Peroxyacetic acid, $\ce{CH3C(O)OOH}$ has its anion form when a proton is detached, like $\ce{CH3C(O)OO-}$.

I think it can have two resonance forms like I drew (even though they have several charge separations, they're at least possible), but several websites address that it cannot have any resonance structures.

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  • $\begingroup$ 1st one would be OK if not for lack of positive formal charge. In 2nd the same issue gets more pronounced... $\endgroup$
    – Mithoron
    Commented Apr 21, 2023 at 18:15
  • $\begingroup$ I don't think first structure ever exists. $\endgroup$ Commented Apr 21, 2023 at 18:20
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    $\begingroup$ @mat I was thinking about first "alternate" one. You mean the anion itself? While maybe not terribly important in solution, afaict there are salts available. $\endgroup$
    – Mithoron
    Commented Apr 21, 2023 at 21:05
  • $\begingroup$ There is resonance in CO, so there should be resonance in RC(O)R. If that description is to be applied. It's just a mathematical model after all. $\endgroup$ Commented May 15, 2023 at 21:51

2 Answers 2


When you are considering resonance structures, you have to maintain certain aspects of the molecule; you have to maintain molecular formula/connectivity, number of electrons, and net charge. You maintain the molecular formula and number of electrons throughout your proposed structures, but what you fail to account for is net charge. In the true structure, there is a net charge of -1 resulting from the rightmost oxygen. In your structures, one example has a net charge of -2 while the other has -3. Here is a good source with the rules for resonance structures.

With that information, you can dismiss the resonance structures as invalid from the outset. This is because a resonance structure is the result of molecules "spreading" their electrons to achieve a more stable state; net charge will not get bigger. When you make the net charge bigger, as you have done, you make the molecule less stable, as there are now more lone pairs and charge is less evenly distributed. Furthermore, carbon will not pull an electron away from oxygen as you have done in your bottom drawing, at least in normal conditions; oxygen has a much higher electronegativity than carbon and will hold onto its electrons too tightly for carbon to remove them like that.

For these reasons, paracetic acid's conjugate does not show resonance because there is no resonance structure that follows the rules of resonance; the potential resonance structures you have considered are destabilizing, which is the opposite of the effect of resonance (resonance will always make a molecule's structure more stable).

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    $\begingroup$ When you are applying the concept of resonance, you have to maintain every aspect of the molecular. Resonance is a mathematical description resulting from an incomplete model. As such the stabilising effect is purely hypothetical. Unfortunately I have to disagree with this answer. To me there is no doubt that there is pi-delocalisation in this molecule. $\endgroup$ Commented May 15, 2023 at 21:47
  • $\begingroup$ What do you mean when you say "maintain every aspect of the molecular?" As well, are you saying that these proposed structures would exist in the molecule? Or are you only saying that to apply the concept of resonance as I have done is incorrect, and that the molecule will have delocalized electrons even if resonance cannot explain that delocalization? $\endgroup$ Commented May 22, 2023 at 15:18
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    $\begingroup$ Oh sorry, there is a typo, it should be: When you are applying the concept of resonance, you have to maintain every aspect of the molecule. This is due to you writing "certain aspects", which gives the impression that there are some aspects that differ between resonance structures. Resonance is a concept stemming from Valence Bond Theory. It captures the delocalisation of electrons within this theory; something which is inherent in Molecular Orbital Theory. I recommend the following: chemistry.stackexchange.com/q/51632/4945 $\endgroup$ Commented May 22, 2023 at 18:46
  • $\begingroup$ Ah yes I see haha. That wasn't my intention when I wrote "certain" but I see what you mean. Cheers. $\endgroup$ Commented Jul 1, 2023 at 4:04

Before we suggested any resonance structures for the conjugate base of peracetic acid (PAA), we have to prove it physically exists. The $\mathrm{p}K_\mathrm{a}$ of peracetic acid is found to be 8.3 (Ref.1). According to Ref.1, PAA is spontaneously decomposed to products given in equation $(1)$ in the $\mathrm{pH}$ range of 5 to 10: $$\ce{2CH3C(O)O-O-H -> 2CH3C(O)OH + O2} \tag1$$

The authors have shown that the maximum rate have occurred at $\mathrm{p}K_\mathrm{a}$ of PAA, and the mechanism of decomposition is gone according to following figure using labelled $\ce{CH3C(O)O^{18}-O^{18}-H}$.

Mechanism for decomposition of PAA

Accordingly, nucleophilic peroxy $\ce{O-}$ would attack the electrophilic carbonyl carbon of acetic part of another molecule to give the product as shown. The mechanism was confirmed by finding 87% of decomposed oxygen has been in $\ce{O^{18}-O^{18}}$ formula. However, a different group (Ref.2) has proved that there is no peroxy $\ce{O-}$ exists during the decomposition during the $\mathrm{pH} \le 7.5$:

[...] in the preparation of PAA, the $\ce{H+}$ concentration in the system is usually more than $\pu{0.1 mol L−1}$. It indicates that PAA is present as molecular form and no peraetic anion is found in the reaction solution.

Thus, the reaction mechanism of PAA spontaneous decomposition in an acid system can be predicted as two-step reaction. Step 1 is as follows:

Step 1 of mechanism

The protonated form $\mathrm{F_1}$ can be attacked by another unprotonated molecule in step 2 to give the active intermediate as follows:

Step 2 of mechanism

Thus, it is very unlikely that significant amount of peracetyl anion exist in these conditions before it self-decompose below its $\mathrm{p}K_\mathrm{a}$. That means it doesn't have a chance to show or form any resonance activity!


  1. E. Koubek, M. L. Haggett, C. J. Battaglia, Khairat M. Ibne-Rasa, H. Y. Pyun, and J. O. Edwards, "Kinetics and Mechanism of the Spontaneous Decompositions of Some Peroxoacids, Hydrogen Peroxide and t-Butyl Hydroperoxide," J. Am. Chem. Soc. 1963, 85(15), 2263–2268 (DOI: https://doi.org/10.1021/ja00898a016).
  2. Xuebing Zhao, Keke Cheng, Junbin Hao, and Dehua Liu, "Preparation of peracetic acid from hydrogen peroxide, part II: Kinetics for spontaneous decomposition of peracetic acid in the liquid phase," Journal of Molecular Catalysis A: Chemical 2008, 284(1-2), 58–68 (DOI: https://doi.org/10.1016/j.molcata.2008.01.003).
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    $\begingroup$ This is all great information, but I disagree with the final takeaway message. Just because something is under certain conditions reactive, it doesn't mean that this something isn't showing a stabilizing effect in the first place. I would be very surprised if there wasn't any pi-delocalisation in paa. $\endgroup$ Commented May 15, 2023 at 21:41

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