Van der Waals' forces

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In chemistry, the term van der Waals' forces (sometimes called London dispersion forces) refers to a particular class of intermolecular forces. The term originally referred to all such forces, and this usage is still sometimes observed, but it is now more commonly used to refer to those forces which arise from the polarization of molecules into dipoles. This includes forces that arise from fixed or angle-averaged dipoles (Keesom forces) and free or rotation dipoles (Debye forces) as well as shifts in electron cloud distribution (London forces). The name refers to the Dutch physicist and chemist Johannes Diderik van der Waals, who first documented these types of forces. The Lennard-Jones potential is often used as an approximate model for the Van der Waals force as a function of distance.

van der Waals' forces are observed in noble gases, which are very stable and tend not to interact. This is why it is difficult to condense them into liquids. However, the larger the atom of the noble gas (the more electrons it has) the easier it is to condense the gas into a liquid. This happens because, when the electron cloud surrounding the gas atom gets large, it does not form a perfect sphere around the nucleus. Rather, it is only spherical if averaged over longer times and generally forms an ellipsoid, which has a slight negative charge on one side of the major axis and a slight positive charge on the other. The atom becomes a temporary dipole. This induces the same shift in neighboring atoms and spreads from one atom to the next. Unlike charges attract, and the induced dipoles are held together by dispersion force (or Van der Waals force). Van der Waals forces are responsible for certain cases of pressure broadening (van der Waals broadening) of spectral lines.

1 London dispersion force
2 Relation to the Casimir effect
3 Use by animals
4 See also
5 Sources

[edit] London dispersion force

Image:Argon dimer potential.png

London dispersion forces, named after the German-American physicist Fritz London, are weak intermolecular forces that arise from the attractive force between transient dipoles (or better multipoles) in molecules without permanent multipole moments. London dispersion forces are also just called dispersion forces or London forces and sometimes van der Waals' forces.

London forces can be exhibited by nonpolar molecules because electron density moves about a molecule probabilistically, see quantum mechanical theory of dispersion forces. There is a high chance that the electron density will not be evenly distributed throughout a nonpolar molecule. When an uneven distribution occurs, a temporary multipole is created. This multipole may interact with other nearby multipoles. London forces are also present in polar molecules, but they are usually only a small part of the total interaction force.

Electron density in a molecule may be redistributed by proximity to another multipole. Electrons will gather on the side of a molecule that faces a positive charge and retreat from a negative charge. Hence, a transient multipole can be produced by a nearby polar molecule, or even by a transient multipole in another nonpolar molecule.

London forces are weaker than other intermolecular forces such as ionic interactions, hydrogen bonding, or permanent dipole-dipole interactions.

This phenomenon is the only attractive intermolecular force at large distances present between neutral atoms (e.g., helium), and is the major attractive force between non-polar molecules, such as nitrogen or methane (to name a couple). Without London forces, there would be no attractive force between noble gas atoms, and they could not then be obtained in a liquid form.

London forces become stronger as the atom (or molecule) in question becomes larger. This is due to the increased polarizability of molecules with larger, more dispersed electron clouds. This trend is exemplified by the halogens (from smallest to largest: F₂, Cl₂, Br₂, I₂). Fluorine and chlorine are gases at room temperature, bromine is a liquid, and iodine is a solid.

[edit] Relation to the Casimir effect

The London-van der Waals' forces is related to the Casimir effect for dielectric media, the former the microscopic description of the latter bulk property. First detailed calculations of this were done in 1955 by E. M. Lifshitz.

[edit] Use by animals

The van der Waals' forces is the force to which the gecko's climbing ability is attributed. A gecko can hang on a glass surface using only one toe. Efforts continue to create a synthetic "gecko tape" that exploits this knowledge. So far, research has produced some promising results - early research yielded an adhesive tape [1] product, which only obtains a fraction of the forces measured from the natural material, and new research [2] has yielded a discovery that purports 200 times the adhesive forces of the natural material.

Researchers at Stanford University recently developed a gecko-like robot which uses synthetic setae to climb walls [3].

[edit] See also

[edit] Sources

Iver Brevik, V. N. Marachevsky, Kimball A. Milton, Identity of the Van der Waals Force and the Casimir Effect and the Irrelevance of these Phenomena to Sonoluminescence, hep-th/9901011
I. D. Dzyaloshinskii, E. M. Lifshitz, and L. P. Pitaevskii, Usp. Fiz. Nauk 73, 381 (1961)
- English translation: Soviet Phys. Usp. 4, 153 (1961)
R. H. French, University of Pennsylvania, Materials Science "Full Spectral London Dispersion Interaction: Forces and Energies".
L. D. Landau and E. M. Lifshitz, Electrodynamics of Continuous Media, Pergamon, Oxford, 1960, pp. 368–376.
Mark Lefers, "Van der Waals dispersion force". Holmgren Lab.
E. M. Lifshitz, Zh. Eksp. Teor. Fiz. 29, 894 (1955)
- English translation: Soviet Phys. JETP 2, 73 (1956)
Western Oregon University's "London force". Intermolecular Forces. (animation)ca:Força de van der Waals