Halogen bond

and trimethylamine. 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. == Donor characteristics ==

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. • Multi-electron donors such as ethers and amines prefer halogen bonds collinear with the lone pair and donor nucleus. • Pyridine derivatives tend to donate halogen bonds approximately coplanar with the ring, and the two C–N–X angles are about 120°. • Carbonyl, thiocarbonyl-, and selenocarbonyl groups, with a trigonal planar geometry around the Lewis donor atom, can accept one or two halogen bonds. 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. == Comparison to other bond-like forces ==

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 ; known halogen bond energies range from 10–200 kJ/mol. ==History==

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. Then, in 1954, Odd Hassel fruitfully applied the distinction to rationalize the X-ray diffraction patterns associated with a mixture of 1,4-dioxane and bromine. The patterns suggested that only 2.71 Å separated the dioxane oxygen atoms and bromine atoms, much closer than the sum (3.35 Å) of the atoms' van der Waals radii; and that the angle between the O−Br and Br−Br bond was about 180°. From these facts, Hassel concluded that halogen atoms are directly linked to electron pair donors in a direction with a bond direction that coincides with the axes of the orbitals of the lone pairs in the electron pair donor molecule. Dumas and coworkers first coined the term "halogen bond" in 1978, during their investigations into complexes of CCl4, CBr4, SiCl4, and SiBr4 with tetrahydrofuran, tetrahydropyran, pyridine, anisole, and di-n-butyl ether in organic solvents. However, it was not until the mid-1990s, that the nature and applications of the halogen bond began to be intensively studied. Through systematic and extensive microwave spectroscopy of gas-phase halogen bond adducts, Legon and coworkers drew attention to the similarities between halogen-bonding and better-known hydrogen-bonding interactions. In 2007, computational calculations by Politzer and Murray showed that an anisotropic electron density distribution around the halogen nucleus — the "σ-hole" This hole was then experimentally observed using Kelvin probe force microscopy. In 2020, Kellett et al. showed that halogen bonds also have a π-covalent character similar to metal coordination bonds. In August 2023 the "π-hole" was too experimentally observed ==Applications==

Crystal engineering The strength and directionality of halogen bonds are a key tool in the discipline of crystal engineering, which attempts to shape crystal structures through close control of intermolecular interactions. Halogen bonds can stabilize copolymers or induce mesomorphism in otherwise isotropic liquids. Indeed, halogen bond-induced liquid crystalline phases are known in both alkoxystilbazoles Controlled polymerization Conjugated polymers offer the tantalizing possibility of organic molecules with a manipulable electronic band structure, but current methods for production have an uncontrolled topology. Sun, Lauher, and Goroff discovered that certain amides ensure a linear polymerization of poly(diiododiacetylene). The underlying mechanism is a self-organization of the amides via hydrogen bonds that then transfers to the diiododiacetylene monomers via halogen bonds. Although pure diiododiacetylene crystals do not polymerize spontaneously, the halogen-bond induced organization is sufficiently strong that the cocrystals do spontaneously polymerize. Monomerand6.jpg|The catalyst-monomer cocrystal. Units repeat every 5.25 Å and are oriented at 51.3˚. PIDA.jpg|Post-polymerization crystal structure: the oxygen atom (purple) forms a hydrogen bond (blue dashed line) and a weak halogen bond with the polymer's iodine substituents. Iodine may also form a halogen bond with the terminal nitriles (red dashed line). Biological macromolecules : a short Br−O halogen bond contributes to inhibitor potency. ==Notes==