Electrostatic interactions London forces

As is demonstrated in Figure 1, electro-optic coefficients increase linearly with chromophore number density (with a slope proportional to pP) only at low chromophore loading. At higher concentrations, deviation from linearity is observed and the problem becomes more severe as chromophores with higher pp are studied. In particular, maximum in plots of r versus N are observed and these shift to lower N for the higher pp chromophores. The problem can be traced to increasing chromophore electrostatic interactions as is illustrated in Table II. In this Table, we illustrate the variation of dipole moment, polarizability, and ionization potential (the parameters which define chromophore-chromophore electrostatic interactions-London forces) with molecular hyperpolarizability. 

At any rate, a minimum does represent a situation in which attractions and repulsions are balanced (Brehmer et al. 2000). The nomenclature of these intermolecular interactions is quite variegated and the terms are not always clearly defined or distinguished from one another. Some in common usage include van der Waals interactions, London forces, dipole-dipole interactions (and higher terms), dispersion forces, steric repulsion, hydrogen bonds, charge-transfer interactions (also called donor-acceptor interactions), electrostatic interactions, exchange repulsion forces, etc. 

When two molecules interact with each other, several types of electrostatic interactions or forces may be involved, some of which have been described in the preceding sections (e.g., charge-charge, charge-dipole, dipole-dipole interactions). Here, we would like to mention two other kinds of electrical interactions which were not described above namely, short-range repulsive interactions and the London dispersion interaction. The latter interaction plays an especially important role in biological systems. 

In general, polarizability increases as the orbital increases in size negative electrons orbit the positive nucleus at a greater distance in such atoms, and consequently experience a weaker electrostatic interaction. For this reason, London dispersion forces tend to be stronger between molecules that are easily polarized, and weaker between molecules that are not easily polarized. 

Attractive or repulsive forces between molecular entities or groups within the same molecular entity (i.e., both intermolecular and intramolecular) not due to bond formation or to electrostatic interactions of ions or ionic groups with one another or with neutral molecules. The origin of van der Waals forces is in electric polarization of uncharged atoms, groups, or molecules and includes dipole-dipole interactions, dipole-induced dipole interactions, and London forces (induced dipole-induced dipole interactions). 

Polar molecules such as ethyl chloride and PVC are attracted to each other by both the London forces, but also by dipole-dipole interactions resulting from the electrostatic... 

The forces involved in the interaction al a good release interface must be as weak as possible. They cannot be the strong primary bonds associated with ionic, covalent, and metallic bonding neither arc they the stronger of the electrostatic and polarization forces that contribute to secondary van der Waals interactions. Rather, they are the weakest of these types of forces, the so-called London or dispersion forces that arise from interactions of temporary dipoles caused by fluctuations in electron density. They are common to all matter. The surfaces that are solid at room temperature and have the lowest dispersion-force interactions are those comprised of aliphatic hydrocarbons and fluorocarbons. 

It has been traditional to define a van der Waals potential (which combines Coulomb s law and the Lennard-Jones 6-12 potential function) and thereby subsume electronic shell repulsion, London forces, and electrostatic interactions under the term van der Waals interaction. Unfortunately, the resulting expression is an oversimplified treatment of the electrostatic interactions, which are only calculated between close neighbors and are considered to be spatially isotropic. Both of these implicit assumptions are untrue and do not represent physically realistic approximations. We prefer to use the term van der Waals distance for the intemuclear separation at which the 6-12 potential function is a minimum (see Fig. 6), the van der Waals radius being one-half this value when the two interacting atoms are identical, and explicitly treat the Lennard-Jones and electrostatic terms separately. While the term van der Waals interaction may have some value as a shorthand in structure description, it should be avoided when energetics are treated quantitatively. 

Electrostatic interactions occur between the ionic head groups of the surfactant and the oppositely charged solid surface (head down adsorption with monolayer structure) [56]. Acid-base interactions occur due to hydrogen bonding or Lewis acid-Lewis base reactions between solid surface and surfactant molecules (head down with monolayer structure) [57]. Polarisation of jt electrons occurs between the surfactant head group which has electron-rich aromatic nuclei and the positively charged solid surface (head down with monolayer structure) [58]. Dispersion forces occur due to London-van der Waals forces between the surfactant molecules and the solid surface (hydrophobic tail lies flat on the hydrophobic solid surface while hydrophilic head orients towards polar liquid) [59]. 

It is of importance for a knowledge of the forces acting between colloidal particles that the greatest distance at which the London forces are still important is not the radius of the atom but in fact of the order of magnitude of the radius of the particle itself, since the interaction between all the atoms in each of the colloidal particles must be summed, and this interaction, therefore, will increase with increasing size of the particles (Hamaker)1. This is quite different from, for example, the interaction between particles with a crystal lattice in which only purely electrostatic forces would act in this case the radius of action remains, even for large particles, of the order of the lattice constant and there is only a question of a surface action. The effect of the more deeply situated parts of the lattice does not appear outside on account of the mutual compensation of the action of the oppositely charged ions. 

In the theory developed by Derjaguin and Landau (24) and Verwey and Overbeek (25.) the stability of colloidal dispersions is treated in terms of the energy changes which take place when particles approach one another. The theory involves estimations of the energy of attraction (London-van der Walls forces) and the energy of repulsion (overlapping of electric double layers) in terms of inter-oarticle distance. But in addition to electrostatic interaction, steric repulsion has also to be considered. 

Even though the components of the total interaction potential between such complex adsorbents as solid carbons and a wide range of adsorbates can be grouped in many different ways 315,316], it is convenient and meaningful to consider only the London dispersion (induced dipole) forces and the electrostatic (double-layer) forces [620,621,76,77]. 

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