Sciencemadness Discussion Board

Can't find a Peroxy acid mechanism of formation

SunriseSunset - 29-10-2015 at 22:09

This question is killin' me. I don't know why, but I cannot find an example of the mechanism to form ANY peroxy acid from ANY said carboxylic acid and hydrogen peroxide... Am I being dumb or something, is the answer to this question really so simple that there are no mechanisms needed for examples -___-


.. could really use some enlightenment thx

[Edited on 30-10-2015 by SunriseSunset]

SunriseSunset - 29-10-2015 at 22:41

IS there not a mechanism? is it a redox reaction?

IrC - 29-10-2015 at 23:22

Did you search? https://en.wikipedia.org/wiki/Peroxy_acid

"Several organic peroxyacids are commercially useful. They can be prepared in several ways. Most commonly, peracids are generated by treating the corresponding carboxylic acid with hydrogen peroxide:

RCO2H + H2O2 <> RCO3H + H2O"

SunriseSunset - 30-10-2015 at 00:05

Yes of course lol but that doesn't say what's going on there. It's just a balanced equation, not a mechanism or the type of reaction it is. It just says "treated" .. But what is really going on? o.O

I want to know if (aq)H2O2 dissociates to +H3O and -O2H in solution.
^ scratch this, H2O2 is too weak of an acid

I basically want to know if any of the conjugate species play a role as an intermediate species to form the final peroxy acid. such as the formate anion?

I don't know the actual steps for the mechanism. It would be great to have this cleared up!!!!!

[Edited on 30-10-2015 by SunriseSunset]

SunriseSunset - 30-10-2015 at 00:07

Yes, it does at least tell me water is produced as a side product. But it doesn't say how the peroxy acid is formed. What I'm trying to say, is there are no descriptive layouts to teach the formation of peroxy acids online. No mechanism information or steps involved.
Everything you search for just brings up the action of peroxy acid on alkenes to form epoxides, it's rather stressful

I can't figure out what is the electrophile or nucleophile

Maybe I'm making this more complicated than it needs to be. Could someone please help me


[Edited on 30-10-2015 by SunriseSunset]

AJKOER - 30-10-2015 at 12:12

Here are some brief comments and sources I cited on a recent thread, that you may find helpful:
Quote: Originally posted by AJKOER  

...........
Another experiment (which I plan on performing) is saturating a solution with some SO2 and lots of N2O, and then treating the mix in the presence of added O2 to photolysis or pulse radiation from a microwave. The expected product could include the aqueous HSO5- anion, or H2SO5 (assuming it has not otherwised reacted, see https://books.google.com/books?
https://books.google.com/books?id=vVvrCAAAQBAJ&pg=PA143&... ), which can be used to create persulfate salts.

Reference, see discussion on the sulfite radical on page 6 and the middle of page 7 at https://www.google.com/url?sa=t&source=web&rct=j&... and also http://pubs.acs.org/doi/abs/10.1021/jp011255h .

[Edited on 11-9-2015 by AJKOER]


Some specifics:

2 SO2 + 2 H2O = 2 H2SO3 = 2 HSO3- + 2 H+
HSO3- + •OH → •SO3- + H2O
•SO3- + O2 → •SO5-
•SO5- + HSO3- → •HSO5- + •SO3-
1/2 •SO3- + 1/2 •SO3- → 1/2 S2O6(2-)
-------------
Net: 2 SO2 + H2O + •OH + O2 → 2 H+ + •HSO5- + 1/2 S2O6(2-)

[Edit] There is also possible a second reaction occurring for the 4th reaction above, but per the reaction constant found in Table 2 on page 135 of "Heterogeneous and Liquid Phase Processes: Laboratory Studies Related to ...", edited by Peter Warneck, link:
https://books.google.com/books?id=vVvrCAAAQBAJ&pg=PA138&... ,it occurs only 4% of the time (.36/(.36+8.6)) given by:

•SO5- + HSO3- → SO4(2-) + •SO4- + H+

Note, the required hydroxyl radical employed above, can be provided either by the action of either sunlight or pulse radiation on N2O in the presence of water, for example, as I have documented previously on SM. Also, hydroxyl radicals can be generated from the photolysis of aqueous nitrate or nitrite per the reactions:

NO3-(aq) + hv → •NO2 + •OH (see http://pubs.acs.org/doi/abs/10.1021/ja073609 )

