Covalent bond

Covalent bonding is a fundamental concept in chemistry describing the sharing of electrons between atoms. Through this sharing, atoms form electron pairs that hold them together, creating stable molecular structures. The balance between attractive and repulsive forces established when atoms share electrons results in a covalent bond, which typically allows atoms to achieve full valence shells resembling stable electronic configurations. Covalent bonding dominates organic chemistry and is essential for understanding molecular behaviour, structure, and reactivity.

Concept and Characteristics of Covalent Bonding

A covalent bond arises when two atoms share one or more pairs of electrons. These shared pairs, also called bonding pairs, enable the atoms to attain complete outer shells. Covalency is most pronounced between atoms of similar electronegativity because the electrons are attracted relatively equally to both nuclei. This feature means that covalent bonds can form between different elements, provided their electronegativities are comparable.
Electron sharing may occur between just two atoms or can extend over several atoms, producing delocalised electrons. Such delocalisation is a characteristic of aromatic systems, conjugated molecules, and certain bonding arrangements found in complex organic and inorganic substances. In simple diatomic molecules such as hydrogen, each atom contributes one electron, forming a shared pair that satisfies the duet rule for hydrogen and the general octet pattern seen in other elements.

Historical Development

The concept of covalent bonding has deep roots in early twentieth-century chemical theory. The term covalence was first introduced in 1919 by Irving Langmuir, who defined it as the number of electron pairs an atom shares with neighbouring atoms. Langmuir’s work built upon the ideas of Gilbert N. Lewis, who in 1916 proposed that atoms bond by sharing electron pairs. Lewis also introduced the widely used Lewis dot structures to depict valence electrons and bonding patterns.
Quantum mechanics later provided mathematical justification for covalent bonding. In 1927, Walter Heitler and Fritz London applied quantum mechanical principles to the hydrogen molecule, demonstrating that bond formation could be explained through overlap of atomic orbitals. Their work formed the basis of valence bond theory, which remains essential for interpreting how atoms combine to form molecules.

Types of Covalent Bonds

The directional characteristics of atomic orbitals lead to different types of covalent bonds:

• Sigma (σ) bonds: These are the strongest type of covalent bond, formed by head-on orbital overlap. A single bond between atoms is typically a sigma bond.
• Pi (π) bonds: Formed by lateral overlap of p or d orbitals, pi bonds are generally weaker than sigma bonds. They occur alongside sigma bonds in double and triple bonds. A double bond contains one sigma and one pi bond, while a triple bond contains one sigma and two pi bonds.

Electronegativity differences also influence covalent bonds. When atoms have equal electronegativity, the bond is non-polar, as seen in the hydrogen molecule. When electronegativity differs, a polar covalent bond emerges, as in hydrogen chloride. Molecular polarity, however, also depends on geometry, since dipoles may cancel in symmetrical shapes.

Covalent Structures

Covalent bonding produces various structural forms depending on how atoms are arranged:

• Individual molecules: These have strong internal bonds but weak intermolecular forces. They tend to be gases or volatile liquids, such as methane, carbon dioxide, or sulfur dioxide.
• Molecular structures: Examples include iodine or solid carbon dioxide, where weak intermolecular interactions give rise to low melting and boiling points.
• Macromolecular structures: Long chains of covalently bonded atoms occur in polymers such as polyethylene and nylon, as well as in biopolymers including proteins and starch.
• Giant covalent structures (network covalent bonding): These consist of extensive networks of covalent bonds in two or three dimensions. Substances like diamond, graphite, and quartz exhibit high melting points, brittleness, and high resistivity. Elements with high electronegativity and the capacity to form multiple electron-pair bonds commonly form these networks.

One-Electron and Three-Electron Bonds

In certain chemical species with an odd number of electrons, unconventional bonding arrangements appear:

• One-electron bonds: Found in species such as the dihydrogen cation, these bonds contain a single shared electron and typically possess about half the bond energy of a typical two-electron bond. Some exceptions exist, such as the stronger one-electron bond in dilithium, which is attributed to hybridisation and inner-shell effects.
• Three-electron bonds: These occur in species like the helium dimer cation and nitric oxide. They consist of two electrons forming a bonding interaction and a third electron occupying an antibonding orbital. The resulting bond order is typically one-half. Dioxygen can be described using two three-electron bonds and one typical two-electron bond, which explains its paramagnetism and observed bond order of 2.

Molecules featuring odd-electron bonds are usually highly reactive and are most stable when formed between atoms of similar electronegativity.

Resonance and Bond Order

Many molecules cannot be satisfactorily represented by a single Lewis structure. Instead, they are described using resonance, where multiple structures contribute to a composite picture that reflects the real electron distribution. Bond order becomes an average value across these contributing structures. For instance, in the nitrate ion, nitrogen forms bonds with three equivalent oxygen atoms. One structure may depict a double bond to one oxygen and single bonds to the others, but resonance requires that each nitrogen–oxygen bond have an average bond order of between single and double.
Resonance is essential for understanding bonding in delocalised systems and provides insight into molecular stability, length, and reactivity patterns.

Aromaticity

Aromaticity describes the enhanced stability seen in planar ring structures that follow Hückel’s rule, which requires that the number of electrons in the delocalised pi system fit the formula 4n + 2, where n is an integer. Compounds such as benzene exemplify aromatic stabilisation through uniform electron distribution across the ring, contributing to distinctive chemical behaviour.
In aromatic systems, delocalisation of electrons leads to equalisation of bond lengths and remarkable thermodynamic stability, which cannot be accounted for by simple single or double bond representations. Aromaticity thereby illustrates the importance of delocalised covalent bonding in organic chemistry.