In 1932, G.N. Lewis introduced a broader concept of acids and bases based on the transfer of electron pairs, overcoming the limitations of the Arrhenius and Bronsted-Lowry theories. According to Lewis:
🔹 Lewis Acid
A Lewis acid is a substance that can accept a pair of electrons.
🔹 Lewis Base
A Lewis base is a substance that can donate a pair of electrons.
🔹 Adduct Formation
When a Lewis acid and base react, they form a coordinate covalent bond. Metal cations (e.g., H⁺) act as Lewis acids, while anions (e.g., non-metals) act as Lewis bases. The result is a single compound called an adduct.
🧬 Example: NH3 + BF3
The reaction between ammonia (NH3) and boron trifluoride (BF3) demonstrates Lewis acid-base interaction:
NH3 acts as a Lewis base by donating its lone electron pair, while BF3 acts as a Lewis acid by accepting it due to its incomplete octet. The bond forms through the empty 2pz orbital of boron. This is a typical Lewis acid-base reaction where no proton transfer occurs.
📌 Key Point:
The Lewis theory is more general than the Arrhenius or Bronsted-Lowry concepts, as it includes reactions where proton transfer is not involved.
⚠️ Limitations of Lewis Concept
- Cannot compare acid/base strengths effectively.
- Fails to explain behavior of acids not forming coordinate bonds.
- Formation of coordinate bonds is slow, unlike rapid acid-base reactions.
- Interferes with coordination chemistry principles.
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