Halogen Bonding in Solution. Группа авторов

concurrently become as fascinated by this topic as all of the authors already are.

      March 2020

       Stefan Huber

      Bochum

1 Halogen Bonding: An Introduction

       Daniel A. Decato, Eric A. John, and Orion B. Berryman

       University of Montana, Department of Chemistry and Biochemistry, 32 Campus Drive, Missoula, MT, 59812 USA

1.1 Introduction

The group 17 elements, known as halogens, are diatomic species in their elemental form with the chemical formula X₂ (where X = F, Cl, Br, I)1. In nature, they are seldom found in this manner due to their reactivity and thus are often presented as covalent or ionic species. However, the physical state of diatomic halogens provides initial insight into their capacity for noncovalent interactions. Moving down group 17, the elemental species exist in different phases from gas (F₂ and Cl₂) to liquid (Br₂) to solid (I₂). For simplicity, this observation is attributed to greater intermolecular dispersion forces afforded by the larger, more polarizable halogens and is quantifiable by physical properties such as boiling and melting point. This is the explanation that young chemists generally receive when being introduced to halogenated species and their physical properties. Later, in many organic curricula, these elements are presented as covalently bound components of molecules and are discussed within the context of molecular and bond dipoles, often in conjunction with concepts of electronegativity. Once again, the enhanced dispersion capacity of the halogens is often highlighted, explaining the higher boiling points of haloalkanes over hydrocarbons of comparable size and shape (e.g. ethane bp = −89 °C and bromomethane bp = 4 °C). Ultimately, the role of halogens in noncovalent interactions is neglected, giving way to their participation in classic reactions such as radical, substitution, and elimination pathways, which predictably leads to the misconception that halogens are simply electronegative leaving groups. This oversight is often reinforced in upper‐level courses, where halogens are shown to be the reactive site in many cross‐coupling reactions. Additionally, in classical inorganic chemistry, halides are depicted as weak field ligands and as prototypical examples in hard–soft acid–base theory. Even discussions of covalently bound halogens participating as hydrogen bond acceptors are atypical in university curricula.

In summary, halogens have traditionally been perceived as electronegative reactive species that participate in weak nondirectional noncovalent interactions (dispersion and as weak hydrogen bond acceptors). So how did scientists discover the ability of halogens to participate in a very directional and potent (comparable with hydrogen bond strength) noncovalent interaction, where the halogen is an electropositive species that is attracted to Lewis bases? To answer this, one should start with the definition of the halogen bond provided by the IUPAC in 2013 [2]. Then, early contributions from scientists who acknowledged an attractive interaction (more significant and directional than dispersion) can be acknowledged. Following the historical contributions, this introduction recounts the rediscovery of the halogen bond near the turn of the twenty‐first century. Finally, the chapter concludes by highlighting impactful nonsolution‐based examples that have helped construct the current understanding of the halogen bond, thereby providing context for the ensuing chapters on solution phase chemistry.

      1.1.1 The Halogen Bond: Definition, Characteristics, Representations, and Parallels to the Hydrogen Bond

      “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”

      IUPAC definition 2013 [2]

The depiction of a covalently bound halogen atom in traditional textbooks is that of an electron‐rich sphere (Figure 1.1a). This simplified description is applicable in many cases and helps account for the behavior of covalently bound halogens as hydrogen bond acceptors [3] and their “side‐on” interactions with metal cations [4]. However, the electron density around halogens is not uniform; the distribution of electron density is anisotropic, resulting in two regions of electron density [5,6] (Figure 1.1b). The region directly involved in halogen bonding is the electropositive region at the tip of the halogen projected away from the covalent bond. This region is termed the sigma hole (σ‐hole), as it is a localized deficit of electron charge opposite a σ‐bond [7]. The second region is the electronegative belt, which forms orthogonal to the covalent bond involving the halogen atom. Thus, an electronic gradient on the surface of the halogen is formed, going from electropositive at the “tip” to electronegative around the equator. The electronic distribution is aptly demonstrated in Figure 1.2, which shows a series of molecular electrostatic potential (ESP) maps of trifluoromethyl halides [8]. Mapping halogenated species in this manner highlights the two distinct regions described above that engender the halogen bond with its characteristic directionality. For instance, if a Lewis base deviates from the σ‐hole region, the interaction becomes less favorable and eventually becomes repulsive as the Lewis base approaches the electron‐rich belt (Figure 1.1b). The σ‐hole is most prominent in the CF₃I molecule and is depicted in this case as the red region of high ESP (Figure 1.2). The region can be quantified and compared with other systems by computing the maximum ESP (V_s,max) on the halogen.

Figure 1.1 Schematic of interactions between a halogen atom and a Lewis base acceptor (A) from a classical view of halogens (a) and from a modern halogen bonding description (b). For comparison the depiction of a hydrogen bond is also included (c). A solid arrow indicates a “stronger” attractive interaction. The dotted arrow is a “less” attractive interaction. The solid arrow with an “X” through it indicates the interaction is repulsive.

Figure 1.2 Molecular electrostatic potential maps drawn at the isodensity surface of 0.001 au for CF₃I, CF₃Br, CF₃Cl, and CF₄. All maps are drawn at the same scale, and values are in kcal/mol.

      Source: From Clark et al. [8]. © 2007 Springer Nature.

      Disclaimer: It is crucial to note that the σ‐hole description does not account for all the nuances of the halogen bond. Therefore, other conceptual approaches and methodologies (e.g. polarizability, charge transfer) can and should be used to fully describe the halogen bond. These particulars are discussed in the computational section of this chapter. Nevertheless, the σ‐hole is widely used, and ESP maps offer a low barrier to