9.2 The Lewis Concept of Acids and Bases

The Arrhenius and Brønsted-Lowry theories of acids and bases, while fundamental, have limitations in explaining the acid-base behavior of certain substances. For instance, compounds like $\mathrm{S}{\mathrm{O}}_2$, $\mathrm{C}{\mathrm{O}}_2$, $\mathrm{CaO}$, and $\mathrm{B}{\mathrm{F}}_3$ act as acids or bases despite not having the ability to donate or accept protons. To address these limitations, G.N. Lewis proposed a more general acid-base theory in 1923, focusing on electron pair transfer.

Definitions

Lewis Acid: A substance that can accept a pair of electrons to form a coordinate covalent bond. Lewis acids are often referred to as electron pair acceptors.

Lewis Base: A substance that can donate a pair of electrons to form a coordinate covalent bond. Lewis bases are often referred to as electron pair donors.

In a Lewis acid-base reaction, a coordinate covalent bond (also known as a dative bond) is formed. This is a covalent bond in which both shared electrons originate from one atom (the Lewis base). The resulting product is often called an Adduct.

Illustrative Example: Boron Trifluoride and Ammonia

Consider the reaction between boron trifluoride ($\mathrm{B}{\mathrm{F}}_3$) and ammonia ($\mathrm{N}{\mathrm{H}}_3$):

$\mathrm{B}{\mathrm{F}}_3+\mathrm{N}{\mathrm{H}}_3\to {\mathrm{F}}_3\mathrm{B}\leftarrow \mathrm{N}{\mathrm{H}}_3$

The bond formed between boron and nitrogen is a coordinate covalent bond, as the electron pair is donated by nitrogen.

Let's analyze this reaction based on Lewis's definitions:

1. Which species is donating an electron pair? $\mathrm{N}{\mathrm{H}}_3$ (specifically, the nitrogen atom in ammonia)

2. Which species is accepting an electron pair? $\mathrm{B}{\mathrm{F}}_3$ (specifically, the boron atom in boron trifluoride)

3. Which species is a Lewis acid? $\mathrm{B}{\mathrm{F}}_3$ (electron pair acceptor)

4. Which species is a Lewis base? $\mathrm{N}{\mathrm{H}}_3$ (electron pair donor)

Gas-Phase Neutralization Reactions

The Lewis theory can effectively explain gas-phase neutralization reactions, which are not typically covered by Arrhenius or Brønsted-Lowry theories. For example, the reaction between hydrogen chloride ($\mathrm{HCl}$) gas and ammonia ($\mathrm{N}{\mathrm{H}}_3$) gas forms ammonium chloride ($\mathrm{N}{\mathrm{H}}_4\mathrm{Cl}$):

$\mathrm{HCl}(\mathrm{g})+\mathrm{N}{\mathrm{H}}_3(\mathrm{g})\to \mathrm{N}{\mathrm{H}}_4\mathrm{Cl}(\mathrm{s})$

In this reaction, the nitrogen atom in ammonia donates an electron pair to the hydrogen atom in $\mathrm{HCl}$ (which then leaves as $\mathrm{C}{\mathrm{l}}^{-}$). Therefore, $\mathrm{N}{\mathrm{H}}_3$ is the Lewis base (electron pair donor), and $\mathrm{HCl}$ (specifically, the $\mathrm{H}$ atom acting as ${\mathrm{H}}^{+}$) is the Lewis acid (electron pair acceptor).

Characteristics of Lewis Acids and Bases

Lewis Acids:

• The central atom of the Lewis acid typically has a deficiency of an electron pair (an incomplete octet).
• It must have an empty orbital to accommodate an unshared electron pair.
• Examples include: $\mathrm{B}{\mathrm{F}}_3$, $\mathrm{AlC}{\mathrm{l}}_3$, ${\mathrm{H}}^{+}$, metal ions ($\mathrm{F}{\mathrm{e}}^{3+}$, $\mathrm{C}{\mathrm{u}}^{2+}$), and molecules with multiple bonds where atoms can expand their octet ($\mathrm{C}{\mathrm{O}}_2$, $\mathrm{S}{\mathrm{O}}_2$).

Lewis Bases:

• The central atom of a Lewis base usually has a complete octet and possesses one or more unshared electron pairs (lone pairs).
• Hence, the base has the ability to donate an unshared electron pair.
• Examples include: $\mathrm{N}{\mathrm{H}}_3$, ${\mathrm{H}}_2\mathrm{O}$, $\mathrm{O}{\mathrm{H}}^{-}$, $\mathrm{C}{\mathrm{l}}^{-}$, and organic molecules with lone pairs on heteroatoms (e.g., ethers, amines).

Worked Example: Classifying Substances as Lewis Acids or Lewis Bases

Problem: Identify the Lewis acid and Lewis base in the following reaction:

$\mathrm{B}{\mathrm{F}}_3+\mathrm{N}{\mathrm{H}}_3\to {\mathrm{F}}_3\mathrm{B}\leftarrow \mathrm{N}{\mathrm{H}}_3$

Strategy:

1. Draw Electronic Structures: Sketch the Lewis structures of both species.
2. Identify Electron Pair Donor (Lewis Base): Look for the species that has a lone pair of electrons or a negative charge. Such a species typically has a complete octet and can donate an electron pair.
3. Identify Electron Pair Acceptor (Lewis Acid): Look for the species that can accommodate an electron pair (e.g., has an incomplete octet or an empty orbital).

Solution:

1. Electronic Structures:

   • For $\mathrm{N}{\mathrm{H}}_3$: Nitrogen has one lone pair of electrons and forms three single bonds with hydrogen atoms.
   • For $\mathrm{B}{\mathrm{F}}_3$: Boron forms three single bonds with fluorine atoms and has only six valence electrons (an incomplete octet).

2. Identify Lewis Base: $\mathrm{N}{\mathrm{H}}_3$ has a lone pair on the nitrogen atom. This lone pair can be donated. Therefore, $\mathrm{N}{\mathrm{H}}_3$ is an electron pair donor — it is the Lewis base.

3. Identify Lewis Acid: Boron in $\mathrm{B}{\mathrm{F}}_3$ has only six valence electrons (three bonding pairs), meaning it has an incomplete octet and an empty orbital available to accept an electron pair. Therefore, $\mathrm{B}{\mathrm{F}}_3$ is the Lewis acid.

4. Product: The coordinate covalent bond forms between N and B, producing the adduct ${\mathrm{F}}_3\mathrm{B}\leftarrow \mathrm{N}{\mathrm{H}}_3$, where the arrow indicates the electron pair was donated by nitrogen.