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Lewis Acid and Base Theory

09th Apr 2022 @ 7 min read

Inorganic Chemistry

In the early 1900s, Gilbert N. Lewis, one of the well-known American chemists, introduced the concept of Lewis acids and bases. Today, his theory has led us to a better understanding of what acids and bases are and has been able to get explanations for many chemical reactions, which remained unclear. In this article, we will study what are Lewis acids and bases with examples in depth.

Lewis Acid

A Lewis acid is a chemical species that can accept a pair of electrons, namely from a Lewis base. Since Lewis acids accept electrons, they must have a positive affinity towards electrons. In other words, they attract electrons, and thus, they are electrophiles.

The positive affinity towards electrons suggests that a Lewis acid must have a positive charge on it or a vacant orbital that could readily accept electrons. In more technical terms, Lewis acids contain Lowest Unoccupied Molecular Orbitals (LUMO).

Examples of Lewis Acids

We have a broad diversity of Lewis acids. Some are simply limited to a single element, while some are complex, composed of multiple groups. Below are a few examples of Lewis acids.


The cation is any species with a net positive charge on it. All cations can be classified as Lewis acids since they all can accept negatively charged electrons to neutralize the charge on them. Some examples of cations are Na+, Mg2+, Fe3+, Cu2+.

Complexed cations, like carbocations (R3C+), are also Lewis acids.

Protonated species

A proton (H+) is a Lewis acid. Other cations that are derived after protonation are also classified as Lewis acids. For example, ammonium (NH4+) ions are produced from the protonation of ammonia.

the formation of ammonium (NH4+) ions

An ammonium ion can readily accept a pair of electrons from anions, like OH−, to form the ammonium salt.

Other example of Lewis acids are phosphonium (PH4+), oxonium (H3O+), chloronium (H2Cl+), Boronium (BH4+)

Species with incomplete octet

Chemical species that have incomplete octet have a natural tendency to accept electrons. Popular examples are boron trifluoride (BF3) and aluminum trichloride (AlCl3).

In boron trifluoride, the central boron atom is surrounded by three fluorine atoms. It is sp2 hybridized and has a trigonal planar structure. Boron shares its three electrons with three fluorine atoms. Each fluorine atom contributes an electron to the bonding. Thus, the total number of electrons available to Boron in the outermost shell are six (three from boron plus three from fluorine), making the octet of Boron incomplete.

structure of boron trifluoride

One p-orbital of Boron remains unhybridized. This orbital acts as the electron-deficient center in BF3 and can accept a pair of electrons from an electron donor.

We can arrive at a similar conclusion for aluminum trichloride.

Pentahalides of group 15

The pentahalides of phosphorus, arsenic, and antimony show characteristics of Lewis acids. A notable example of the Lewis acid in this category is phosphorus pentachloride (PCl5).

Phosphorus pentachloride is sp3d hybridized. Each of the five hybridized orbitals is bonded with a chlorine atom. PCl5 has a trigonal bipyramidal structure; three chlorine atoms are in a plane with the central phosphorus atom, and the other two chlorine atoms are in the axial position. The unhybridized vacant d orbitals of phosphorus and five electronegative chlorine atoms give PCl5 acidic characteristics. In many reactions, PCl5 readily accepts electrons, making it Lewis acid.

Species with electronegative groups

Chemical species in which a central atom is surrounded by relatively strong electronegative atoms or groups also act as Lewis acids.

Examples are carbon dioxide (CO2) and sulfur dioxide (SO2).

In carbon dioxide, the central carbon atom is bonded with two oxygen atoms, which are more electronegative than the central element. So, the bonding electrons are shifted more toward oxygen. This creates a distortion of charge neutrality around carbon. The central carbon atom experiences electron deficiency and can accept electrons, making it a Lewis acid.

Lewis Base

The Lewis base is the counterpart of the Lewis acid. It is a chemical species that can donate a pair of electrons, namely to a Lewis acid. Lewis bases have a positive affinity toward positive charge and are categorized as nucleophiles.

Since a Lewis base can readily donate a pair of electrons, it must have a lone pair of electrons in its outermost orbital called HOMO (Highest Occupied Molecular Orbital). The HOMO of the Lewis base interacts with the LUMO of the Lewis acid to form the molecular orbital.

HOMO of the Lewis base interacts with the LUMO of the Lewis acid to form the molecular orbital

Examples of Lewis Bases

Diverse chemical species belong to Lewis bases, but the most common are anions. Below mentions some of them.


Simple anions carrying negatively charged electrons are all Lewis bases. Examples include H−, F−, Cl−, O2−, OH−, CN−, CH3COO−.

Species with lone pair of electrons

Many chemical species have a lone pair of electrons in the outer orbitals. These electrons can be readily denoted to an electron acceptor. A popular example is ammonia.

Ammonia has a trigonal pyramidal structure as shown below.

structure of Ammonia

The nitrogen atom is bonded to three hydrogen atoms by sigma bonds. The hybridization of ammonia is sp3. Of four hybridized orbitals, three are utilized in bonding with hydrogen And the remaining orbital is used for holding a lone pair of electrons. This pair of electrons can be donated to electron-deficient species.

Other examples are pyridine, water, amines, and ether.

Π systems

Certain π systems rich in electron conjugation can donate electrons. These systems attain extra stability from conjugation. Examples are ethylene, ethyne, benzene.


Usually, when a Lewis acid reacts with a Lewis base, the resulting product is the direct addition of the two and is called a Lewis adduct. The bond between the Lewis acid and the Lewis base is the coordinate covalent bond, aka dative bond. In a coordinate covalent bond, one reactant is solely responsible for the contribution of both electrons. Some reactions are discussed below.

Formation of ammonium ion

The ammonia is a Lewis base. It has a lone pair of electrons in one of its orbitals, which can be readily donated. In the presence of the proton-rich environment, ammonium ions are formed. The reaction is as follows:

formation of ammonium ion reaction

Formation of hydronium ion

Similar to ammonium ions, oxygen in water molecules donates its lone pair to H+ under protonic medium to hydronium ions.

Formation of hydronium ion

Similar to ammonium ions, oxygen in water molecules donates its lone pair to H+ under protonic medium to hydronium ions.

Formation of hydronium ion

Formation of complex polyatomic adducts

Many a time, the Lewis acid reacts with the Lewis base to form a complex structure, composed of multiple groups of the Lewis base neighboring the central Lewis acid. One such example is aluminum hexahydrate ions.

Formation of polyatomic adducts
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