A covalent bond is a chemical bond in which atoms are held together through the sharing of electrons. More precisely, it involves relatively high electron density between atomic nuclei, arising at least partly from electron sharing and associated with attraction and a characteristic internuclear distance. Covalent bonding forms the framework of many molecules and also occurs in extended solids. The familiar shared-electron-pair picture is useful, but not every covalent interaction consists of one pair localized between two atoms. (goldbook.iupac.org)
Physical basis and theoretical descriptions
Covalent bonding is explained by quantum mechanics. Electrons are described by states rather than fixed trajectories, and bond formation depends on the energy of the combined system. In hydrogen, attraction between electrons and both nuclei stabilizes the molecule, while electron–electron and nucleus–nucleus repulsions oppose excessively close approach. The resulting energy minimum determines an equilibrium bond length; electron sharing alone does not guarantee a stable bond. (openstax.org)
Two complementary models describe this electronic structure. Valence bond theory emphasizes overlap between atomic orbitals and the pairing of electrons in localized bonding interactions. Orbital hybridization combines orbitals on an atom into directional orbitals useful for describing molecular geometry. These orbitals are mathematical constructions, not separate physical objects. (openstax.org)
Molecular orbital theory instead describes electrons in orbitals extending over a molecule. Bonding orbitals stabilize the system, whereas antibonding orbitals oppose bonding when occupied. This approach accounts for properties not captured by simple electron-pair diagrams, including the unpaired electrons and resulting paramagnetism of molecular oxygen. Both theories are models of the same underlying quantum system. (openstax.org)
Electron-pair notation and its limits
A Lewis structure represents valence electrons as dots and shared pairs as lines between atomic symbols. A single line denotes one shared pair; two or three lines denote double or triple bonds. Nonbonding pairs, called lone pairs, are shown separately. Thus H–H represents hydrogen, O=O represents oxygen, and N≡N represents nitrogen. (openstax.org)
The octet rule describes a common pattern in which main-group atoms have eight electrons in their valence-shell bookkeeping after bonding. Hydrogen instead commonly has two. The rule is not universal: electron-deficient compounds, odd-electron species, and structures conventionally drawn with expanded valence shells require exceptions. Lewis diagrams record electron allocation and connectivity rather than the complete electron distribution. (openstax.org)
Sometimes several Lewis structures describe the same atomic arrangement with different electron placements. Resonance represents this situation; the actual structure is not a molecule repeatedly switching between drawings. Delocalization can make equivalent bonds intermediate in character between conventional single and double bonds. Consequently, a localized pair between each bonded atom pair is not always an adequate description. (openstax.org)
Single and multiple bonds
In the usual orbital description, a sigma bond has electron density distributed around the internuclear axis. A pi bond arises from sideways orbital overlap, with a nodal plane containing that axis. Ordinary single bonds are σ bonds, double bonds contain one σ and one π component, and triple bonds contain one σ and two π components. (openstax.org)
Bond order expresses the net extent of bonding. In elementary Lewis notation, single, double, and triple bonds have orders of one, two, and three. In a simple molecular-orbital calculation, bond order equals half the difference between the numbers of bonding and antibonding electrons. Fractional values are possible. For comparable bonds between the same elements, greater bond order generally corresponds to shorter, stronger bonds, although molecular environment also matters. (openstax.org)
Polarity and ionic character
Electron sharing need not be equal. Electronegativity measures an atom’s tendency to attract bonding electron density. Identical atoms in an isolated homonuclear diatomic molecule share it symmetrically. Unlike atoms often produce polar covalent bonds: the more electronegative atom carries a partial negative charge, δ−, and its partner a partial positive charge, δ+. These are partial charges, not necessarily the integer charges used in an ionic bonding model. (openstax.org)
Covalent and ionic descriptions are limiting models, with many bonds showing characteristics of both. Electronegativity differences provide a useful guide but do not supply a universal sharp boundary. Bond polarity also differs from molecular polarity: the polar bonds in linear carbon dioxide cancel by symmetry, whereas the bent structure of water produces a net molecular dipole. (openstax.org)
Bond energies and bulk properties
The equilibrium bond length lies near a minimum on the relevant potential-energy curve. A bond dissociation energy measures the energy required to separate a specified bond into defined fragments; tabulated bond enthalpies commonly describe gas-phase homolytic cleavage. Average values vary from molecule-specific values because neighboring atoms affect bond strength. Breaking bonds requires energy, while forming bonds releases it; their balance helps estimate the enthalpy change of a chemical reaction. (openstax.org)
Strong covalent bonds do not necessarily imply a high melting point. Molecular solids contain discrete molecules held together by interactions such as hydrogen bonding and van der Waals forces. Melting can disrupt these interactions without breaking the molecules’ covalent frameworks. Network solids instead have extended covalent structures: diamond has a three-dimensional network, while graphite has strongly bonded sheets with weaker attractions between them. Their contrasting hardness and electrical behavior show why bulk properties depend on bonding topology and electron delocalization, not simply on the label “covalent.” (openstax.org)