Or, NO2-(aq) + hv → NO + •OH

[Edited on 31-10-2015 by AJKOER]

Darkstar - 30-10-2015 at 19:57

I haven't studied this reaction at all, but my guess is that the H2O2 first disassociates into two hydroxyl radicals:

HOOH → HO• + •OH

Then one hydroxyl radical reacts with the carboxylic acid and removes its acidic proton, cleaving the oxygen-hydrogen bond homolytically and producing a resonance-stabilized carboxyl radical and a water molecule:

RCOOH + •OH → RCOO• + H2O

Finally, the second hydroxyl radical combines with the carboxyl radical to form the peroxy acid:

RCOO• + •OH → RCO3H

Another possible route I thought of could be something like the carbonyl oxygen on the carboxylic acid getting protonated followed by a nucleophilic attack on carbon by H2O2. The intermediate then undergoes a proton shift to convert one of the hydroxyl groups to an oxonium group, which then leaves as water as a double bond forms between carbon and the other hydroxyl group oxygen. The charged carbonyl oxygen is then deprotonated to give the peroxy acid.

If it helps you visualize what I mean, this is a post that I made recently in another thread showing the mechanism for Fischer esterifications. Just replace the ethanol in that reaction with H2O2.

[Edited on 10-31-2015 by Darkstar]

SunriseSunset - 31-10-2015 at 06:52

Quote:

I haven't studied this reaction at all, but my guess is that the H2O2 first disassociates into two hydroxyl radicals:HOOH → HO• + •OH


What could cause this to happen?

For how reactive the radical is, I would think it would take a lot to split that sigma bond

[Edited on 31-10-2015 by SunriseSunset]

Darkstar - 31-10-2015 at 12:04

Quote: Originally posted by SunriseSunset  
What could cause this to happen?


Light, heat, harsh language . . . pretty much any input of energy can potentially cause the bond to cleave homolytically. The oxygen-oxygen single bond in peroxides is extremely weak for the same reason the bonds in diatomic halogens are weak. The similar electronegativity between the two oxygen atoms makes the sigma bond between them easily breakable. Why do you think organic peroxides are so unstable? Some of them seem to detonate if you so much as even look at them the wrong way.

SunriseSunset - 31-10-2015 at 17:17

Ok.. That clears up a ton.. Thank you!! btw

UC235 - 31-10-2015 at 17:33

I'm pretty sure that there is no radical mechanism at play here. Sulfuric acid is a commonly used catalyst (and would have no effect on a radical mechanism). Probably the carboxylic acid is protonated, attacked at the carbonyl by peroxide's oxygen lone pair, followed by proton transfer and elimination of water. The entire thing is reversible and mechanistically looks a lot like Fischer Esterification but with hydrogen peroxide in place of an alcohol.

Darkstar - 31-10-2015 at 20:47

Quote: Originally posted by UC235  
I'm pretty sure that there is no radical mechanism at play here. Sulfuric acid is a commonly used catalyst (and would have no effect on a radical mechanism). Probably the carboxylic acid is protonated, attacked at the carbonyl by peroxide's oxygen lone pair, followed by proton transfer and elimination of water. The entire thing is reversible and mechanistically looks a lot like Fischer Esterification but with hydrogen peroxide in place of an alcohol.


The Fischer esterification-esque route was the second mechanism I proposed. I do tend to agree that protonation followed by nucleophilic attack will likely dominate in acidic media; however, I'm not entirely sure I'd completely discount the possibility of at least some kind of competing radical mechanism. And it's just as reversible via a radical mechanism as it is a reverse Fischer esterification. (peroxy acid disassociates back into a carboxyl radical and a hydroxyl radical. the carboxyl radical grabs a proton from water to give back the carboxylic acid and create a second hydroxyl radical, and then two hydroxyl radicals recombine to give back H2O2)

SunriseSunset - 1-11-2015 at 04:36

Could this b about right?

fische2.GIF - 2kB

Darkstar - 1-11-2015 at 07:21

Quote: Originally posted by SunriseSunset  
Could this b about right?


Yes, that very likely could be the mechanism.

SunriseSunset - 1-11-2015 at 10:11

It's interesting though, with a strong enough carboxylic acid such as Formic acid, you don't need an acid catalyst. So formic acid would be the source for Hydronium to kick this reaction off.