralization, internet suggests somewhere around 5-8% oxygen content (which, mind you is a decent portion for a solid salt, considering gaseous oxygens density at STP)

i find chlorate to perchlorate <very> unlikely, ferrate - maybe. we have yet to prove that one, alibaba is open for selling 25kg bags of ferrate

the bromine formed, it formed liquid bromine from NaBr sol mixed with KHSO5 sol? this is very noteworthy if true

KHSO5 appears to dissolve copper metal in hot solution, i tried earlier to dissolve copper with NaHSO4 which has a pH of 11 at 1M IIRC, but no luck. this should leave me with KHSO4, CuSO4 and K2SO4 - in which i could then ppt CuCO3*OH (basic copper carbonate) adding potassium carbonate to it, at about 20 euro/kg being readily available it has great potential, but for once im gonna attempt to not be too much of a collector

KIO3 to KIO4 which is done with chlorine gas usually should however be doable with KHSO5, KI would as you say form iodine, but so would KI and chlorine usually? now periodate being 0.17g/100mL 0*C, where K2SO4 is 7.4 would be doable to seperate, KHSO4 easily leeched out with solubility of 36g/100mL at 0

using this oftenly if not always mixed salt that i call "KHSO5", mean to say its about 40% of KHSO5 would prove difficult to accurately calculate amounts with for reactions

it states that peroxymonosulfate is formed from H2O2 and persulfate - reversing that reaction could be interesting

edit: addition--v
actually KHSO5 and water could decompose followingly:
KHSO5 + H2O = KHSO4 + H2O2
now thats quite something, if you keep water levels very low you could end up with highly concentrated H2O2, 80% combusts on contact with many organics.
KHSO5 didnt appear to react very much if at all with alcohols, i didnt try very hard however

[Edited on 5-4-2018 by Antiswat]

symboom - 5-4-2018 at 10:39

I let methanol sit for a couple of days and and it turned into paraformydahyde

Here is a post of the experiment
http://www.sciencemadness.org/talk/viewthread.php?tid=71716#...

Im glad to see other try out tge experience and i hope it helps others
Oxone is very powerful oxidizer a mix of an acid and oxidizer
It will generatee chlorine gas from sodium chloride

And calcium hypochlorite and oxone im woried to form chlorine dioxide but hopefully its oxygen

Chlorites with oxone would be interesting

A side thought fluorine gas can form potassium peroxysulfate from potassium sulfate which would be getting oxidized
If that gets you imagination

Solutions can be very concentrated due to both being solids and reaction is started by only water desolving both

[Edited on 5-4-2018 by symboom]

[Edited on 5-4-2018 by symboom]

AJKOER - 6-4-2018 at 16:28

One can activate peroxymonosulfate with mineral rich tap water, or try carbonated mineral water.

Apparently, however, one of the best way to activate peroxydisulfate (PDS) and peroxymonosulfate (PMS) is siderite, (Fe,Mg,Ca,Mn,Zn,Co)CO3, where Fe, Mn and cobalt are the key metal ions (apparently, a trace of the toxic cobalt ions in the presence of HSO5- is capable of producing the powerful sulfate radical anion, see "COBALT/PEROXYMONOSULFATE AND RELATED OXIDIZING REAGENTS FOR WATER TREATMENT" a thesis by Georgios P. Anipsitakis, https://etd.ohiolink.edu/rws_etd/document/get/ucin1130533674... ). The ensuing advanced oxidation process can then produce both the usual sulfate radicals (see https://www.sciencedirect.com/science/article/pii/S004313540... ) and the even more powerful hydroxyl radical (see "Activation of Persulfates Using Siderite as a Source of Ferrous Ions: Sulfate Radical Production, Stoichiometric Efficiency, and Implications", at https://pubs.acs.org/doi/abs/10.1021/acssuschemeng.7b03948 ).

The siderite functions as a source of Fe2+ (think fenton with KHSO5 replacing H2O2) with the sulfate radicals as the major active radical. In place of siderite, also consider FeS2 (Fool's Gold, see for example, https://pubs.acs.org/doi/abs/10.1021/ie100740d to which I would also pump in air/O2 to expand possible reactive oxygen species creation. Further, I would suspect adding the Bravoite variation of Fool's Gold, which contains toxic nickel and cobalt, would be favorable).

