Other Intermolecular Forces

SCHEME 7.R.3 van der Waals interactions between two I2 molecules (iodine atoms shown in red), which induce dipoles on each other. Both dipole orientations are shown, on the left and on the right. 

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Comment on the validity of the following statement Dispersion forces are weak in comparison to other intermolecular forces. ... 

In the absence of electron donor-acceptor interactions, the London dispersive energy is the dominant contributor to the overall attractions of many molecules to their surroundings. Hence, understanding this type of intermolecular interaction and its dependency on chemical structure allows us to establish a baseline for chemical attractions. If molecules exhibit stronger attractions than expected from these interactions, then this implies the importance of other intermolecular forces. To see the superposition of these additional interactions and their effect on various partitioning phenomena below, we have to examine the role of dispersive forces in more detail,... 

As described in Section 3.1.2, intermolecular interactions between a solute and a solvent or solvents play an important role to dissolve the solute molecule. A number of intermolecular forces have been proposed that invariably contained the masses of the molecules. The molecular properties of materials depend not on the quantity of molecules but on the forces between molecules in close proximity to each other. Intermolecular forces may be divided into three categories ... 

We are cognizant that the structural data on hydrogen bonds derived from crystal structure analyses refer particularly to the hydrogen bond in the solid state. These data are subject to crystal field effects caused by other intermolecular forces. Just as with any discussion of covalent or ionic bonds observed in crystals, these crystal field effects have to be taken into consideration when extrapolating from the precise data available from the crystalline state to the imprecise data that applies to the liquid state in which most chemical and biochemical reactions take place. 

Directional properties of acceptor groups are very soft or nonexistent. Just as the H A lengths and X-ft- -A angles for individual hydrogen bonds are perturbed by the other intermolecular forces in the crystal, so are the directional acceptor properties of the functional groups. These properties appear to be even softer than the hydrogen-bond geometries, and in consequence it is not possible to extract any characteristic trends from the surveys of the type described in this chapter. 

The no-bond function includes the electronic energy of the component molecules, plus terms representing the effect of dipole interactions, dispersion forces, hydrogen bonding and other intermolecular forces. The dative bond functions represent states where an electron has been transferred from the donor molecule to the acceptor, introducing electrostatic interactions and forming a weak covalent link between the resulting radical ions ... 

To provide a more quantitative explanation of the magnitudes of the properties of different materials, we must consider several types of intermolecular forces in greater detail than we gave to the Lennard-Jones model potential in Chapter 9. The Lennard-Jones potential describes net repulsive and attractive forces between molecules, but it does not show the origins of these forces. We discuss other intermolecular forces in the following paragraphs and show how they arise from molecular structure. Intermolecular forces are distinguished from intramolecular forces, which lead to the covalent chemical bonds discussed in Chapters 3 and 6. Intramolecular forces between atoms in the covalent bond establish and maintain... 

This enhancement of a metal coordination binding site with the potential for molecular recognition via other intermolecular forces allows the synthesis of anion receptors with novel anion selectivities. 

Other intermolecular forces can also be used for the purpose of enhancing selectivity. One approach to enhancing electrostatic interactions in order to enforce a further degree of selectivity is to use hydrogen bonding. This methodology is outlined in Section V.D. 

When this theory was used to predict the solubility of polymers in a variety of solvents, it was only partially successful. It was apparent that other intermolecular forces were at work which could not be calculated by this simple procedure. Hydrogen bonding, probably the strongest type of intermolecular force in a nonelectrolyte, was the clue for making solubility parameter theory work. 

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