The given statement is incorrect. While there is an overlap between Bronsted-Lowry acids and Lewis acids, not all Bronsted-Lowry acids are Lewis acids. Bronsted-Lowry acids are defined as species that can donate a proton (H+), whereas Lewis acids are substances that can accept an electron pair.

While many Bronsted-Lowry acids, such as HCl or H2SO4, also meet the criteria of Lewis acids as they can accept an electron pair from a lone pair of electrons on a base, there are some Bronsted-Lowry acids that do not.

For example,

HSO4- ( hydrogen sulfate anion) acts only as a Bronsted-Lowry acid, donating a proton to form H2SO4, but it cannot accept an electron pair and therefore not behaves as a Lewis acid. Another example would be water (H2O) in acid dissociation equilibrium (Autoprotolysis of water)

H2O + H2O ⇌ H3O+ + OH-

In such cases where only proton transfer occurs, the term Bronsted-Lowry acids/bases is more appropriate and not all the Bronsted-Lowry acids will be a Lewis acid.