It has been my understanding that when nitric, perchloric, or sulfuric acids act as oxidizing agents, they must exist as whole $\ce{HNO_3}$ or $\ce{HClO_4}$ or $\ce{H_2SO_4}$ molecules that are still protonated. For example, putting copper in dilute sulfuric acid does nothing, but using concentrated sulfuric acid does cause a redox reaction. This agrees with the fact that only concentrated forms of these acids can oxidize; only concentrated forms have undissociated/whole molecules. A dilute solution would contain mostly hydronium and the conjugate base of the strong acid, thus not be able to oxidize. However, the "General Properties" section of this Wikipedia article flat out contradict that. It says,

However, even dilute nitric acid can oxidize copper to $\ce{Cu^{2+}}$ ions, with the nitrate ions acting as the effective oxidant. (Emphasis: nitrate ions).

Then below it says that concentration is a factor for sulfuric acid. Is something amiss here?

UPDATE: I now embrace what user @matt_black is saying, and by now I'm not even surprised that something like this could happen. Chemical guidelines are just guidelines. However, I'm now even more curious as to this departure from the trend. If he or someone else would like to give an explanation of the mechanism of reactions involving oxidation (of copper, for instance) by $\ce{HNO_3}$ and $\ce{H_2SO_4}$ that accounts for this anomaly, that would be great.

  • $\begingroup$ If only the protonated forms are oxydants why all the perchlorate salts are strong oxidants/potential explosives? This chemistry does not add up $\endgroup$
    – Greg
    Commented May 13, 2016 at 20:54
  • $\begingroup$ @Greg It explains some things, such as how sulfuric acid exhibits its oxidizing powers only when fairly concentrated, and perchloric acid beyond 70%. $\endgroup$
    – Yunfei Ma
    Commented May 13, 2016 at 23:37
  • $\begingroup$ A strong acid is more reactive when concentrated than when the same amount of that acid is diluted in water. The acid is strong precisely because dissociation in water makes it give up free energy. $\endgroup$ Commented May 14, 2016 at 0:37

1 Answer 1


It is chemistry that is being weird not Wikipedia

The trouble with simple explanations of oxidising acids is that there is no simple explanation that covers all the cases. The specific reactions leading to oxidation are complicated and have a variety of different mechanisms.

So copper in sulphuric acid can be oxidised by a reaction that involves SO2 production. That has a completely different mechanism to the reaction with nitric acid and a different concentration dependence.

And, given the variety of reactions, the need to have protonated molecules of the acid is not a good generalisation: it depends on the specific reaction that is happening.

The oxidising reactions depend on the substrate being oxidised as well as the acid. Which mechanism will dominate isn't easy to determine without trying the reaction. So don't trust simple generalisations.

  • $\begingroup$ Which is why we sometimes describe something as acting as a Lewis acid - it is not the "proton donating" characteristic or the hydronium ion that is the mechanism of reaction. $\endgroup$
    – Stian
    Commented Aug 25, 2019 at 9:59

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