3
$\begingroup$

Among the amino acids, histidine and lysine, which is more basic? And why? I tried protonating and checking electronic effects but couldn't reach a satisfactory conclusion.

$\endgroup$
5
$\begingroup$

As their name imply, each amino acid has at least one amine and one acid functional group, but in most amino acids, the basicity of the amine is offset by the carboxylic acid group at physiological $\mathrm{pH}$, and hence considered neutral.

Amines are basic because the nitrogen has an unshared electron pair that can accept an $\ce{H+}$ more readily than that water does. They will even, to some extent, take an $\ce{H+}$ from water, leaving a hydroxide ($\ce{OH-}$) ion. There are three amino acids that have basic side chains at neutral $\mathrm{pH}$, which are named arginine (Arg), lysine (Lys), and histidine (His).

Arginine is the most basic among them because it contains guanidine side group, $\ce{-(CH2)4NHC(=NH)NH2}$, which is basic.

Lysine has two amine groups, which makes it overall basic because of the second isolated amine group $(\ce{-(CH2)4NH2})$.

Histidine, on the other hand, contains imidazole group, which is also basic.

The $\mathrm{p}K_\mathrm{a}$'s of all three side chains are high enough that they tend to bind protons, gaining a positive charge in the process (See below chart for $\mathrm{p}K_\mathrm{a}$ values). The $\mathrm{p}K_\mathrm{a}$ values are in order of

$$\ce{-(CH2)4NHC(=NH)NH2} > \ce{-(CH2)4NH2} > \ce{-(CH2)-(\text{4-imidazole})}$$

(see the chart), thus, lysine is considered to be more basic than histidine in physiological conditions. Please also note that, in the following chart: $\mathrm{p}K_\mathrm{a1} = \alpha$-carboxyl group, $\mathrm{p}K_\mathrm{a2} = \alpha$-amine group, $\mathrm{p}K_\mathrm{a3} =$ side chain group, and $\mathrm{pI} =$ the isoelectronic point.

$$ \begin{array}{lccccc} \hline \textrm{Amino Acid} & \mathrm{p}K_\mathrm{a1} & \mathrm{p}K_\mathrm{a2} & \mathrm{p}K_\mathrm{a3} & \mathrm{pI} & \textrm{acid/base} \\ \hline \textrm{Aspartic acid} & 1.88 & 9.60 & 3.65 & 2.77 & \textrm{acidic} \\ \textrm{Glutamic acid} & 2.19 & 9.67 & 4.25 & 3.22 & \textrm{acidic} \\ \textrm{Arginine} & 2.17 & 9. 04 & 12.48 & 10.76 & \textrm{basic} \\ \textrm{Lysine} & 2.18 & 8.95 & 10.53 & 9.74 & \textrm{basic} \\ \textrm{Histidine} & 1.88 & 9.17 & 6.00 & 7.59 & \textrm{basic} \\ \hline \end{array} $$

I thank University of Calgary, Canada and University of Wisconsin, Madison, WI for relevant data.

$\endgroup$
  • 2
    $\begingroup$ I took the liberty to improve this great answer's formatting and code semantics. Please note that \ce{...} is used for chemical formulas only, to put math operators in upright mode \mathrm{...} should be used instead. I also cleared-up the table a little according to ACS style recommendations, see also Small Guide to Making Nice Tables (PDF), which is pretty much a TL;DR version of those. $\endgroup$ – andselisk Apr 11 '19 at 21:03
  • 1
    $\begingroup$ @andselisk: Great edition! Thank you. You learn something everyday! :-) $\endgroup$ – Mathew Mahindaratne Apr 11 '19 at 21:06
  • $\begingroup$ Great explanation! I was wondering if there is any way to theoretically(by electronic effects) predict the basic strength without looking at the K values. $\endgroup$ – user226375 Apr 12 '19 at 4:52
  • $\begingroup$ @Mathew Mahindaratne , can you reply to my above question please? $\endgroup$ – user226375 Apr 21 '19 at 6:48

Your Answer

By clicking “Post Your Answer”, you agree to our terms of service, privacy policy and cookie policy

Not the answer you're looking for? Browse other questions tagged or ask your own question.