Note, the action of the sulfate radical anion, for example, on chloride (or other halogens, in reverse order of their reactivity, so the action on iodide is faster than with bromide...) proceeds as follows:

.SO4- + Cl- --> SO4= + .Cl

.Cl + .Cl --> Cl2

On hypochlorite:

.SO4- + OCl- --> SO4= + .ClO

.ClO + .ClO --> Cl2O2 or Cl2 + O2

With chlorite, (explosive) chlorine dioxide:

.SO4- + ClO2- --> SO4= + .ClO2

On bicarbonate, the carbonate radical:

HCO3- + .SO4- --> HSO4- + .CO3-

[Edited on 7-4-2018 by AJKOER]

DavidJR - 7-4-2018 at 20:38

Chemplayer did a video on producing bromine from NaBr with a persulphate as the oxidizer: https://www.youtube.com/watch?v=Au2Hb6XKohE

I can also confirm from personal experience that NaBr + H2SO4 + H2O2 (i.e. NaBr + piranha solution) results in very fast evolution of elemental bromine.

RogueRose - 7-4-2018 at 21:31

I found this at a local hardware store as a pool chem, so I looked for some other manufacturers and it seems that many brands list the content as 32 - 64% oxone (depending upon brand), the rest "inactive" ingredients. I did find one that listed 92%+ Potassium Peroxymonosulfate and it was actually priced better lb/lb than the lower concentrations.

When you were testing your oxone, did you use something from a chem supplier or from pool/spa chemicals? There are a lot of "inactive" ingredients in the pool chems, some of them containing 8+ compounds in quantities of 3 - 26%. has anyone found a relatively pure supplier of this or brand"?

XeonTheMGPony - 8-4-2018 at 04:28

cheaper ones tend to be purer htt spa chems all list just 9x purity how much that can be trusted how ever

woelen - 8-4-2018 at 22:54

Yesterday I ordered some "Bellaqua Sauerstoff Granulat" from amazon.de. If you look at the MSDS of this stuff, then it only lists one chemical (oxone: K2SO4.KHSO4.2KHSO5) and in the section about mixtures it tells that it does not apply. So, this must be quite pure stuff. The price is appr. EUR 10 per kg, plus shipping. If you live in Germany, then you just pay a few euros for shipping, I had to pay nealry EUR 14 for shipping :(. In NL you cannot buy this stuff anymore, this time not because of environmental concerns or terrorosm, but because it only is an incomplete treatment of swimming pools and does not kill all bacteria. If people only use oxone, then they risk trouble with their swimming pools. So, in NL we only can buy complete treatments, being all kinds of variations of chlorine-based products (TCCA, Na-DCCA, Ca(OCl)2).

I already have very impure oxone, purchased locally as an oxygen treatment for swimming pools many years ago, but that is a blend with two ingredients (oxone and K2S2O8). I want the pure oxone. I do not know what I will do with the blend. The blend is very stable on storage, and not hygroscopic at all. It still is as good as when I purchased it. I hope that pure oxone also will be easy on storage.

zed - 9-4-2018 at 16:49

I saw Oxone for sale recently at Walmart.

Much less stock, than in years past.

Seems like Chlorides of Isocyanuric Acid, have become more available. While Oxone is fading away.

Though perhaps, it is just a local phenomena.

AvBaeyer - 9-4-2018 at 19:37

Useful information on the utility of oxone in organic synthesis:

http://reag.paperplane.io/00002385.htm

AvB

woelen - 10-4-2018 at 12:37

I just received oxone this afternoon and did some exploratory tests:

- with solid NaCl and a few drops of water you get chlorine gas
- with solid NaBr and a few drops of water you get bromine. The mix heats up a lot and the bromine boils away as gas and condenses on the glass and runs back in little drops of liquid. Nice to see this violent reaction.

An interesting reaction occurs with nickel(II) at high pH, which clearly demonstrates the difference with free peroxide:
- Dissolve some nickel sulfate (or the nitrate, but not the chloride) in water.
- In a separate test tube, dissolve quite some NaOH or KOH in water.
- In a third test tube, dissolve some oxone (be sure to have excess hydroxide, relative to the oxone).
- Mix the solution of NaOH or KOH and the solution of oxone. The resulting liquid must be strongly alkaline. You see very fine bubbles of gas (oxygen) in the solution.
- Add the solution of the nickel(II) salt.
The result is a black precipitate (really black, like carbon black) and a lot of bubbles of gas (oxygen). The precipitate remains black, also when production of oxygen stops.

A similar control experiment was done with H2O2 instead of oxone:
- Dissolve some NaOH or KOH in water and add a few drops of 10% H2O2.
- In a separate test tube dissolve some nickel sulfate or nickel nitrate.
- Mix the two solutions
The result is a pale green precipitate of Ni(OH)2. No special reaction occurs.
I tried with a very small amount of H2O2 and an excess amount of H2O2, relative to the nickel. The result is the same in both tests.

