A New Noncovalent Force: Comparison of P∙∙∙N Interaction with Hydrogen and Halogen Bonds

When PH(3) is paired with NH(3), the two molecules are oriented such that the P and N atoms face one another directly, without the intermediacy of a H atom. Quantum calculations indicate that this attraction is due in part to the transfer of electron density from the lone pair of the N atom to the σ(∗) antibond of a P-H covalent bond. Unlike a H-bond, the pertinent hydrogen is oriented about 180° away from, instead of toward, the N, and the N lone pair overlaps with the lobe of the P-H σ(∗) orbital that is closest to the P. In contrast to halogen bonds, there is no requirement of a σ-hole of positive electrostatic potential on the P atom, nor is it necessary for the two interacting atoms to be of differing potential. In fact, the two atoms can be identical, as the global minimum of the PH(3) homodimer has the same structure, characterized by a P···P attraction. Natural bond orbital analysis, energy decomposition, and visualization of total electron density shifts reveal other similarities and differences between the three sorts of molecular interaction

Ability of IR and NMR Spectral Data to Distinguish Between a Tetrel Bond and a Hydrogen Bond

The placement of a nucleophile X along the R-CH3 axis of a methyl group can be described as either a trifurcated H-bond or as a tetrel bond between C and X. Quantum calculation of the vibrational and NMR spectral features of a number of systems are used to provide a framework by which to distinguish between the presence of a tetrel or H-bond. Both cationic and neutral methyl-containing Lewis acids (SMe3+, NMe4+, SMe2) are paired with both neutral and anionic Lewis bases (NH3, OH2, OCHNH2, OCH3-, OH-, HCOO-). As the base is moved away from the R-CH3 axis of the Lewis acid (tetrel bond) toward a position along a CH axis (CH··X H-bond), the methyl stretching frequencies shift to the red and the bending modes to the blue. This modification in the geometry also causes the methyl C atom NMR chemical shielding to increase, while the methyl H atom is progressively deshielded. These changes are of sufficiently large magnitude that they should provide a means to distinguish these two bonding situations from one another

On the Capability of Metal-Halogen Groups to Participate in Halogen Bonds

A number of halogen (X) atoms were covalently attached to a metal (M) and the ability of the X atom to act as electron acceptor in a halogen bond to nucleophile NCH was assessed. Both Cl and Br were considered as halogen atom, with NH3 and CO as other ligands attached to the metal. Metals tested were Ti, Mn, and Zn in various combinations of oxidation state, coordination, and overall charge. In the majority of cases, the strong electron-releasing power of the metal imbues the halogen atom with a high negative partial charge and minimizes the development of a σ-hole. As such, the M atom is generally a stronger attractor for the incoming nucleophile than is the halogen. Nonetheless, there are cases where a halogen bond can form such as Ti(CO)4Br+, TiCl3+, and MnCl4+, each with a different coordination. A requisite of halogen bond formation is generally an overall positive charge, although neutral species can engage in such bonds, albeit much weaker

Interpretation of Spectroscopic Markers of Hydrogen Bonds

Quantum calculations are used to examine whether a AH∙∙D H-bond is unambiguously verified by a downfield shift of the bridging proton’s NMR signal or a red (or blue) shift of the AH stretching frequency in the IR spectrum. It is found that such IR band shifts will occur even if the two groups experience weak or no attractive force, or if they are drawn in so close together that their interaction is heavily repulsive. The mere presence of a proton-acceptor molecule can affect the chemical shielding of a position occupied by a proton-donor by virtue of its electron density, even if there is no H-bond present. This density-induced shielding is heavily dependent on position around the proton-acceptor atom, and varies from one group to another. Evidence of a H-bond rests on the measurement of a proton deshielding in excess of what is caused purely by the presence of the proton acceptor species

The Hydrogen Bond: A Hundred Years and Counting

Since its original inception, a great deal has been learned about the nature, properties, and applications of the H-bond. This review summarizes some of the unexpected paths that inquiry into this phenomenon has taken researchers. The transfer of the bridging proton from one molecule to another can occur not only in the ground electronic state, but in various excited states as well. Study of the latter process has developed insights into the relationships between the nature of the state, the strength of the H-bond, and the height of the transfer barrier. The enormous broadening of the range of atoms that can act as both proton donor and acceptor has led to the concept of the CH···O HB, whose properties are of immense importance in biomolecular structure and function. The idea that the central bridging proton can be replaced by any of various electronegative atoms has fostered the rapidly growing exploration of related noncovalent bonds that include halogen, chalcogen, pnicogen, and tetrel bonds

