The strength of a Bronsted-Lowry acid is measured using the pH scale.
Lewis acids are electron pair acceptors. They have an unoccupied low-energy atomic or molecular orbitals.
The Lewis definition is more flexible than the traditional Bronsted-Lowry definition.
The Lewis definition includes compounds that are Bronsted-Lowry acids but also encompasses compounds that do not donate protons, but still exhibit acid behavior.
Lewis Acid Example 1
In the traditional acid/base neutralization reaction of H+ and OH-, H+ is a Lewis acid because it accepts an electron pair from the OH-.
Lewis Acid Example 2
Boron trifluoride acts as a Lewis acid when it accepts an electron pair from ammonia to form a co-ordinate covalent bond - i.e. a bond in which both bonding electrons come from the same atom:
The co-ordinate bond may be drawn as a double-headed arrow to represent the fact that both electrons come from one atom, in this case the nitrogen atom:
The reactions of Lewis acids with bases are interpreted with the HSAB (hard and soft acids and bases) concept.