Lewis acids and bases

The most general definition of acids and bases, which encompasses the Arrhenius and Bronsted-Lowry definitions is due to our old friend, Lewis and his dot structures. A Lewis acid is defined to be any species that accepts lone pair electrons. A Lewis base is any species that donates lone pair electrons. Thus, ${\rm H}^+$ is a Lewis acid, since it can accept a lone pair, while ${\rm OH}^-$ and NH$_3$ are Lewis bases, both of which donate a lone pair:

$${\rm H}^++ \stackrel{..}{\stackrel{:{\rm O}:}{..}}{\rm H}^- \longrightarrow {\rm H}_2{\rm O}$$

Interestingly, however, is that species which have no hydrogen to donate (a la the Bronsted-Lowry scheme) can still be acids according to the lewis scheme. As an example, consider the molecule BF$_3$. If we determine Lewis structure of BF$_3$, we find that B is octet deficient and can accept a lone pair. Thus it can act as a Lewis acid. Thus, when reacting with ammonia, the reaction would look like:

Figure 1:

In fact octet deficient molecules are often strong Lewis acids because they can achieve an octet configuration by accepting a lone pair from a Lewis base. Compounds involving elements in periods lower then the second period can act as Lewis acids as well by expanding their valence shells. Thus, SnCl$_4$ acts as a Lewis acid according to the reaction:

$${\rm SnCl}_4(l) + 2{\rm Cl}^-(aq)\longrightarrow {\rm SnCl}_6^{2-}(aq)$$

Figure 2:

The central tin atom is surrounded by a valence shell of 12 electrons rather than 8.