Table of Contents

• Objective

• Introduction

• Lewis Bases

  • Lone Pairs

  • Pi Bonds

  • Sigma Bonds

• Lewis Acids

  • Empty Orbitals

  • Partial Positive Charges

    • Double Bond

    • Single Bond

      • Halogens as Lewis Acids

Objective

Welcome to our guide on Lewis acids and bases. By the end of this guide, you’ll understand the broader Lewis definition of acids and bases, how to identify them, and the roles they play in organic chemistry reactions. This guide is perfect for chemistry students looking to deepen their understanding or prepare for assessments.

Introduction

The Lewis definition of acids and bases is broader than the Brønsted-Lowry definition. While the Brønsted-Lowry theory describes acidity and basicity in terms of proton transfer, the Lewis theory focuses on electron transfer.

• A Lewis base donates a pair of electrons

• A Lewis acid accepts a pair of electrons

• The donated electron pair is shared between the acid and base in a covalent bond

All Brønsted-Lowry acids and bases can be classified as Lewis acids and bases.

For example, consider the reaction between HCl and NH₃ above.

HCl acts as Brønsted-Lowry acid because it donates a proton to NH₃. HCl also acts as a Lewis acid because it accepts a pair of electrons from NH₃.

The Lewis definition expands the concept of acids and bases to include reagents that would otherwise not be classified as such under the Brønsted-Lowry definition.

For example, consider the reaction between H₂O and BF₃.

According to the Brønsted-Lowry theory, BF₃ is not an acid because it has no protons and cannot donate a proton. However, according to the Lewis theory, it is an acid because it accepts a pair of electrons from the Lewis base, H₂O.

Lewis Bases

A Lewis base, also known as a nucleophile, donates a pair of electrons to a Lewis acid. Therefore, Lewis bases have electron-rich centers, which can be lone pairs, pi bonds, or sigma bonds. Lewis bases can be negatively charged or neutral.

Lone Pairs

The most common type of Lewis base has a lone pair of electrons. Examples include ammonia (NH₃), water (H₂O), alcohols (ROH), and amines (RNH₂).

Other atoms later in the periodic table that carry lone pairs, such as phosphine (PH₃), thiols (RSH), and sulfides (RSR'), also make good Lewis bases.

Anions with lone pairs, such as Cl^- , Br^- , OH^- , and CN^- , are often strong Lewis bases.

Pi Bonds

Molecules containing carbon-carbon pi (π) bonds, such as alkenes and alkynes, can act as Lewis bases. The pi bonds are rich in electron density and can be donated to a Lewis acid.

Sigma Bonds

Sigma (σ) bonds in certain Lewis bases can donate electrons, especially when the bond is between a carbon atom and a metal. Examples include Grignard reagents and alkylithiums.

Grignard reagents (RMgX) contain a carbon-magnesium (C-Mg) bond, while alkylithiums (RLi) contain a carbon-lithium (C-Li) bond. In both cases, the carbon is more electronegative than the metal, making it electron-rich and capable of donating electrons.

Reducing agents like NaBH₄ and LiAlH₄ also fall under this category.

Lewis Acids

A Lewis acid, also known as an electrophile, accepts a pair of electrons from a Lewis base. Therefore, Lewis acids have electron-poor centers, which can include atoms with empty orbitals and atoms that bear partial positive charges due to induction or polarization. Lewis acids can be neutral or positively charged.

Empty Orbitals

The simplest Lewis acid is the hydrogen cation (H⁺), commonly referred to as a proton. Because H⁺ has no electrons and an empty 1s orbital, it is highly reactive and can readily accept a pair of electrons from a Lewis base.

Carbocations, which are positively charged carbon atoms with an empty p orbital, also act as Lewis acids. The empty p orbital allows them to accept electron pairs, making them highly reactive intermediates in organic reactions.

Many metal cations, such as Mg²⁺ (magnesium), Ca²⁺ (calcium), and Al³⁺ (aluminum), act as Lewis acids. These cations have vacant orbitals that can accept electron pairs from Lewis bases.

Additionally, compounds of group 3A elements, like boron trifluoride (BF₃) and aluminum chloride (AlCl₃), are also Lewis acids. These compounds possess empty orbitals that can accept electron pairs from Lewis bases

Partial Positive Charges

Molecules that make good Lewis acids or electrophiles might have a double bond or single bond to an electronegative atom such as O, N, Cl, or Br.

Double Bond 

The most important molecules with a double bond to an electronegative atom include carbonyl compounds (C=O).

The oxygen atom, being more electronegative, withdraws electron density from the carbon, making it electron-poor. As a result, carbonyl compounds act as Lewis acids because the electron-poor carbon can accept electrons from a Lewis base.

Single Bond 

Molecules with a single bond to electronegative atoms can also make good Lewis acids. Examples include alkyl halides (R-X) where X is a halogen (Cl, Br, I).

The electronegative halogen withdraws electron density from the carbon, making it electron-poor. As a result, alkyl halides act as Lewis acids because the electron-poor carbon can accept electrons from a Lewis base.

Halogens as Lewis Acids

Some molecules that contain no dipoles can still function as Lewis acids. Examples are halogens like I₂, Br₂, and Cl₂. Bromine (Br₂), for example, is strongly electrophilic and reacts rapidly with alkenes.

Why is Br₂ a Lewis acid? The answer lies in polarization. Bromine is a large, diatomic molecule with a significant number of electrons. This makes it highly polarizable, meaning its electron cloud can be easily distorted.

When bromine approaches a region of high electron density, such as a nucleophile, the electron cloud shifts, creating a temporary dipole.

This shift makes one bromine atom slightly positive (δ⁺) and the other slightly negative (δ⁻), allowing the positive end to accept electrons from a Lewis base.