The Crystal Structure of the 1:1 Addition Compound Formed by Benzyl Sulphide and Iodine.

Rømming, Chr.; Hjortaas, K. E.; Lindgren, Jan‐Erik; Varde, E.; Westin, Gertrud

doi:10.3891/acta.chem.scand.14-2145

Cited by 38 publications

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“…A number of complexes between halides and oxygen/nitrogen Lewis bases were reported and characterized crystallographically in the 1960s, much of this being done by Hassel and his colleagues. Some examples are as follows: SbI 3 ···1,4-dithiane, [4] IÀCCÀI···1,4-dioxane, [5] CHI 3 ···1,4-dithiane (1), [6] CHI 3 ···quinoline, [7] I 2 C=CI 2 ···1,4-diselenane, [8] ClÀC(O)ÀC(O)À Cl···1,4-dioxane, [9] BrÀC(O)ÀC(O)ÀBr···1,4-dioxane (2), [9] and Br 2 C=CBr 2 ···pyrazine (3). [10] These form chain-like structures in the crystal held together by weak X···B bonds, for example, compounds 1-3.…”

Section: The Enigma Of Halogen Bonding: Early Experimental Observationsmentioning

confidence: 99%

“…A remarkably consistent pattern was found. Close contacts with electrophilic sites, for example, metal ions, were largely at angles of 90-1208 with the RÀX axis (4). Close contacts with nucleophilic sites, for example, oxygen and nitrogen atoms, were primarily close to linear, with angles of 160-1808 (5).…”

Section: The Enigma Of Halogen Bonding: Early Experimental Observationsmentioning

confidence: 99%

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Halogen Bonding: An Interim Discussion

2013

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Halogen bonding is a noncovalent interaction that is receiving rapidly increasing attention because of its significance in biological systems and its importance in the design of new materials in a variety of areas, for example, electronics, nonlinear optical activity, and pharmaceuticals. The interactions can be understood in terms of electrostatics/polarization and dispersion; they involve a region of positive electrostatic potential on a covalently bonded halogen and a negative site, such as the lone pair of a Lewis base. The positive potential, labeled a σ hole, is on the extension of the covalent bond to the halogen, which accounts for the characteristic near-linearity of halogen bonding. In many instances, the lateral sides of the halogen have negative electrostatic potentials, allowing it to also interact favorably with positive sites. In this discussion, after looking at some of the experimental observations of halogen bonding, we address the origins of σ holes, the factors that govern the magnitudes of their electrostatic potentials, and the properties of the resulting complexes with negative sites. The relationship of halogen and hydrogen bonding is examined. We also point out that σ-hole interactions are not limited to halogens, but can also involve covalently bonded atoms of Groups IV-VI. Examples of applications in biological/medicinal chemistry and in crystal engineering are mentioned, taking note that halogen bonding can be "tuned" to fit various requirements, that is, strength of interaction, steric factors, and so forth.

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Section: The Enigma Of Halogen Bonding: Early Experimental Observationsmentioning

confidence: 99%

Section: The Enigma Of Halogen Bonding: Early Experimental Observationsmentioning

confidence: 99%

Halogen Bonding: An Interim Discussion

2013

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“…This positive region increases in the order F<Cl<Br<I. The atom A is usually carbon but other bonds such as for example SiX24 or SbX32 have been also considered. The same halogen X atom that possesses the region of the positive electrostatic potential is surrounded by the negative electrostatic potential in the direction perpendicular to the AX bond owing to the existence of “a belt” of free p orbitals.…”

Section: Introductionmentioning

confidence: 99%

“…[23][24][25] This positive region increases in the order F < Cl < Br < I. The atom A is usually carbon but other bonds such as for example Si À X [24] or Sb À X [32] have been also considered. The same halogen X atom that possesses the region of the positive electrostatic potential is surrounded by the Abstract: Ab initio calculations were performed on complexes of ZH 4 + (Z = N, P, As) and their fluoro derivatives, ZFH 3 + and ZF 4 + , with a HCN (or LiCN) molecule acting as the Lewis base through the nitrogen electronegative center.…”

Section: Introductionmentioning

confidence: 99%

σ‐Hole Bond Versus Hydrogen Bond: From Tetravalent to Pentavalent N, P, and As Atoms

Grabowski

2013

Chemistry A European J

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Ab initio calculations were performed on complexes of ZH4 (+) (Z=N, P, As) and their fluoro derivatives, ZFH3 (+) and ZF4 (+) , with a HCN (or LiCN) molecule acting as the Lewis base through the nitrogen electronegative center. It was found that the complexes are linked by the ZH⋅⋅⋅N hydrogen bond or another type of noncovalent interaction in which the tetravalent heavy atom of the cation acts as the Lewis acid center, that is, when the Z⋅⋅⋅N link exists, which may be classified as the σ-hole bond. The formation of the latter interaction is usually preferable to the formation of the corresponding hydrogen bond. The Z⋅⋅⋅N interaction may be also considered as the preliminary stage of the SN 2 reaction. This is supported by the observation that for a short Z⋅⋅⋅N contact, the corresponding complex geometry coincides with the trigonal-bipyramidal geometry typical for the transition state of the SN 2 reaction. The Z⋅⋅⋅N interaction for some of complexes analyzed here possesses characteristics typical for covalent bonds. Numerous interrelations between geometrical, topological and energetic parameters are discussed. The natural bond orbital method as well as the Quantum Theory of "Atoms in Molecules" is applied to characterize interactions in the analyzed complexes. The experimental evidences of the existence of these interactions, based on the Cambridge Structure Database search, are also presented. In addition, it is justified that mechanisms of the formation of the Z⋅⋅⋅N interactions are similar to the processes occurring for the other noncovalent links. The formation of Z⋅⋅⋅N interaction as well as of other interactions may be explained with the use of the σ-hole concept.

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“…Structural results for N-donor Lewis base molecules (B) with dihalogens (X 2 ) and interhalogens (XY) began to appear with regularity in the late 1950s and 1960s with the extensive work of Hassel [36][37][38][39][40][41][42], Stromme [39][40][41][42], Romming [42,45,46], McCullough [47][48][49][50][51], and others [52][53][54]. Since then, a large number of such complexes have been studied.…”

Section: General Observationsmentioning

confidence: 99%