In theory, all Brønsted–Lowry acids can function as Lewis acids, but not all Lewis acids are Brønsted–Lowry acids. The reason is grounded in their definitions. A Brønsted–Lowry acid is one that can donate a proton (H⁺), and since a proton is essentially a nucleus without electrons, it's eager to accept an electron pair, adhering to the Lewis definition. However, Lewis acids, like BF₃, don't always donate protons. They may accept an electron pair without any proton transfer, making them exclusively Lewis acids and not Brønsted–Lowry acids.

Lewis acids can play a pivotal role in catalysis by accepting electron pairs and thus increasing the reactivity of other species in the system. By doing so, they can facilitate certain reactions by stabilising intermediates or activating specific reactants. A common application is in the petrochemical industry where Lewis acid catalysts assist in cracking hydrocarbons or polymerising olefins. Their electron-accepting nature can make certain bonds in the reactant molecules more susceptible to breaking or facilitate the formation of new bonds, thereby speeding up the reaction.

The primary difference stems from the nature of their interactions. Lewis bases interact by donating a pair of electrons, which can lead to the formation of coordinate bonds with Lewis acids. This donation isn't limited to protons; it can involve other species like metal cations or molecules. In contrast, Brønsted–Lowry bases specifically interact by accepting a proton from an acid. This proton acceptance is limited to reactions where a proton (H⁺) transfer occurs. For instance, OH⁻ is a Brønsted–Lowry base when it accepts a proton from an acid like HCl but acts as a Lewis base when it donates an electron pair to a metal ion like Al³⁺.

Some molecules possess regions that can accept electron pairs and other regions that can donate electron pairs, making them amphoteric in the Lewis sense. For instance, water (H₂O) can act as a Lewis base by donating a pair of its non-bonding electrons, like when it binds to metal ions. Conversely, it can also act as a Lewis acid in the presence of stronger Lewis bases like OH⁻. This dual nature arises from the molecular structure and the distribution of electrons, allowing certain molecules to interact in multiple ways depending on the reaction context.

The dative or coordinate bond holds unique importance in Lewis acid-base reactions. It represents the donation of an electron pair from the Lewis base to the Lewis acid. This type of bond forms when one reactant (the Lewis base) donates a pair of non-bonding electrons to another reactant (the Lewis acid). An illustrative example is the reaction between ammonia (NH₃) and boron trifluoride (BF₃). Ammonia, acting as a Lewis base, donates an electron pair to the boron atom in BF₃, which serves as a Lewis acid. The bond between nitrogen and boron in the resulting complex is a dative bond, which illustrates the Lewis acid-base reaction.