Lewis Concept of Acids and Bases:
This concept was proposed by G.N. Lewis, in 1939. According to this concept, a base is defined as a substance which can furnish a pair of electrons to form a coordinate bond whereas an acid is a substance which can accept a pair of electrons. The acid is also known as electron acceptor or electrophile while the base is electron donor or nucleophile.
A simple example of an acid-base is the reaction of a proton with hydroxyl ion.
Some other examples are:
H3N: + BF3 = H3N → BF3
H+ +: NH3 = [H ← NH3]+
BF3 + [F]- = [F → BF3]+
Lewis concept is more general than the Bronsted Lowry concept.
According to Lewis concept, the following species can act as Lewis acids.
(i) Molecules in which the central atom has incomplete octet: All compounds having central atom with less than 8 electrons are Lewis acids, e.g., BF3, BC13, A1C13, MgCl2. BeCL. etc.
(ii) Simple cations: All cations are expected to act as
Lewis acids since they are deficient in electrons. However, cations such as Na+, K+, Ca2+, etc., have a very little tendency to accept electrons, while the cations like H+, Ag+, etc., have greater tendency to accept electrons and, therefore, act asLewis acids.
(iii) Molecules in which the central atom has empty
d-orbitals: The central atom of the halides such as SiX4,
GeX4, TiCl4, SnX4, PX3, PF5, SF4, SeF4, TeCl4, etc., have vacant d-orbitals. These can, therefore, accept an electron pair and act as Lewis acids.
(iv) Molecules having a multiple bond between atoms of dissimilar electronegativity: Typical examples of molecules falling in this class of Lewis acids are C02, S02 and S03. Under the influence of attacking Lewis base, one -electron pair will be shifted towards the more negative atom.
Lewis acid Lewis base
The following species can act as Lewis bases.
(i) Neutral species having at least one lone pair of electrons: For example, ammonia, amines, alcohols, etc., act as Lewis bases because they contain a pair of electrons
(ii) Negatively charged species or anions: For example, chloride, cyanide, hydroxide ions, etc., act as Lewis bases.
CN-, CI-, OH-
It may be noted that all Bronsted bases are also Lewis bases but all Bronsted acids are not Lewis acids.
Since the strength of the Lewis acids and bases is found to depend on the type of reaction, it is not possible to arrange them in any order of their relative strength.
The choice of which definition of acids and bases one wishes to use in a particular instance depends largely on the sort of chemistry that is studied. But Arrhenius concept is perfectly satisfactory and simplest for dealing with reactions in aqueous solutions. It explains satisfactorily the strength of acids and bases in aqueous solutions, neutralisation, salt hydrolysis, etc.