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As per PubChem: $pK_1 = 2.18$; $pK_2 = 8.95$; $pK_3 = 10.53$ at 38 °C, indicating that the side chain amino group is deprotonated at a higher pH, indicating it is more basic. I tried to justify this by analysing electronic effects in the conjugate acid:

Attempt I) The carboxylic acid group, at a higher pH would exist as $\ce{-COO}^{-}$ and exert a +I effect on nearby atoms. Then the alpha-amino group should be more stable in the protonated form and be more basic.

Attempt II) The carboxylate anion and $\ce{-NH_3}^{+}$ group may form hydrogen bonds. This again leads me to conclude that the alpha-amino group should be more basic.

But the experimental data contradict this. Perhaps the carboxylate anion H-bonds with (and stabilises) the side-chain amino group instead, but this seems unlikely (9-membered ring?).

Could anyone provide a decent theoretical explanation of the $pKa$ values?

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  • $\begingroup$ chemistry.stackexchange.com/questions/63181/…. $\endgroup$
    – Gaurav Sai Maddipati
    Commented Jan 21 at 13:11
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    $\begingroup$ Do you have a source for COO- being a +I group? $\endgroup$
    – Gaurav Sai Maddipati
    Commented Jan 21 at 13:12
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    $\begingroup$ @GauravSaiMaddipati Solomons/Fryhle/Snyder adapted by MS Chouhan, page 93 $\endgroup$
    – Sid
    Commented Jan 21 at 14:56
  • $\begingroup$ We must have different editions, my page 93(2016) is about Lewis acids and bases, btw all the best for your jee $\endgroup$
    – Gaurav Sai Maddipati
    Commented Jan 21 at 16:45
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    $\begingroup$ In aqueous solution, the solvation of the COO- group ensures that it maintains an I- effect. This is clearly observed in a comparison of the ammonium group pKa on glycine (~9.8) vs methylamine (~ 10.7). This effect is much stronger for the a-amino group of lysine vs the e-amino group, whose pKa is close to what is typical for an alkyl amine $\endgroup$
    – Andrew
    Commented Jan 25 at 12:28

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