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Halogen bond

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A halogen bond 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.[1] Like a hydrogen bond, the result is not a formal chemical bond, but rather a strong electrostatic attraction.[2][3] 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;[2][3][4] drug design and biochemistry;[5][6] crystal engineering[6] and liquid crystals;[2] and organic catalysis.[6]

Definition

Halogen bond in complex between iodine monochloride 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.[7] 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.[2][8] Consequently halogen-bond propensity follows the trend[9][Note 1]

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.[10][11][12][13]

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 angles are about 120°.[14]
  • Carbonyl, thiocarbonyl-, and selenocarbonyl groups, with a trigonal planar geometry around the Lewis donor atom, can accept one or two halogen bonds.[15]

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.[16]

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 20 kJ/mol; known halogen bond energies range from 10–200 kJ/mol.[15]

The σ-hole concept readily extends to pnictogen, chalcogen and aerogen bonds, corresponding to atoms of Groups 15, 16 and 18 (respectively).[17]

History

Chains in the 1:1 adduct of 1,4-dioxane and bromine, the first crystallographic evidence of halogen bonding.

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.[18] 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.[19]

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.[20] The patterns suggested that only 2.71 Å separated the dioxane oxygens 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.[21] For this work, Hassel was awarded the 1969 Nobel Prize in Chemistry.[22]

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.[23]

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.[24]

In 2007, computational calculations by Politzer and Murray showed that an anisotropic electron density distribution around the halogen nucleus — the "σ-hole"[8] — underlay the high directionality of the halogen bond.[25] This hole was then experimentally observed using Kelvin probe force microscopy.[26]

Applications

Crystal engineering

Br···O halogen bonds observed in the 3D crystal structure of certain silsesquioxanes.[27]

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.[28] Halogen bonds can stabilize copolymers[29][30] or induce mesomorphism in otherwise isotropic liquids.[31] Indeed, halogen bond-induced liquid crystalline phases are known in both alkoxystilbazoles[31] and silsesquioxanes (pictured).[27] Alternatively, the steric sensitivity of halogen bonds can cause bulky molecules to crystallize into porous structures; in one notable case, halogen bonds between iodine and aromatic π-orbitals caused molecules to crystallize into a pattern that was nearly 40% void.[32]

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.[33]

  • The catalyst-monomer cocrystal. Units repeat every 5.25 Å and are oriented at 51.3˚.
    The catalyst-monomer cocrystal. Units repeat every 5.25 Å and are oriented at 51.3˚.
  • 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).
    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

IDD 594 binding to human aldose reductase: a short Br−O halogen bond contributes to inhibitor potency.[34]

Most biological macromolecules contain few or no halogen atoms. But when molecules do contain halogens, halogen bonds are often essential to understanding molecular conformation. Computational studies suggest that known halogenated nucleobases form halogen bonds with oxygen, nitrogen, or sulfur in vitro. Interestingly, oxygen atoms typically do not attract halogens with their lone pairs, but rather the π electrons in the carbonyl or amide group.[35]

Halogen bonding can be significant in drug design as well. For example, inhibitor IDD 594 binds to human aldose reductase through a bromine halogen bond, as shown in the figure. The molecules fail to bind to each other if similar aldehyde reductase replaces the enzyme, or chlorine replaces the drug halogen, because the variant geometries inhibit the halogen bond.[34]

Notes

  1. ^ Although hydrogen is sometimes considered a halogen, convention excludes hydrogen bonds from the category of halogen bonds. For a complete analysis, see § Comparison to other bond-like forces.

