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I'm reading Bard's Electrochemical Methods. He cites the oxidation of p-aminophenol in acid solution. I haven't taken organic chemistry in a while and don't understand why oxidation takes place in acid. The pKa's of the amine and alcohol are 5 and 10, respectively. If you have a solution, say pH 1, wouldn't both moieties remain protonated and nothing happen? Thank you.

oxidation of p-aminophenol

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    $\begingroup$ The reaction (1.5.15) is not spontaneous. It only happens if an oxidant is in contact with the p-aminophenol $\endgroup$
    – Maurice
    Aug 20 at 12:13
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    $\begingroup$ Non-negligible amounts of double protonation will almost certainly require a superacidic medium, like pure triflic acid or carborane superacids in non-protic solvents. You need to check the pKb values of anilines and phenols (or equivalently the pKa values of aniliniums and phenyloxoniums). The pKa of phenol and aniline are irrelevant. Note that your reported "aniline pKa" is actually the anilinium pKa. $\endgroup$
    – Nicolau Saker Neto
    Aug 20 at 12:53
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    $\begingroup$ There is quite a difference in p$K_a$ between an (alkyl) amine, and a phenyl amine. Even more so the difference between an (alkyl) alcohol, and a phenol (a reference), to an extent phenol is a weak acid willing to offer protons and is labelled with a GHS05/corrosive sign (e.g. to skin and eye) while e.g., ethanol does not get it. This is why I suggest to call the starting material in 1.5.15 a phenol (when describing its chemistry around the OH group). $\endgroup$
    – Buttonwood
    Aug 20 at 15:34

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You can read about the quinhydrone electrode here: https://en.wikipedia.org/wiki/Quinhydrone_electrode

This electrode reaction is similar to the oxidation of Para-aminophenol and is measurably reversible over a pH range of about 2-10. It would be reasonable that the aminophenol coupled with the acid requirements of the oxidant would have an optimum pH [range] for the redox reaction.

Standard electrode potentials are when all activities are at unit levels. Some reactions can actually be measured under those conditions but sometimes the standard potentials must be calculated from other thermodynamic data. [Imagine trying to measure the standard potential of sodium metal immersed in one normal acid!] Research is needed to determine what is going on here.

About " and nothing happen.", Acid-base reactions in water involving proton transfers are very fast and frequent; something is always happening [sometimes catalysis is helpful].

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You are right in your assumption that in a solution with an acidic pH nothing would happen as is. But the example is considering a redox reaction.

If a suitable oxidant oxidises the p-aminophenol first, the resulting positively charged p-aminophenol will have drastically different pKa values because of the positive charge. The pKa might be low enough, that the deprotonation can even happen in acidic solution.

Alternatively, the reaction might also proceed via a concerted mechanism, where both a proton and an electron are transfered at the same time. In that case, pKa values don't apply anymore.

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