Finally, I added a few drops of 10% H2O2 to the test tube with the black precipitate. This results in immediate destruction of the black precipitate. Bubbles of oxygen are produced and the precipitate becomes pale green.

In a final experiment I added the liquid with green precipitate from NiSO4 and H2O2/NaOH solution to the liquid with black precipitate from NiSO4 and oxone/NaOH solution. This also results in immediate destruction of the black precipitate and only a lot of green nickel(II) hydroxide remains.

So, one can conclude that oxone at high pH and hydrogen peroxide at high pH are very different species, which give totally different reactions with nickel(II).

A similar result also is obtained with K2S2O8. This also gives a black precipitate with nickel(II) at high pH. This black precipitate also is destroyed by H2O2.

Interesting stuff and a lot to be investigated: H2O2 vs. S2O8(2-) vs. SO5(2-)

woelen - 12-4-2018 at 11:48

Another interesting experiment with oxone, active ion is HSO5(-), demonstrating the difference with peroxodisulfate, active ion is S2O8(2-).

- Dissolve some AgNO3 in 50% HNO3. The AgNO3 dissolves with difficulty, some heating is required to get a decent amount dissolved.
- Allow the solution to cool down or cool it under a running cold water tap before proceeding further.
- Split the solution in two equal parts.

- To one half of the solution of AgNO3 add a small pinch of oxone: The oxone dissolves, also with some difficulty, the liquid remains colorless.
- To the other half of the solution of AgNO3 add a small pinch of Na2S2O8 or K2S2O8: The solid slowly dissolves and it makes the solution dark brown. A deep brown nitrato complex of silver(III) is formed.

- To the colorless solution with oxone also add a little Na2S2O8: The solid partly dissolves, the liquid remains colorless.
- To the brown solution with Na2S2O8 add a little oxone: The oxone dissolves and the dark brown color almost instantly disappears. There is production of a little amount of a colorless gas (most likely this is oxygen).

So, peroxodisulfate gives a brown complex of silver in nitric acid, while oxone does not give such a brown complex. Even more so, the oxone destroys the brown complex.

It is remarkable that two peroxo sulfates show such different behavior, and also that there is such a difference with H2O2, while both compounds are said to produce H2O2 when mixed with water. I cannot confirm this in my experiments. E.g. there is no formation of a blue peroxo complex with chromium(VI) in acidic solution, not with oxone, nor with peroxodisulfate.

[Edited on 12-4-18 by woelen]

AJKOER - 14-4-2018 at 03:24

Woelen:

Some interesting experiments!
------------------------------------------

Also, just came across this study, "Oxone/Co(2+) oxidation as an advanced oxidation process: comparison with traditional Fenton oxidation for treatment of landfill leachate", by Sun J, Li X, Feng J, Tian X.

My opinion on the use of Oxone/Cobalt(ll) as an AOP has been raised, see https://www.ncbi.nlm.nih.gov/pubmed/19595430 . To quote:

"From this work, it can be concluded that Oxone/Co(2+) oxidation process demonstrated higher degradation efficiencies of the COD, SS and color for landfill leachate treatment than that by Fenton oxidation process. "

[Edited on 14-4-2018 by AJKOER]

Metallus - 17-4-2018 at 23:52

Quote: Originally posted by woelen  
Another interesting experiment with oxone, active ion is HSO5(-), demonstrating the difference with peroxodisulfate, active ion is S2O8(2-).

- Dissolve some AgNO3 in 50% HNO3. The AgNO3 dissolves with difficulty, some heating is required to get a decent amount dissolved.
- Allow the solution to cool down or cool it under a running cold water tap before proceeding further.
- Split the solution in two equal parts.

- To one half of the solution of AgNO3 add a small pinch of oxone: The oxone dissolves, also with some difficulty, the liquid remains colorless.
- To the other half of the solution of AgNO3 add a small pinch of Na2S2O8 or K2S2O8: The solid slowly dissolves and it makes the solution dark brown. A deep brown nitrato complex of silver(III) is formed.

- To the colorless solution with oxone also add a little Na2S2O8: The solid partly dissolves, the liquid remains colorless.
- To the brown solution with Na2S2O8 add a little oxone: The oxone dissolves and the dark brown color almost instantly disappears. There is production of a little amount of a colorless gas (most likely this is oxygen).

