The Lewis theory of acids and bases gives us a wider view compared to older definitions like Arrhenius and Bronsted-Lowry. According to Lewis, an acid is something that accepts electron pairs, while a base donates electron pairs. This change helps us see acid-base reactions in a new way.

Key Differences:

1. Electron Trading: Instead of just looking at how protons (like $H^+$) move, Lewis theory focuses on how electron pairs are exchanged. This helps explain some relationships that don’t fit the older definitions.

2. Broader Definitions: Many substances that act like acids or bases don’t follow the classic definitions. For example, aluminum chloride ($AlCl_3$) works like a Lewis acid because it accepts electron pairs, even though it doesn’t give away protons.

Examples:

• Lewis Acid: $BF_3$ can accept an electron pair because it has an empty part called a p-orbital.
• Lewis Base: Ammonia ($NH_3$) has a pair of electrons it can give away.

Illustrations:

Think about what happens when $NH_3$ and $BF_3$ react. In this case, $NH_3$ donates its electron pair to $BF_3$. This creates a special bond called a coordinate covalent bond. This example shows how the Lewis theory covers a wider range of acid-base behaviors beyond just exchanging protons.

In short, the Lewis theory helps us understand acids and bases better by focusing on how electrons interact. This makes it an important idea in both organic and inorganic chemistry.