Forty Years of Progress in the Study of the Hydrogen Bond

The author looks back at developments over the last few decades concerning the H–bond. The list of atoms involved as proton donor and acceptor has broadened dramatically, including most electronegative atoms and even metals. The factors that control the transfer of the proton across the H–bond have been elucidated and show the importance of even minor changes in its geometry. Small stretches can shut down the transfer entirely, and certain bends can force a proton to transfer against a pK gradient. Along with the recognition that a CH..O interaction can represent a true H–bond, and one with strength comparable to more traditional H–bonds, has come an understanding of its contributions to protein structure and function. The replacement of the bridging of H by any of a litany of electronegative atoms leads to similarly strong interactions, with many features virtually indistinguishable from a true H–bond. These noncovalent interactions are typically referred to as halogen, chalcogen, pnicogen, and tetrel bonds, depending upon the identity of the substitute bridging atom

An Updated Description of the Hydrogen Bond and Related Noncovalent Bonds

The hydrogen bond is typically introduced briefly in General Chemistry as a simple electrostatic phenomenon involving a small and select group of atoms, a definition which is typically unchanged through higher levels in the curriculum. But this definition has undergone dramatic modernization of which students should be made aware. The original formulation in terms of only F, O, and N atoms has broadened very considerably, encompassing C as well as atoms from lower rows in the periodic table. The influence of hybridization, substituents, and overall charge cannot be overlooked. In addition to the Coulombic attraction, there are other “covalent” contributors such as charge transfer and polarization. Further broadening has occurred with the recognition that the bridging H can be replaced by a host of electronegative atoms in what have come to be denoted halogen, chalcogen, pnicogen, and tetrel bonds, with behavior very similar to H-bonds

Special Issue: Intramolecular Hydrogen Bonding 2017

Even after more than a century of study [1–6], scrutiny, and detailed examination, the H-bond continues [7–12] to evoke a level of fascination that surpasses many other phenomena. Perhaps it is the ability of the simple H atom, with but a single electron, to act as a glue that maintains contact between much more complicated species. Or it might be its geometry, which prefers to hold the bridging proton on a direct line between the two heavy atoms. Not to be ignored are the spectral features of the H-bond: the large red shift of the stretching frequency of the covalent A–H bond, coupled with its intensification, or the downfield shift of the proton’s NMR signal. Yet study of this bond is far from complete, with one surprise after another continuing to emerge. As it turns out, the aforementioned red shift, for example, long considered as the trademark of this bond, is not so characteristic after all. H-bonds that shift in the opposite direction, to the blue, have been observed [13–16] in a variety of systems. The long held belief that only very electronegative atoms like F, O, and N can participate in these bonds has been shattered, as one atom after another, S and Cl and even metals to name just a few, have been added [17–20] to the rapidly growing list

Halogen Bonds Formed Between Substituted Imidazoliums and N Bases of Varying N-Hybridization

Heterodimers are constructed containing imidazolium and its halogen-substituted derivatives as Lewis acid. N in its sp3, sp2 and sp hybridizations is taken as the electron-donating base. The halogen bond is strengthened in the Cl \u3c Br \u3c I order, with the H-bond generally similar in magnitude to the Br-bond. Methyl substitution on the N electron donor enhances the binding energy. Very little perturbation arises if the imidazolium is attached to a phenyl ring. The energetics are not sensitive to the hybridization of the N atom. More regular patterns appear in the individual phenomena. Charge transfer diminishes uniformly on going from amine to imine to nitrile, a pattern that is echoed by the elongation of the C-Z (Z=H, Cl, Br, I) bond in the Lewis acid. These trends are also evident in the Atoms in Molecules topography of the electron density. Molecular electrostatic potentials are not entirely consistent with energetics. Although I of the Lewis acid engages in a stronger bond than does H, it is the potential of the latter which is much more positive. The minimum on the potential of the base is most negative for the nitrile even though acetonitrile does not form the strongest bonds. Placing the systems in dichloromethane solvent reduces the binding energies but leaves intact most of the trends observed in vacuo; the same can be said of ∆G in solution

Proton transfers in hydrogen‐bonded systems. VI. Electronic redistributions in (N2H7)+ and (O2H5)+

Electronic rearrangements accompanying transfer of the central proton between the two XHn units of (H3NHNH3)+ and (H2OHOH2)+ are studied using ab initio molecular orbital methods. Electron density difference maps are calculated by subtracting the density of the equilibrium structure (X–H‐‐‐X) from that of the midpoint geometry (X‐‐H‐‐X) using the split‐valence 4‐31G basis set. Some of the features revealed by the maps are common to both systems while others indicate significant differences between nitrogen and oxygen. Decomposition of the total electron density into contributions from individual occupied molecular orbitals (MOs) provides insight into the factors responsible for the overall charge migrations. The orbitals of a1 symmetry lead to density shifts in a direction parallel to the H bond axis. Among the features attributed to these MOs are the charge transfer across the H bond from one molecule to the other and characteristic density changes in the lone pair regions of the first‐row X atoms. Internal polarizations of the XH bonds of each molecule arise from the density shifts perpendicular to the H bond axis associated with the MOs of non‐a1 symmetry. Simple arguments involving electrostatic and covalent effects are used to explain the redistributions observed in the various MOs. Mulliken analyses provide information, complementary to the difference maps, concerning the relative involvement of various atoms and atomic orbitals in the electronic redistributions associated with each MO