No abstract

This article argues against a close relationship between Lewis dot structures and electron structure obtained from quantum mechanical calculations. Lewis structures are a powerful tool for structure prediction, though they are classical models of bonding and do not predict electronic structure. The "best" Lewis structures are those that, when combined with the VSEPR model, allow the accurate prediction of molecular properties, such as polarity, bond length, bond angle, and bond strength. These structures are achieved by minimizing formal charges within the molecule, even if it requires an expanded octet on atoms beyond the second period. Lewis structures that show an expanded octet do not imply full d-orbital involvement in the bonding. They suggest that the presence of low-lying d-orbitals is important in producing observed molecular structures.Based on this work, the presence of electron density, not a large separation in charge, is responsible for the short bond lengths and large angles in species containing nonmetal atoms from beyond the second period. This result contradicts results obtained from natural population analysis, a method that attempts to derive Lewis structures from molecular orbital calculations.

Intermolecular forces play an important role in many aspects of chemistry ranging from inorganic to biological chemistry. These forces dictate molecular conformation, species aggregation (including self-assembly), trends in solubility and boiling points, adsorption characteristics, viscosity, phase changes, surface tension, capillary action, vapor pressure, ion-pairing, and clathrate formation. Understanding these attractions is important for chemists; yet, visualizing these forces and developing a qualitative "feel" for them often can be very difficult for students. Furthermore, the topic of intermolecular forces can be found throughout all levels of chemical education, from general chemistry to advanced courses. Until recently, these interactions could be described only using static images or mathematical calculations. However, the recent development of widely-available, user-friendly software has provided a new tool for chemical educators to help students investigate and visualize intermolecular forces. This new tool is molecular dynamics simulation (MDS). MDS predicts the motion of chemical species based on the application of empirical rules and a physical analysis of the forces that act between the species. These motions can then be shown in vivid graphical form. This article presents a series of experiments that have been used in courses taught by the authors that can be incorporated readily into the curriculum at almost any high school, college, or university.

The complex (III), formed when (I) is used in excess, crystallizes in the space group P21/n with Z=4; the complex (IV) with semiquinone and catecholate ligands in P1, Z=1.

How to draw the most appropriate Lewis structure for molecules and ions in which nonmetal atoms beyond the second period are bonded to a terminal oxygen atom has been the subject of considerable debate. Recently there has been a shift toward satisfying the octet rule even if formal charges remain on the structure. In this work, the internuclear electron densities of a series of X-O bonds (where X = P, S, or Cl) have been calculated using quantum mechanics and compared to Lewis structures for which the formal charges have been minimized. A direct relationship was found between the internuclear electron density and the bond order predicted from Lewis structures in which formal charges are minimized. This relationship suggests that obeying the octet rule at the expense of having formal charges does not yield the most appropriate Lewis structure. Instructors are urged to teach students to minimize formal charges where possible.