Waals Interactions and London Dispersion Forces

When the hydrophobic effect brings atoms very close together, van der Waals interactions and London dispersion forces, which work only over very short distances, come into play. This brings things even closer together and squeezes out the holes. The bottom line is a very compact, hydrophobic core in a protein with few holes. 

The polarity of molecules like water has very significant effects on the behavior of these compounds. If you recall in ionic compounds, the oppositely charged ions attract each other and form large crystalline structures. A similar process occurs between polar molecules, but we describe these as intermolecular forces. There are three main intermolecular forces we need to examine. All three of these forces are known as van der Waals forces and are specifically called hydrogen-bonding forces, dipole-dipole interactions, and London dispersion forces. 

Attractive Forces. Attractive forces, collectively called van der Waal s forces, exist between two oil droplets. Simplistically, these forces may be thought of as the attraction between oil molecules at the o/w interfaces that have lower energy when in contact with each other than when in contact with water. Several phenomena are involved hydrophobic interactions and London dispersion forces are most commonly considered. These are effective as (roughly) the fourth power of the distance between the surfaces and are unaffected by ionic strength. The attraction due to van der Waal s forces is shown in Figure 6. Suspensions of solids (cellulose fibers, finely divided CaCOa, etc.) are stabilized in the same way. Ionic surfactants are used that selectively adsorb to the solid surface, generating a v[/ potential and making possible a stable suspension. 

Both dipole-dipole and London Dispersion forces are subclasses of van der Waal interactions. When two polar molecules approach one another, a natural attraction known as dipole-dipole forces is created between oppositely charged ends. The relative intensity of dipole-dipole forces may be represented by Eq. 2 ... 

The main forces responsible for adhesion are van der Waals, which for convenience are considered to be made of three main contributions Dipole-dipole interaction (Keesom force), dipole-induced-dipole interaction (Debye force) and London dispersion force. A hydrogen-bonding force can also be induded in the interaction. 

Intermolecular forces acting between atoms or molecules in a pure substance are called van der Waals forces and include dipole-dipole interactions (including hydrogen bonding) and London dispersion forces (also called simply dispersion forces). 

Attractive forces between neutral molecules may include three contributions according to the nature of the molecule. These three interaction energies are dipole-dipole or Keesom interactions, dipole-induced dipole or Debye interactions, and induced dipole-induced dipole interactions or London dispersion forces (Table 3.1). Since all of them depend on the inverse of the sixth power of the intramolecular distance, they are generally combined in only one term, representing the total van der Waals attraction and this term is the sum of the three energies ... 

Attractive and Repulsive Forces. The force that causes small particles to stick together after colliding is van der Waals attraction. There are three van der Waals forces (/) Keesom-van der Waals, due to dipole—dipole interactions that have higher probabiUty of attractive orientations than nonattractive (2) Debye-van der Waals, due to dipole-induced dipole interactions (ie, uneven charge distribution is induced in a nonpolar material) and (J) London dispersion forces, which occur between two nonpolar substances. 

Although the fact that the cycloamyloses include a variety of substrates is now universally accepted, the definition of the binding forces remains controversial. Van der Waals-London dispersion forces, hydrogen bonding, and hydrophobic interactions have been frequently proposed to explain the inclusion phenomenon. Although no definitive criteria exist to distinguish among these forces, several qualitative observations can be made. 

To design more effective medicines, chemists need to consider the structure of compounds that are active for a particular receptor or active site and to determine the Important functional groups present in these compounds. These functional groups allow the compound to interact with the active site. These Interactions can be any van der Waals forces (hydrogen, permanent dipole-permanent dipole interactions, London dispersion forces) and even ionic bonds. 

Dispersion interactions (also known as van der Waals interactions or London forces) play an incredibly important role in our everyday lives. Consider the gasoline or diesel that all of us rely on as transportation fuels. The huge global infrastructure that exists to deliver and use these fuels relies on the fact that they are liquids. The fact that nonpolar molecules such as hexane readily form liquids is a signature of the net attractive interactions that exist between hexane molecules. These attractive interactions arise directly as a result of electron correlation. 

Two later sections (1.6.5 and 1.6.6) look at the crystalline structures of covalently bonded species. First, extended covalent arrays are investigated, such as the structure of diamond—one of the forms of elemental carbon—where each atom forms strong covalent bonds to the surrounding atoms, forming an infinite three-dimensional network of localized bonds throughout the crystal. Second, we look at molecular crystals, which are formed from small, individual, covalently-bonded molecules. These molecules are held together in the crystal by weak forces known collectively as van der Waals forces. These forces arise due to interactions between dipole moments in the molecules. Molecules that possess a permanent dipole can interact with one another (dipole-dipole interaction) and with ions (charge-dipole interaction). Molecules that do not possess a dipole also interact with each other because transient dipoles arise due to the movement of electrons, and these in turn induce dipoles in adjacent molecules. The net result is a weak attractive force known as the London dispersion force, which falls off very quickly with distance. 

Van der Waals forces are intermolecular and are classified as (i) dipole-dipole interactions, (ii) dipole-induced-dipole interactions and (iii) London dispersion forces which operate between atoms as the result of the nucleus not always being at the centre of mass of the surrounding electrons. The hydrogen bond is regarded as a special form of dipole-dipole interaction, because the positive end of dipolar species containing hydrogen atoms is the relatively unshielded proton. 

Relatively weak forces of attraction that exist between nonpolar molecules are called van der Waals forces or London dispersion forces. Dispersion forces between molecules are much weaker than the covalent bonds within molecules. Electrons move continuously within bonds and molecules, so at any time one side of the molecule can have more electron density than the other side, which gives rise to a temporary dipole. Because the dipoles in the molecules are induced, the interactions between the molecules are also called induced dipole-induced dipole interactions. 

There are two types of solute-solvent interactions which affect absorption and emission spectra. These are universal interaction and specific interaction. The universal interaction is due to the collective influence of the solvent as a dielectric medium and depends on the dielectric constant D and the refractive index n of the solvent. Thus large environmental perturbations may be caused by van der Waals dipolar or ionic fields in solution, liquids and in solids. The van der Waals interactions include (i) London dispersion force, (ii) induced dipole interactions, and (iii) dipole-dipole interactions. These are attractive interactions. The repulsive interactions are primarily derived from exchange forces (non bonded repulsion) as the elctrons of one molecule approach the filled orbitals of the neighbour. If the solute molecule has a dipole moment, it is expected to differ in various electronic energy states because of the differences in charge distribution. In polar solvents dipole-dipole inrteractions are important. 

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