Question

Explain why all Bronsted acids are Lewis acids, but not all Lewis acids are Bronsted acids. i need examples

Answer 1

A Lewis acid is hungry for electrons, a Bronsted-Lowry acid releases an H+ in a proton-exchange reaction. The reason the BL acid is always a Lewis acid is that the only way a molecule will release an H+ is because the atom to which the H is bonded "wants" the electrons in that bond so much it readily kicks the H+ out without them.

For example, consider the classic BL acid, HCl reacting with water:

HCl(aq) + H2O(l) -> H3O+(aq) + Cl-(aq)

The Cl is very electronegative, and is willing to release the H+ to the H2O because it (the Cl atom) gets to keep the electrons in the H-Cl bond, resulting in a Cl- anion. The Cl atom here is a Lewis acid, because it is "hungry" for the electrons in the H-Cl bond.

However, it is possible for an atom to be electron hungry without involving the release of an H+. The classic example are metal cations, e.g. Al+3. These have empty valence orbitals which are strongly attracted to available valence electrons (as would be supplied by a Lewis base). Hence, Al+3 can also react with water and result in the release of H+:

Al+3(aq) + 10 H2O(l) -> Al(H2O)2(OH)4+3(aq) + 4 H3O+(aq)

Which looks in the overall sense very much like the HCl reaction. But the Al+3 did not release any H+, of course. What it did was form a dative bond with a Lews base (water), and then the resulting complex released an H+.

Lewis acidity enlarges the idea of acidity from the BL concept, which in turn enlarged it from the Arrhenius concept. Arrhenius said an acid was something that released an H+ when it dissolved. BL enlarged this, focussing on the fact that generally the release of H+ required a chemical reaction with some other compound, the BL base, which was eager to make a bond with an H+. That allows you to recognize things as acids that don't release H+ on dissolution, but can do so when reacting with a strong BL base, for example H2O. Lewis enlarged this still further, by focussing on the behaviour of the electrons, not the proton, and saying that the key fact that produces a proton transfer reaction is the considerably greater attraction of one compound (the Lewis acid) for valence electrons that can be supplied by another (the Lewis base). That allows you to recognize things as acids that are very likely to form dative bonds, which, willy nilly, can (but need not) result in the release of an H+ by the resulting complex.