References

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  2. ^ a b c d Metrangolo P, Neukirch H, Pilati T, Resnati G (May 2005). "Halogen bonding based recognition processes: a world parallel to hydrogen bonding". Accounts of Chemical Research. 38 (5): 386–395. doi:10.1021/ar0400995. PMID 15895976.
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  13. ^ Turunen L, Hansen JH, Erdélyi M (May 2021). "Halogen Bonding: An Odd Chemistry?". Chemical Record. 21 (5): 1252–1257. doi:10.1002/tcr.202100060. PMID 33939244. S2CID 233483539.
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    • Mulliken RS (1952). "Molecular Compounds and their Spectra. II". J. Am. Chem. Soc. 74 (3): 811–824. doi:10.1021/ja01123a067.
    • Mulliken RS (1952). "Molecular Compounds and their Spectra. III. The Interaction of Electron Donors and Acceptors". J. Phys. Chem. 56 (7): 801–822. doi:10.1021/j150499a001.
  20. ^ Hassel O, Hvoslef J (1954). "The Structure of Bromine 1,4-Dioxanate" (PDF). Acta Chem. Scand. 8: 873. doi:10.3891/acta.chem.scand.08-0873.
  21. ^ Hassel O (October 1970). "Structural aspects of interatomic charge-transfer bonding". Science. 170 (3957): 497–502. Bibcode:1970Sci...170..497H. doi:10.1126/science.170.3957.497. PMID 17799698.
  22. ^ Hassel O (1972). "Structural Aspects of Interatomic Charge-Transfer Bonding". In Nobel Lectures, Chemistry 1963-1970: 314–329.
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  24. ^ Legon AC (September 1999). "Prereactive Complexes of Dihalogens XY with Lewis Bases B in the Gas Phase: A Systematic Case for the Halogen Analogue B···XY of the Hydrogen Bond B···HX". Angewandte Chemie. 38 (18): 2686–2714. doi:10.1002/(sici)1521-3773(19990917)38:18<2686::aid-anie2686>3.0.co;2-6. PMID 10508357.
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  26. ^ Mallada B, Gallardo A, Lamanec M, de la Torre B, Špirko V, Hobza P, Jelinek P (November 2021). "Real-space imaging of anisotropic charge of σ-hole by means of Kelvin probe force microscopy". Science. 374 (6569): 863–867. doi:10.1126/science.abk1479. PMID 34762455. S2CID 244039573.
  27. ^ a b Janeta M, Szafert S (2017-10-01). "Synthesis, characterization and thermal properties of T8 type amido-POSS with p-halophenyl end-group". Journal of Organometallic Chemistry. 847: 173–183. doi:10.1016/j.jorganchem.2017.05.044. ISSN 0022-328X.
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  29. ^ Corradi, Eleanora; Meille, Stefano V.; Messina, Maria T.; Metrangolo, Pierangelo; Resnati, Giuseppe (2000). "Halogen Bonding versus Hydrogen Bonding in Driving Self-Assembly Processes". Communications. Angewandte Chemie International Edition. 39 (10). Weinheim: Wiley-VCH: 1782–1786.
  30. ^ Amico, Vincenzo; Meille, Stefano Valdo; Corradi, Eleonora; Messina, Maria Teresa; Resnati, Giuseppe (1998-08-01). "Perfluorocarbon−Hydrocarbon Self-Assembling. 1D Infinite Chain Formation Driven by Nitrogen···Iodine Interactions". Journal of the American Chemical Society. 120 (32): 8261–8262. doi:10.1021/ja9810686. ISSN 0002-7863.
  31. ^ a b Nguyen HL, Horton PN, Hursthouse MB, Legon AC, Bruce DW (January 2004). "Halogen bonding: a new interaction for liquid crystal formation". Journal of the American Chemical Society. 126 (1): 16–17. doi:10.1021/ja036994l. PMID 14709037.
  32. ^ Pigge FC, Vangala VR, Kapadia PP, Swenson DC, Rath NP (October 2008). "Hexagonal crystalline inclusion complexes of 4-iodophenoxy trimesoate". Chemical Communications. 38 (39): 4726–4728. doi:10.1039/b809592b. PMID 18830473. S2CID 40424594.
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  34. ^ a b Howard EI, Sanishvili R, Cachau RE, Mitschler A, Chevrier B, Barth P, et al. (June 2004). "Ultrahigh resolution drug design I: details of interactions in human aldose reductase-inhibitor complex at 0.66 A". Proteins. 55 (4): 792–804. doi:10.1002/prot.20015. PMID 15146478. S2CID 38388856. The electrostatic interaction between the Br atom of the inhibitor and the OG of Thr 113 has an unusually short distance of 2.973(4) Å. The short contact between Br and Thr 113 OG explains the selectivity of IDD 594 towards AR, because in aldehyde reductase the Thr residue is replaced by Tyr....The IDD 594-Br/Thr 113-OG interaction also contributes to the potency of the inhibitor. Other halogens, such as chlorine, cannot engage in a similar interaction (due to its lower polarizability).
  35. ^ Auffinger P, Hays FA, Westhof E, Ho PS (November 2004). "Halogen bonds in biological molecules". Proceedings of the National Academy of Sciences of the United States of America. 101 (48): 16789–16794. Bibcode:2004PNAS..10116789A. doi:10.1073/pnas.0407607101. PMC 529416. PMID 15557000.

Further reading

  • An early review: Bent, H. A. (1968). "Structural Chemistry of Donor-Acceptor Interactions". Chem. Rev. 68 (5): 587–648. doi:10.1021/cr60255a003.
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