Basicity of Amines

Amines are basic in character specifically because of the lone pair of electrons carried on their nitrogen atom, and this basicity can be described from two related theoretical viewpoints. In terms of Lewis acid-base theory, an amine is a base because its nitrogen lone pair can be SHARED with an electron-deficient species -- for example, trimethylamine donates its nitrogen lone pair to the electron-deficient boron atom of boron trifluoride, Me3N + BF3 gives the Lewis adduct Me3N-BF3(+/-) (a dative, coordinate covalent N-B bond forms). In terms of Bronsted-Lowry acid-base theory, an amine is instead a base because its nitrogen lone pair can ACCEPT a proton, written as the equilibrium: amine + H2O is in equilibrium with the protonated ammonium-type conjugate acid + OH(-) -- equation (13.1) in the source. Basic strength is expressed QUANTITATIVELY through the equilibrium (or dissociation) constant Kb, or more conveniently its negative logarithm, pKb: the FURTHER this equilibrium shifts toward the right (i.e. the more the amine is protonated at equilibrium), the STRONGER a base the amine is, corresponding to a LARGER Kb value and, correspondingly, a SMALLER pKb value (and vice versa for a weaker base). Table 13.4 tabulates measured pKb values, in aqueous …