By Timothy Burns
Updated Mar 24, 2022
Covalent bonds form when two or more atoms share one or more pairs of electrons, allowing each atom to achieve a stable outer electron configuration. Think of this as a balanced stool: each leg (electron pair) is essential for stability.
The simplest covalent compounds are diatomic molecules, such as O₂, H₂, and Cl₂, which naturally exist as pairs of identical atoms. Their shared electrons satisfy the octet rule for each constituent atom.
A single covalent bond involves two atoms sharing one electron pair. Classic examples include hydrogen chloride (HCl) and methane (CH₄). In CH₄, each hydrogen atom shares one electron with the central carbon, giving carbon eight electrons and each hydrogen two, thereby fulfilling the octet rule.
When two atoms share two electron pairs, a double bond forms, which is markedly stronger than a single bond. The O₂ molecule features a double bond with a bond energy of approximately 498 kJ mol⁻¹, contributing to its high stability. This strength means that breaking the O=O bond—such as in electrolysis of water—requires substantial energy input.
Covalently bonded molecules often remain gases at room temperature because the forces between individual molecules are weak. Their strong intramolecular bonds leave no incentive for the molecules to interact closely, resulting in low melting points and persistent gaseous states.
Unlike ionic compounds, covalent substances do not dissociate into free ions when dissolved in water. Consequently, aqueous solutions of covalent molecules are typically non‑conductive, as the electrons remain bound within the molecules rather than moving freely.
