I know that anions such as $\ce{CH3COO-}, \ce{OH-}, \ce{CN-}$ are basic in nature, since they are conjugate bases of weak acids. Similarly, cations such as $\ce{NH4+}, \ce{H3O+}$ are acidic in nature , since they conjugate acids of weak bases.

However, we also have another case whereby, anions such as $\ce{HSO3-}, \ce{HS-}$ are amphoteric in nature. This is due to the fact that their conjugate acids are polyprotic acids. Thus, one can have the two reactions:

$\ce{HS- + H2O <=> S^2- + H3O+}$ and $\ce{HS- + H2O <=> H2S + OH-}$

Since, for this case, we got amphoteric anions from polyprotic acids, is there a possibility that we can do the same for polyprotic bases so as to get amphoteric cations? i.e. Can we have a conjugate acid of a polyprotic base, which can be regarded as an amphoteric cation? I haven't seen any such examples of amphoteric cations yet, hence I am asking this question.

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    $\begingroup$ Polyamines. Or basic aminoacids. $\endgroup$
    – Poutnik
    Jan 9 '20 at 14:16
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    $\begingroup$ $\ce{H3O+}$ and $\ce{H3S+}$ are technically amphoteric (as well as likely many others), though they're (always?) extremely weak bases. I wonder if there are more natural examples which are limited to few atoms. $\endgroup$ Jan 9 '20 at 14:27
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    $\begingroup$ Or CaOH+, as Ca(OH)2 has 2 dissociation constants. $\endgroup$
    – Poutnik
    Jan 9 '20 at 15:04
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    $\begingroup$ What about singly protonated DABCO? (ie 1,4-diazabicyclo[2.2.2]octane). It is cationic, and the neutral conjugate base has no additional acidic protons. With one proton, it is weakly basic (pKa of conjugate acid is ~3) and weakly acidic (pKa ~ 9). $\endgroup$
    – Andrew
    Jan 9 '20 at 15:27
  • $\begingroup$ Thing is, almost any cation can be protonated. Only the smallest ones may be to unstable to have minimum on PES. Another thing is that "amphoteric" tends to be used for stuff easily / commonly acting as acid and base, so what exactly you'd like to get here? $\endgroup$
    – Mithoron
    Jan 9 '20 at 18:15

Yes, for example lysine.

Lysine has two amino groups and one carboxyl group. When dissolved in aprotic solvents, all of these will be uncharged so the carboxyl group will be $\ce{-COOH}$ and the two amino groups $\ce{-NH2}$. When dissolved in water, these three groups have distinct $\mathrm pK_\mathrm a$ values of $2.15, 9.16$ and $10.67$. Depending on the acidity of the solution, the following (acid-sided) equilibriums occur:

Two deprotonations of lysine dication

When dissolved in water, the carboxyl group’s proton will immediately be removed and used to immediately protonate the ε-amino group (the far end). However, the α-amino group is still more basic than water and thus will be partially protonated; meaning that the right-hand equilibrium will be happening predominantly if lysine is dissolved in water. Dissolved lysine will be a mixture of uncharged lysine and lysine cation.

When an acid such as hydrochloric acid is added, at first the α-amino group is completely protonated and then the left-hand equilibrium occurs predominantly as more and more lysine cation is converted to lysine dication.

If a base is added (such as sodium hydroxide), the ε-amino group’s proton can also be removed yielding lysine anion; I have not shown that equilibrium in the scheme above.

Lysine cation thus fits your question as it is an amphoteric cation, the conjugate acid of a dibasic compound.


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