So, peroxodisulfate gives a brown complex of silver in nitric acid, while oxone does not give such a brown complex. Even more so, the oxone destroys the brown complex.

It is remarkable that two peroxo sulfates show such different behavior, and also that there is such a difference with H2O2, while both compounds are said to produce H2O2 when mixed with water. I cannot confirm this in my experiments. E.g. there is no formation of a blue peroxo complex with chromium(VI) in acidic solution, not with oxone, nor with peroxodisulfate.

[Edited on 12-4-18 by woelen]

The dark brown compound you obtained with reacting AgNO3 with Na2S2O8 is reported to be a mix of silver I and III oxide with the formula Ag4O4 http://en.wikipedia.org/wiki/Silver(I,III)_oxide . You saw no formation of this compound with the oxone alone, but observed that adding the oxone to the brown solution would bleach it. The presence of transition metal ions (such as Ag2+) in solution catalyses the decomposition of the oxone to give sulfate and oxygen. The brown complex is a mix of Ag I and III, but the compound is solube in nitric acid as Ag2+. This would explain why you see the decomposition of the oxone when you mix it to the brown complex, but see nothing when it's mixed with the simple Ag+.
http://www.cruso.es/FICHAS/33Oxone.htm
https://www.researchgate.net/publication/279889195_The_free-...
https://pubs.acs.org/doi/abs/10.1021/ic50218a012

As to why the oxone bleaches it, I would innocently think that the H2SO5 behaves as a mix of H2SO4 + H2O2 where the latter acts as a reductant in those conditions. H2O2 has a very wild behaviour depending on the situation and this has troubled me very often. However, recently I came across a paper that erased most of my doubts about it.

You have carried out experiments on nickel (II) salts and observed the different behaviours between oxone, H2O2 and Na2S2O8 in basic medium. I would think that the oxone and peroxodisulfate are strong enough even in basic environment to oxidize NiII (green) to NiIII (black) but the H2O2 alone might not.

While in acid solution one might even consider H2SO5 (which is the result of acidification of the oxone) to behave as a mix of H2SO4 + H2O2, in basic solution I'd imagine that the peroxide is "locked" on the sulfate ion, since it's present as a salt. Afterall oxone is sold as the potassium salt of H2SO5 because it is more stable. This would explain why the oxone behaves differently from H2O2 in basic solution (where there is no formation of peroxosulfate ion), but similarly in H2SO4 medium; after all during your last test (the bleaching of the brown complex with H2O2), you had H2O2 in a solution of HNO3 and sulfate ion SO42- (originating from the reduction of the Na2S2O8). You have H+, NO3-, HSO4- and H2O2 + some leftover neutralized salt that remains in solution. Isn't that basically the same situation of an acidified solution of oxone? This would explain why in acidic environment in the presence of H2SO4 (no matter how it's originated), oxone and H2O2 give out the same reactions.

There is still one unanswered question: if acidified oxone and piranha solution (H2SO4 + H2O2 mix) are indeed equivalent as I claim above, why did you observe different reactivity with CrVI (no formation of the blue complex)? In order to answer this, I'd like to ask some detail about your experiment, because when I was playing with dichromate and hydrogen peroxide back in the days, there were some situations in which the blue complex would not form at all and the solution would turn instantly green. This usually depended on the acid concentration or in the order in which I added the reactants (e.g. 1 drop of H2O2 in a pool of acidified dichromate or viceversa).

By the way, something we can certainly agree on is that the peroxodisulfate is another beast and behaves very differently (e.g. it works perfectly even in neutral solution and is way more stable than it may look like. Even its reactions aren't as violent as the ones you get with classic oxidizers)

[Edited on 19-4-2018 by Metallus]

AJKOER - 22-4-2018 at 08:05

Apparently, H2SO5 behaving like a mix of H2O2 + H2SO4 can be a problem as H2O2 can interact with H2SO5 in the presence of select metals like platinum.

See discussion at https://www.researchgate.net/post/Can_water_be_a_reducing_ag... , and in particular reference to, “The effect of colloidal platinum on mixtures of Caro’s per-sulphuric acid with hydrogen peroxide.” by Thomas Slater Price and John Albert Newton Friend, Proc. Chem. Soc., London, 1904, 20, 173-212 (p 187).

woelen - 22-4-2018 at 11:21

@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.

Metallus - 23-4-2018 at 06:04

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.


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]

woelen - 23-4-2018 at 11:42

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:

------------------------------------------------------------------------------------

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).

----------------------------------------------------------------------------------

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]

Metallus - 24-4-2018 at 01:53

@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).