Biological Amines and the Henderson-Hasselbalch Equation

We saw in Section 20.3 that the extent of dis.sociation of a carboxylic acid MA in an aqueous solution buffered to a given pH can be calculated with the Hender.son-Hasscllialch equation. Furthermore, we concluded that at the physiological 

What about amine bases In what form do they exist at the physiological pH inside cells—as the amine (A = RNH2) or as the ammonium ion (HA = RNH3+) Let s take a 0.0010 M solution of methylamine at pH = 7.3, for example. -According to Table 24.1, the pK of mcthylammonium ion is 10.64, so from the Hcnderson-Hasselbalch equation, we liave 

Substituted arylamines can be either more basic or less basic than aniline, depending on the substituent. FJectron-donating substituents, such as -CH3, -NH2, and -OCH3, which increase the reactivity of an aromatic ring toward electrophilic substitution (Section 16.4), also increase the basicity of the corresponding arylamine. Electron-withdrawing substituents, such as -Cl, -NO2, and -CN, which decrease ring reactivity toward electrophilic substitution, also decrease arylamine basicity. Table 24.2 considers only p-substituted anilines, but similar trends are observed for ortho and meta derivatives. 

Without looking at Table 24.2, rank the following compounds in order of ascending basicity. 

Solving the two simultaneous equations gives [RNH3+] = 0.0010 M and [RNH2] = 5 X 10 M. In other words, at a physiological pH of 7.3, essentially 100% of the methylamine in a 0.0010 M solution exists in its protonated form as methylammonium ion. The same is tme of other amine bases, so we always write cellular amines in their protonated form and amino acids in their ammonium carboxylate form to reflect their stmctures at physiological pH. 

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Biological amines and the Henderson-Hasselbalch equation (Section 24.5). 

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