LEWIS CONCEPT OF ACIDS AND BASES
G.N. Lewis, in 1932, defined acid-base based on the electron pair acceptor or donor, as Arrhenius and Bronsted-Lowry concept is only applicable to certain substances.
Acids:
An acid is a substance that can accept a pair of electrons.
Bases:
A base is a substance that can donate a pair of electrons.
Adduct formation:
A coordinate covalent bond is formed between Lewis acid and base. The cations (metals or H+ itself) acts as Lewis acids (accepts a lone pair). The anions (non-metals) etc act as Lewis bases (donate lone pair). The product of any Lewis acid-base reaction is a single species called adduct.
For example:
A reaction between ammonia and boron trifluoride is an example of Lewis acid and base.
In this example, NH3 has a lone pair on N-atom. So it is an electron-pair donor. NH3 is a Lewis base. Boron in BF3 has an incomplete octet. It has six electrons ( 3 electron pairs) so it needs an electron pair to complete its octet. Hence BF3 is an electron pair acceptor or Lewis acid. The empty 2pz orbital of a boron atom in Boron trifluoride accepts the electron pair from ammonia. Thus in this reaction ammonia is a Lewis base and boron trifluoride is a Lewis acid. Although no proton transfer is observed here.
Note:
The Lewis concept of acid and bases is more generalized than Arrhenius and Bronsted-Lowry's concept because it applies to such acid-base reactions which are not covered by the other two concepts.
Limitations:
- Lewis's concept is unable to compare the strength of different acids/bases because of the formation of coordinate covalent bonds.
- Acids do not form coordinate covalent bonds so their acidic character cannot be explained.
- The coordinate covalent bond formation is slow but acid-base reactions are fast. So it does not fit into the concept.
- Lewis's concept interferes with coordination chemistry.