Lewis definition Knowledge-Based Study

Study notes

1. The Need for a Broader Definition

   Earlier models defined acid-base behavior in terms of specific particles:

   • Arrhenius definition: Limited to aqueous systems where acids produce $H^+$ ions and bases produce $OH^-$ ions.
   • Brønsted–Lowry definition: Broadened acids and bases to proton ($H^+$) donors and proton acceptors, allowing gas-phase and non-aqueous reactions.

   However, many chemical reactions show characteristic acid-base properties—such as neutralization, complexation, and indicator color changes—without any proton transfer taking place. In 1923, G. N. Lewis proposed a comprehensive definition based entirely on electron pairs.

   Electron-Pair Transfer

   The Lewis model shifts focus from the movement of protons to the interaction of electron pairs:

   • Lewis acid: A chemical species that accepts an electron pair to form a new bond (an electron-pair acceptor).
   • Lewis base: A chemical species that donates an electron pair to form a new bond (an electron-pair donor).

   Because every Brønsted–Lowry acid transfers an $H^+$ ion (which accepts an electron pair) and every Brønsted–Lowry base supplies a lone pair to bond with that proton, all Brønsted–Lowry bases are also Lewis bases. However, the Lewis definition encompasses many substances that lack hydrogen altogether.

   Structural Characteristics

   To identify whether a molecule or ion can act as a Lewis acid or base, look closely at its valence electron distribution:

   Lewis Bases (Electron-Pair Donors)

   • Molecules with lone pairs on electronegative atoms, such as $NH_3$, $H_2O$, and $CO$.
   • Anions possessing unshared electron pairs, such as $F^-$, $Cl^-$, and $OH^-$.

   Lewis Acids (Electron-Pair Acceptors)

   • Cations with accessible, empty orbitals, such as $H^+$, $Cu^{2+}$, and $Fe^{3+}$.
   • Molecules with incomplete octets on the central atom, such as boron and aluminum compounds ($BF_3$, $AlCl_3$).
   • Molecules with polar multiple bonds where the central atom can accept electrons upon bond reorganization, such as $CO_2$ and $SO_3$.

   Formation of an Adduct

   When a Lewis base donates a pair of electrons to a Lewis acid, both shared electrons in the resulting covalent bond originate from the base. This type of bond is called a coordinate covalent bond (or dative bond).

   The product of a Lewis acid-base reaction is commonly called a Lewis adduct or complex ion.

   Example: Reaction of Ammonia with Boron Trifluoride

   Consider the reaction between ammonia ($NH_3$) and boron trifluoride ($BF_3$):

   $$NH_3 + BF_3 ightarrow H_3N\text{—}BF_3$$

   • In $BF_3$, the boron atom has only six valence electrons and possesses an empty $2p$ orbital, making it a classic Lewis acid.
   • In $NH_3$, the nitrogen atom has an unshared lone pair of electrons, making it a Lewis base.
   • Nitrogen shares its lone pair with boron's empty orbital, forming a stable coordinate covalent bond and completing boron's octet.

   Key Takeaways

   • A Lewis acid is an electron-pair acceptor (electrophile).
   • A Lewis base is an electron-pair donor (nucleophile).
   • The reaction between them forms a coordinate covalent bond, yielding a Lewis adduct.
   • Metal cations acting as centers in coordination complexes (e.g., $[Cu(H_2O)_6]^{2+}$) are Lewis acids, while the surrounding ligands ($H_2O$) are Lewis bases.

   Quick Practice

   In the formation of the hydronium ion: $H^+ + H_2O ightarrow H_3O^+$

   The hydrogen ion ($H^+$) has an empty valence shell and accepts an electron pair, acting as the Lewis acid. The water molecule ($H_2O$) donates one of its lone pairs on oxygen, acting as the Lewis base.