Define Lewis acid and Lewis base with examples.

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Points to Remember:

  • Definition of Lewis acid and Lewis base.
  • Examples of Lewis acids and bases.
  • Distinction between Lewis and Brønsted-Lowry definitions.

Introduction:

The concept of acids and bases is fundamental to chemistry. While the Brønsted-Lowry theory defines acids as proton donors and bases as proton acceptors, the Lewis theory provides a broader definition, encompassing a wider range of chemical reactions. Gilbert N. Lewis proposed his theory in 1923, expanding the understanding of acid-base interactions beyond the limitations of the proton transfer model. This theory focuses on the donation and acceptance of electron pairs.

Body:

1. Definition of Lewis Acid:

A Lewis acid is defined as a species that can accept a pair of electrons to form a coordinate covalent bond. This means it has an empty orbital that can accommodate the electron pair. Lewis acids are often electron-deficient species, possessing a positive charge or an incomplete octet.

Examples of Lewis Acids:

  • Metal cations: Al³⁺, Fe³⁺, Zn²⁺. These ions have empty orbitals and readily accept electron pairs from Lewis bases.
  • Boron trifluoride (BF₃): Boron has only six electrons in its valence shell, making it electron-deficient and capable of accepting an electron pair.
  • Carbon dioxide (CO₂): The carbon atom in CO₂ can accept electron pairs from Lewis bases, forming carbonate ions.
  • Aluminum chloride (AlCl₃): Similar to BF₃, AlCl₃ is an electron-deficient compound that acts as a Lewis acid.

2. Definition of Lewis Base:

A Lewis base is defined as a species that can donate a pair of electrons to form a coordinate covalent bond. This means it possesses a lone pair of electrons that it can share with a Lewis acid.

Examples of Lewis Bases:

  • Ammonia (NH₃): The nitrogen atom in ammonia has a lone pair of electrons that it can donate.
  • Water (H₂O): The oxygen atom in water has two lone pairs of electrons available for donation.
  • Hydroxide ion (OH⁻): The oxygen atom in the hydroxide ion possesses a lone pair of electrons.
  • Chloride ion (Cl⁻): The chloride ion has four lone pairs of electrons and can act as a Lewis base.

3. Distinction between Lewis and Brønsted-Lowry Definitions:

It’s crucial to understand that all Brønsted-Lowry acids and bases are also Lewis acids and bases, but not vice-versa. The Lewis definition is more inclusive. For example, BF₃ is a Lewis acid but not a Brønsted-Lowry acid because it doesn’t donate a proton. Similarly, many molecules with lone pairs, such as NH₃, can act as Lewis bases but may not be Brønsted-Lowry bases in all reactions.

Conclusion:

The Lewis theory of acids and bases provides a more comprehensive understanding of acid-base reactions by focusing on electron pair donation and acceptance. While the Brønsted-Lowry theory is useful for many reactions involving proton transfer, the Lewis theory expands the scope to include a broader range of chemical species and reactions. Understanding both theories is essential for a complete grasp of acid-base chemistry. Further research into the applications of Lewis acid-base chemistry in various fields, such as catalysis and materials science, will continue to advance our understanding of chemical interactions and pave the way for innovative applications. This broader perspective promotes a more holistic understanding of chemical principles and their impact on various scientific disciplines.

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