Covalent bond

A covalent bond is a fundamental type of chemical bond characterized by the sharing of electron pairs between atoms, primarily between nonmetal atoms, to achieve a stable electronic configuration.

Nature and Formation

A covalent bond forms when two atoms, each with incomplete valence shells, approach each other closely enough that their atomic orbitals overlap. This overlap allows a pair of electrons, with opposite spins, to be shared between the nuclei of both atoms. The shared electron pair is attracted to the positive charges of both nuclei, creating a strong, directional bond that holds the atoms together in a molecule. The energy minimum reached at a specific internuclear distance represents the bond length, and the energy released upon bond formation is the bond energy. This model is central to valence bond theory.

Types of Covalent Bonds

Covalent bonds are categorized based on the number of shared electron pairs and the electronegativity difference between the bonded atoms.

• Single, Double, and Triple Bonds: A single bond involves one shared pair of electrons (e.g., H–H in H₂). A double bond involves two shared pairs (e.g., O=O in O₂), and a triple bond involves three shared pairs (e.g., N≡N in N₂). Increasing bond order generally correlates with shorter bond lengths and higher bond energies.
• Polar and Nonpolar Covalent Bonds: In a nonpolar covalent bond, electrons are shared equally between identical atoms (e.g., Cl–Cl). A polar covalent bond occurs between atoms with different electronegativities (e.g., H–Cl), resulting in an unequal sharing of electrons. This creates a dipole moment, with a partial negative charge (δ−) on the more electronegative atom and a partial positive charge (δ+) on the less electronegative atom.

Bonding Theories and Representation

Two primary quantum mechanical models explain covalent bonding. Valence Bond (VB) Theory emphasizes the pairing of electrons in overlapping atomic orbitals, including hybrid orbitals (e.g., sp³ in methane) to account for molecular geometry. Molecular Orbital (MO) Theory describes electrons as delocalized over the entire molecule, occupying molecular orbitals formed from the linear combination of atomic orbitals. This theory effectively explains the properties of molecules like O₂. Covalent structures are commonly depicted using Lewis structures, which show bonding electron pairs as lines and non-bonding lone pairs as dots.

Properties of Covalently Bonded Substances

Substances with extensive covalent bonding networks exhibit distinct physical properties. Molecular covalent compounds (discrete molecules like H₂O or CO₂) typically have low melting and boiling points, are often gases or liquids at room temperature, and are poor conductors of electricity. In contrast, covalent network solids (e.g., diamond, silicon dioxide) consist of a giant lattice of atoms connected by covalent bonds, resulting in extremely high melting points and hardness, though they are also electrical insulators. Graphite is an exceptional network solid with delocalized electrons, making it a conductor.

Comparison with Other Bond Types

Covalent bonding differs significantly from ionic and metallic bonding. Ionic bonds involve the complete transfer of electrons from a metal to a nonmetal, forming charged ions held together by electrostatic forces, producing brittle, high-melting-point solids that conduct electricity when molten or dissolved. Metallic bonds involve a "sea" of delocalized electrons surrounding positive metal ions, granting properties like malleability and high electrical conductivity. The continuum from nonpolar covalent to ionic bonding is described by the concept of electronegativity and percent ionic character.