A chemical bond is an interaction that holds atoms or groups of atoms together in a stable chemical structure. Bonds organize matter into molecules and extended solids, including ionic crystals and metals. They involve the distribution of electrons and interactions between electrons and atomic nuclei. In chemistry, bonding provides a framework for explaining composition, structure, and reactivity. The term encompasses several kinds of interaction rather than a single universal mechanism. (goldbook.iupac.org)
Physical basis
Bond formation is associated with a lowering of energy relative to an appropriate separated reference state. In a simple diatomic molecule, the energy varies with internuclear distance: attraction favors association at intermediate separations, while repulsion dominates at very short distances. The minimum of this energy curve defines an equilibrium separation. Atoms in a bond are not motionless; they vibrate around this configuration. (openstax.org)
A detailed description requires quantum mechanics. Electrons are represented by quantum states rather than classical paths connecting nuclei. Bonding reflects electron–nucleus attraction, electron–electron and nucleus–nucleus repulsion, and quantum constraints on electronic states. The familiar picture of atoms sharing electrons is therefore a useful model, not a literal account of particles being exchanged back and forth. (openstax.org)
Principal bonding types
Covalent bonding involves shared electron density between atoms. The simplest example is hydrogen, H₂, whose two electrons occupy a bonding state. Conventional structural formulas represent single, double, and triple bonds with one, two, and three lines. Covalent bonding occurs both in discrete molecules and in extended networks, such as diamond, where each carbon atom is connected to neighboring atoms throughout the solid. (openstax.org)
Ionic bonding emphasizes electrostatic attraction between oppositely charged ions. Solid sodium chloride is an extended arrangement of sodium cations and chloride anions, not a collection of isolated NaCl molecules. Electron transfer is a useful way to describe the formation of ions, but the attraction between those ions is the bonding interaction. Ionic and covalent descriptions are limiting models: actual bonds can possess both ionic and covalent character. (goldbook.iupac.org)
Metallic bonding involves electrons delocalized across many atomic sites. This collective electronic structure helps explain why metals commonly conduct heat and electricity and can deform without immediately breaking apart. Unlike a localized two-atom bond, metallic bonding is naturally described over an extended solid. (openstax.org)
A coordinate covalent bond is described as forming when both electrons of a shared pair originate from one donor. This description is particularly useful in coordination chemistry, where a ligand donates electron density to a central atom or ion. It identifies the origin of the electron pair rather than establishing a fundamentally separate force. (goldbook.iupac.org)
Models of electronic structure
A Lewis structure represents valence electrons as dots and shared pairs as lines. The octet rule helps account for many compounds of main-group elements, but it is not universal: electron-deficient species, odd-electron molecules, and other exceptions occur. Resonance uses several contributing structures when a single Lewis formula cannot adequately represent the electronic distribution. These contributors are not separate molecules alternating in time. (openstax.org)
Valence bond theory describes bonding through overlapping atomic orbitals and electron pairing. Orbital hybridization constructs combinations of orbitals useful for representing directional bonds. Molecular orbital theory instead describes electrons in orbitals extending across a molecule. Its bonding and antibonding states account for stability and electronic properties, including the unpaired electrons responsible for oxygen’s magnetic behavior. These approaches provide different representations of the same underlying quantum system. (openstax.org)
Bond properties and molecular shape
Bond length is the distance between bonded nuclei. Bond order expresses the degree of bonding between atomic sites; in a simple molecular-orbital treatment, it equals half the difference between bonding and antibonding electron populations. For comparable bonds between the same elements, higher bond order generally corresponds to shorter, stronger bonds, although the chemical environment also matters. (openstax.org)
Electronegativity describes an atom’s tendency to attract electrons in a bond. Unequal electron sharing produces bond polarity, but molecular polarity also depends on three-dimensional shape. For example, the polar bonds of linear carbon dioxide cancel one another’s dipole contributions, whereas the bent arrangement of water produces a net molecular dipole. Electron-pair repulsion models predict many molecular geometries by considering bonding regions and lone pairs around a central atom. (openstax.org)
Bond energies and weaker interactions
Bond dissociation energy measures the energy required to break a specified bond under defined conditions. Breaking bonds requires energy; forming bonds releases energy relative to the corresponding separated fragments. The net energy change of a chemical reaction therefore depends on both processes. Average bond energies offer approximate reaction-enthalpy estimates, while ionic solids are commonly characterized using lattice energies for the collective crystal structure. (openstax.org)
Hydrogen bonds and van der Waals interactions are often distinguished from the primary bonds defining molecular connectivity. Nevertheless, a broad definition of chemical bonding can include weak associations. Hydrogen bonding involves attraction between a hydrogen attached to a relatively electronegative atom and an appropriate acceptor; it may occur between molecules or within one molecule. Such interactions influence molecular association and material structure without necessarily changing conventional covalent connectivity. (goldbook.iupac.org)