In chemistry, a halogen bond (XB or HaB) occurs when there is evidence of a net attractive interaction between an electrophilic region associated with a halogen atom in a molecular entity and a nucleophilic region in another, or the same, molecular entity. Like a hydrogen bond, the result is not a formal chemical bond, but rather a strong electrostatic attraction. Mathematically, the interaction can be decomposed in two terms: one describing an electrostatic, orbital-mixing charge-transfer and another describing electron-cloud dispersion. Halogen bonds find application in supramolecular chemistry; drug design and biochemistry; crystal engineering and liquid crystals; and organic catalysis.

Halogen bonds occur when a halogen atom is electrostatically attracted to a partial negative charge. Necessarily, the atom must be covalently bonded in an antipodal σ-bond; the electron concentration associated with that bond leaves a positively charged "hole" on the other side. Although all halogens can theoretically participate in halogen bonds, the σ-hole shrinks if the electron cloud in question polarizes poorly or the halogen is so electronegative as to polarize the associated σ-bond. Consequently halogen-bond propensity follows the trend F < Cl < Br < I.

There is no clear distinction between halogen bonds and expanded octet partial bonds; what is superficially a halogen bond may well turn out to be a full bond in an unexpectedly relevant resonance structure.

A halogen bond is almost collinear with the halogen atom's other, conventional bond, but the geometry of the electron-charge donor may be much more complex.

Anions are usually better halogen-bond acceptors than neutral species: the more dissociated an ion pair is, the stronger the halogen bond formed with the anion.

A parallel relationship can easily be drawn between halogen bonding and hydrogen bonding. Both interactions revolve around an electron donor/electron acceptor relationship, between a halogen-like atom and an electron-dense one. But halogen bonding is both much stronger and more sensitive to direction than hydrogen bonding. A typical hydrogen bond has energy of formation 20 kJ/mol; known halogen bond energies range from 10–200 kJ/mol.

The σ-hole concept readily extends to pnictogen, chalcogen and aerogen bonds, corresponding to atoms of Groups 15, 16 and 18 (respectively).

In 1814, Jean-Jacques Colin discovered (to his surprise) that a mixture of dry gaseous ammonia and iodine formed a shiny, metallic-appearing liquid. Frederick Guthrie established the precise composition of the resulting I₂···NH₃ complex fifty years later, but the physical processes underlying the molecular interaction remained mysterious until the development of Robert S. Mulliken's theory of inner-sphere and outer-sphere interactions. In Mulliken's categorization, the intermolecular interactions associated with small partial charges affect only the "inner sphere" of an atom's electron distribution; the electron redistribution associated with Lewis adducts affects the "outer sphere" instead.