2.11 Acids and Bases: The Lewis Definition

Lewis Acids and Bases

The Lewis definition of acids and bases is broader than the Brønsted-Lowry definition you've already seen. Instead of focusing on proton transfer, it focuses on electron pair transfer. This wider lens lets you explain reactions that don't involve protons at all, which comes up constantly in organic mechanisms.

Lewis Acids and Bases

A Lewis acid is an electron pair acceptor. It has an empty or electron-poor orbital that can accept a pair of electrons. Common examples include $\ce{BF3}$, $\ce{AlCl3}$, and $\ce{H+}$. Notice that $\ce{BF3}$ has only six electrons around boron, so it's hungry for two more to complete an octet.

A Lewis base is an electron pair donor. It has a lone pair available to share with an electron-deficient species. Common examples include $\ce{NH3}$, $\ce{H2O}$, and $\ce{OH-}$.

The driving force behind many Lewis acid-base reactions is the octet rule: atoms tend to react in ways that give them a complete valence shell. When a Lewis base donates its lone pair into the empty orbital of a Lewis acid, both species often end up with more stable electron configurations.

Curved Arrow Notation

Curved arrows are how you track electron pair movement in these reactions. Two rules to remember:

1. The arrow starts at the electron source (the lone pair on the Lewis base).
2. The arrow points toward the electron acceptor (the Lewis acid).

For example, in the reaction $\ce{NH3 + BF3 -> H3N-BF3}$:

• Nitrogen on $\ce{NH3}$ has a lone pair (this is the Lewis base).
• Boron on $\ce{BF3}$ has an empty p orbital (this is the Lewis acid).
• The curved arrow starts at the lone pair on nitrogen and points to boron, showing that nitrogen donates its electrons into boron's empty orbital.

The result is a new coordinate covalent bond (a bond where both electrons came from the same atom). In the product, boron now has a full octet and carries a formal negative charge, while nitrogen carries a formal positive charge.

Oxygen and Nitrogen as Lewis Bases

In organic chemistry, the most common Lewis bases are atoms with available lone pairs, especially oxygen and nitrogen.

Oxygen has two lone pairs in most functional groups, making it a reliable Lewis base:

• In ethanol ($\ce{CH3CH2OH}$), the oxygen can donate a lone pair to a Lewis acid like $\ce{BF3}$ or $\ce{H+}$.
• In dimethyl ether ($\ce{CH3OCH3}$), oxygen similarly uses a lone pair to form an adduct with a Lewis acid.
• Carbonyl oxygens ($\ce{C=O}$) also act as Lewis bases, which is central to many carbonyl addition reactions you'll see later.

Nitrogen typically has one lone pair available for donation:

• Methylamine ($\ce{CH3NH2}$) donates its nitrogen lone pair to Lewis acids readily.
• Pyridine ($\ce{C5H5N}$) has a lone pair on nitrogen that sits in the plane of the ring (not part of the aromatic system), making it available for donation.

Nitrogen is generally a stronger Lewis base than oxygen in comparable molecules because nitrogen is less electronegative and holds its lone pair less tightly, making it more willing to share.

Coordination Chemistry and Transition Metal Complexes

Lewis acid-base theory extends beyond organic molecules to metal chemistry. When a transition metal ion (the Lewis acid) accepts electron pairs from surrounding molecules or ions called ligands (the Lewis bases), the result is a coordination complex.

You won't need deep detail on this for organic chemistry, but recognizing that metal-catalyzed reactions rely on Lewis acid-base interactions helps connect the concepts. For instance, $\ce{AlCl3}$ acting as a Lewis acid catalyst in Friedel-Crafts reactions is a direct application of this idea.