London interactions

Komasa J and Thakkar A J 1995 Accurate Heitler-London interaction energy for He2 J. Mol. Struct. (Theochem) 343 43... 

The total van der Waals interaction potential is obtained by simply adding the individual contributions arising from the Keesom, Debye, and London interactions. Because the radial power-law dependencies of all these interactions vary as 1 /r, the total van der Waals interaction can be expressed simply as... 

We assume that the double bonds in 1,3-butadiene would be the same as in ethylene if they did not interact with one another. Introduction of the known geometry of 1,3-butadiene in the s-trans conformation and the monopole charge of 0.49 e on each carbon yields an interaction energy <5 — 0.48 ev between the two double bonds. Simpson found the empirical value <5 = 1.91 ev from his assumption that only a London interaction was present. Hence it appears that only a small part of the interaction between double bonds in 1,3-butadiene is a London type of second-order electrical effect and the larger part is a conjugation or resonance associated with the structure with a double bond in the central position. 

Closely related to the London interaction is the dipole-induced-dipole interaction, in which a polar molecule interacts with a nonpolar molecule (for example, when oxygen dissolves in water). Like the London interaction, the dipole—induced-dipole interaction arises from the ability of one molecule to induce a dipole moment in the other. However, in this case, the molecule that induces the dipole moment has a permanent dipole moment. The potential energy of the interaction is... 

Once again, the potential energy is inversely proportional to the sixth power of the separation. Notice that the potential energies of the dipole-dipole interaction of rotating polar molecules in the gas phase, the London interaction, and the dipole-induced-dipole interaction all have the form... 

SOLUTION The data in Fig. 2.12 show that electronegativity differences decrease from 1 HC1 to HI, and so the dipole moments decrease as well. Therefore, dipole-dipole forces decrease, too, a trend suggesting that the boiling points should decrease from HQ to HI. This prediction conflicts with the data so we examine the London forces. The number of electrons in a molecule increases from HQ to HI, and so the strength of the London interaction increases, too. Therefore, the boiling points should increase from HCl to HI, in accord with the data. This analysis suggests that London forces dominate dipole-dipole interactions for these molecules. 

Answer The strength of the London interaction increases as the number of electrons increases. ... 

The London interaction arises from the attraction between instantaneous electric dipoles on neighboring molecules and acts between all types of molecules its strength increases with the number of electrons and occurs in addition to any dipole-dipole interactions. Polar molecules also attract nonpolar molecules by weak dipole-induced-dipole interactions. 

The London interaction is universal in the sense that it applies to all molecules regardless of their chemical identity. Similarly, the dipole-dipole interaction depends only on the polarity of the molecule, regardless of its chemical identity. However, there is another very strong interaction between molecules that is specific to molecules with certain types of atoms. 

FIGURE 8.19 Sulfur, which is a solid with nonpolar molecules, does not dissolve in water (left), but it does dissolve in carbon disulfide (right), with which the S8 molecules have favorable London interactions. 

IB 1, l-dichloroethane, because it has a dipole moment 5.2B Unlike CF4, CHh has a net dipole moment. The resulting dipole-dipole interactions account for the higher boiling point of CHF, even though one might expect the CF4 molecule (with more electrons) to exhibit stronger London interactions. 

Adsorption can be attributed to the following interactions van der Waals-London interactions, charge transfer/hydrogen bonding, ligand exchanges, ion exchange, direct and induced ion-dipole... 

Van der Waals-London interactions are due to fluctuations in electron distribution as the electrons circulate within their orbits. These instantaneous dipoles are usually weak, but are, regardless, the most common interaction resulting in adsorption.31 Stronger interactions result from charge transfer. 

Certain types of non-covalent interactions such as hydrogen bonds, London interactions and van der Waals interactions are enthalpy driven interactions (26) heat is released during bond formation. The heat released during bond formation stabilizes the bonds. Hydrogen bonds, London interactions and van der Waals interactions are variants on the dipole-dipole interaction model, which include permanent and induced dipoles. 

Almost all interfacial phenomena are influenced to various extents by forces that have their origin in atomic- and molecular-level interactions due to the induced or permanent polarities created in molecules by the electric fields of neighboring molecules or due to the instantaneous dipoles caused by the positions of the electrons around the nuclei. These forces consist of three major categories known as Keesom interactions (permanent dipole/permanent dipole interactions), Debye interactions (permanent dipole/induced dipole interactions), and London interactions (induced dipole/induced dipole interactions). The three are known collectively as the van der Waals interactions and play a major role in determining material properties and behavior important in colloid and surface chemistry. The purpose of the present chapter is to outline the basic ideas and equations behind these forces and to illustrate how they affect some of the material properties of interest to us. 

FIG. 10.4 A linear arrangement of two dipoles used to define the potential energy in the Schrodinger equation for the London interaction energy. 

As we have already noted, all molecules display the dispersion component of attraction since all are polarizable and that is the only requirement for the London interaction. Not only is the dispersion component the most ubiquitous of the attractions, but it is also the most important in almost all cases. Only in the case of highly polar molecules such as water is the dipole-dipole interaction greater than the dispersion component. Likewise, the mixed interaction described by the Debye equation is generally the smallest of the three. 

Assume that the intermolecular attraction in this case is dominated by the London interaction, that is, = -Cr 6, for which C is the London parameter (Equation (25)). Estimate the London constant from the equation-of-state data and compare it with the coefficient from Equation (25). The polarizability [cco tteo)] for CH4 is 2.6 10-3° m3. The ionization energy / is 2.0185 10 18J (Israelachvili 1991, Chapter6 and Table 6.1). 

The strength of the London interaction depends on the polarizability, a (alpha), the ease with which the electron cloud can be distorted. This dependence is reasonable, because the nuclei in highly polarizable molecules have only weak control over the surrounding electrons, so there can be big fluctuations in electron density and hence large instantaneous partial charges. It turns out that the potential energy of the London interaction varies as the sixth power of the separation of two molecules ... 

An important point to note is that, like the potential energy of dipole-dipole interactions between rotating molecules, the potential energy of the London interaction also decreases very rapidly with distance (as 1/r6 see Fig. 5.1). 

The strength of the London interaction also depends on the shapes of the molecules. Both pentane (8) and 2,2-dimethylpropane (9), for instance, have the molecular formula C5H12, so they each have the same number of electrons. 

London interactions. However, although 2-propanone (acetone), C3H60, has approximately the same number of electrons as butane, it is a liquid that boils at 56°C. In its case, the interactions between the polar carbonyl groups occur in addition to the London interaction, so the total interaction between the molecules is stronger than for butane. 

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