Attractive forces orientational

Van der Waals complexes can be observed spectroscopically by a variety of different teclmiques, including microwave, infrared and ultraviolet/visible spectroscopy. Their existence is perhaps the simplest and most direct demonstration that there are attractive forces between stable molecules. Indeed the spectroscopic properties of Van der Waals complexes provide one of the most detailed sources of infonnation available on intennolecular forces, especially in the region around the potential minimum. The measured rotational constants of Van der Waals complexes provide infonnation on intennolecular distances and orientations, and the frequencies of bending and stretching vibrations provide infonnation on how easily the complex can be distorted from its equilibrium confonnation. In favourable cases, the whole of the potential well can be mapped out from spectroscopic data. 

There are two principal forces that govern the abdity of a polymer to crystallise the interchain attractive forces, which are a function of the chain stmcture, and the countervailing kinetic energy of the chain segments, which is a function of the temperature. The fact that polymers consist of long-chain molecules also iatroduces a third parameter, ie, the imposition of a mechanical force, eg, stretching, which can also enhance interchain orientation and favor crystallisation. 

Polar molecules, like nonpolar molecules, are attracted to one another by dispersion forces. In addition, they experience dipole forces as illustrated in Figure 9.9, which shows the orientation of polar molecules, such as Id, in a crystal. Adjacent molecules line up so that the negative pole of one molecule (small Q atom) is as dose as possible to the positive pole (large I atom) of its neighbor. Under these conditions, there is an electrical attractive force, referred to as a dipole force, between adjacent polar molecules. 

Solids and liquids are called condensed phases. The attractive forces in a condensed phase, either a solid or a liquid, tend to hold the molecules close together. In liquids, molecules are irregularly spaced and randomly oriented. In a crystalline solid, the molecules occupy regular positions, resulting in additional stability (relative to the liquid). 

London23 has treated the case of the attractive force between anisotropic molecules on the dipole-dipole interaction basis as well as on the monopole basis mentioned above. The small anisotropy found for the chlorine atom makes the dipole-dipole formulation appropriate. For the symmetrical orientation in the Cl2 molecule the London formula is... 

One of the primary features of the Gay-Berne potential is the presence of anisotropic attractive forces which should allow the observation of thermally driven phase transitions and this has proved to be the case. Thus using the parametrisation proposed by Gay and Berne, Adams et al. [9] showed that GB(3.0, 5.0, 2, 1) exhibits both nematic and isotropic phases on varying the temperature at constant density. This was chosen to be close to the transitional density for hard ellipsoids with the same ellipticity indeed it is generally the case that to observe a nematic-isotropic transition for Gay-Berne mesogens the density should be set in this way. The long range orientational order of the phase was established from the non-zero values of the orientational correlation coefficient, G2(r), at large separations and the translational disorder was apparent from the radial distribution function. 

Trouton s rule is obeyed most closely by liquids that do not have a high degree of order within the liquid. In both HF and CH3OH, hydrogen bonds create considerable order within the liquid. In C6H5CH3, the only attractive forces are non-directional London forces, which cause the molecules to attract each other, but have no preferred orientation as... 

The value of the effective van der Waals radius of an atom in a crystal depends on the strength of the attractive forces holding the molecules together, and also on the orientation of the contact relative to the covalent bond or bonds formed by the atom (as discussed below) it is accordingly much more variable than the corresponding covalent radius. In Table 7-20 there are given the ionic radii of nonmet llic elements for use as van der Waals radii. They have been rounded off... 

In addition to being additive, intermolecular forces can also be competitive. More specifically, the orientation of the molecules that is optimal to satisfy one type of interaction may not be ideal for another type. The result is a compromise amongst the individual attractive forces to achieve the lowest potential energy state for the crystal. Furthermore, AHm is not the energy required to completely eliminate intermolecular attraction at Tm but the energy required to reduce the attraction to its level in the liquid that level being a function of intermolecular distance. 

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