# London dispersion force

Interaction energy of argon dimer. The long-range section is due to London dispersion forces

London dispersion forces (LDF, also known as dispersion forces, London forces, instantaneous dipole–induced dipole forces, or loosely van der Waals forces) are a type of force acting between atoms and molecules.[1] They are part of the van der Waals forces. The LDF is named after the German-American physicist Fritz London.

## Introduction

The electron distribution around an atom or molecule undergoes fluctuations in time. These fluctuations create instantaneous electric fields which are felt by other nearby atoms and molecules, which in turn adjust the spatial distribution of their own electrons. The net effect is that the fluctuations in electron positions in one atom induce a corresponding redistribution of electrons in other atoms, such that the electron motions become correlated. While the detailed theory requires a quantum-mechanical explanation (see quantum mechanical theory of dispersion forces), the effect is frequently described as the formation of instantaneous dipoles that (when separated by vacuum) attract each other. The magnitude of the London dispersion force is frequently described in terms of a single parameter called the Hamaker constant, typically symbolized A. For atoms that are located closer together than the wavelength of light, the interaction is essentially instantaneous and is described in terms of a "non-retarded" Hamaker constant. For entities that are farther apart, the finite time required for the fluctuation at one atom to be felt at a second atom ("retardation") requires use of a "retarded" Hamaker constant.[2][3]

While the London dispersion force between individual atoms and molecules is quite weak and decreases quickly with separation (R) like ${\displaystyle {\frac {1}{R^{6}}}}$ , in condensed matter (liquids and solids), the effect is cumulative over the volume of materials, or within and between organic molecules, such that London dispersion forces can be quite strong in bulk solid and liquids and decays much more slowly with distance. For example, the total force per unit area between two bulk solids decreases like ${\displaystyle {\frac {1}{R^{3}}}}$ [citation needed] where R is the separation between them. The effects of London dispersion forces are most obvious in systems that are very non-polar (e.g., that lack ionic bonds), such as hydrocarbons and highly symmetric molecules such as bromine (Br2, a liquid at room temperature), iodine (I2, a solid at room temperature). In hydrocarbons and waxes the dispersion forces are sufficient to cause condensation from the gas phase into the liquid or solid phase. Sublimation heats of e.g. hydrocarbon crystals reflect the dispersion interaction. Liquification of oxygen and nitrogen gases into liquid phases is also dominated by attractive London dispersion forces.

When atoms/molecules are separated by a third medium(rather than vacuum), the situation becomes more complex. In aqueous solutions, the effects of dispersion forces between atoms/molecules are frequently less pronounced due to competition with polarizable solvent molecules. That is, the instantaneous fluctuations in one atom or molecule are felt both by the solvent (water) and by other molecules.

London forces become stronger as the atom in question becomes larger, and to a smaller degree for large molecules.[4] This is due to the increased polarizability of molecules with larger, more dispersed electron clouds. The polarizability is a measure of ease with which electrons can be redistributed; a large polarizability implies that the electrons are more easily redistributed. This trend is exemplified by the halogens (from smallest to largest: F2, Cl2, Br2, I2). The same increase of dispersive attraction occurs within and between organic molecules in the order RF<RCL<RBr<RI, or with other more polarizable heteroatoms.[5] Fluorine and chlorine are gases at room temperature, bromine is a liquid, and iodine is a solid.The London forces are thought to be arise from the motion of electrons.

## Quantum mechanical theory of dispersion forces

The first explanation of the attraction between noble gas atoms was given by Fritz London in 1930.[6][7][8] He used a quantum-mechanical theory based on second-order perturbation theory. The perturbation is because of the Coulomb interaction between the electrons and nuclei of the two moieties (atoms or molecules). The second-order perturbation expression of the interaction energy contains a sum over states. The states appearing in this sum are simple products of the stimulated electronic states of the monomers. Thus, no intermolecular antisymmetrization of the electronic states is included and the Pauli exclusion principle is only partially satisfied.

London wrote Taylor series expansion of the perturbation in ${\displaystyle {\frac {1}{R}}}$ , where ${\displaystyle R}$  is the distance between the nuclear centers of mass of the moieties.

This expansion is known as the multipole expansion because the terms in this series can be regarded as energies of two interacting multipoles, one on each monomer. Substitution of the multipole-expanded form of V into the second-order energy yields an expression that resembles somewhat an expression describing the interaction between instantaneous multipoles (see the qualitative description above). Additionally, an approximation, named after Albrecht Unsöld, must be introduced in order to obtain a description of London dispersion in terms of dipole polarizabilities and ionization potentials.

In this manner, the following approximation is obtained for the dispersion interaction ${\displaystyle E_{AB}^{\rm {disp}}}$  between two atoms ${\displaystyle A}$  and ${\displaystyle B}$ . Here ${\displaystyle \alpha _{A}}$  and ${\displaystyle \alpha _{B}}$  are the dipole polarizabilities of the respective atoms. The quantities ${\displaystyle I_{A}}$  and ${\displaystyle I_{B}}$  are the first ionization potentials of the atoms, and ${\displaystyle R}$  is the intermolecular distance.

${\displaystyle E_{AB}^{\rm {disp}}\approx -{3 \over 2}{I_{A}I_{B} \over I_{A}+I_{B}}{\alpha _{A}\alpha _{B} \over {R^{6}}}}$

Note that this final London equation does not contain instantaneous dipoles (see molecular dipoles). The "explanation" of the dispersion force as the interaction between two such dipoles was invented after London arrived at the proper quantum mechanical theory. The authoritative work[9] contains a criticism of the instantaneous dipole model[10] and a modern and thorough exposition of the theory of intermolecular forces.

The London theory has much similarity to the quantum mechanical theory of light dispersion, which is why London coined the phrase "dispersion effect". In physics, the term "dispersion" describes the variation of a quantity with frequency, which is the fluctuation of the electrons in the case of the London dispersion.

## Relative magnitude

Dispersion forces are usually dominant of the three van der Waals forces (orientation, induction, dispersion) between atoms and molecules, with the exception of molecules that are small and highly polar, such as water. The following contribution of the dispersion to the total intermolecular interaction energy has been given:[11]

Contribution of the dispersion to the total intermolecular interaction energy
Molecule pair % of the total energy of interaction
Ne-Ne 100
CH4-CH4 100
HCl-HCl 86
HBr-HBr 96
HI-HI 99
CH3Cl-CH3Cl 68
NH3-NH3 57
H2O-H2O 24
HCl-HI 96
H2O-CH4 87

## References

1. ^ "Chemguy Chemistry P5T8S9". YouTube. Retrieved 2013-04-01.
2. ^ Israelachvili, Jacob N. (2011), "Interactions of Biological Membranes and Structures", Intermolecular and Surface Forces, Elsevier, pp. 577–616, doi:10.1016/b978-0-12-375182-9.10021-1, ISBN 9780123751829
3. ^ Gelardi, G.; Flatt, R.J. (2016), "Working mechanisms of water reducers and superplasticizers", Science and Technology of Concrete Admixtures, Elsevier, pp. 257–278, doi:10.1016/b978-0-08-100693-1.00011-4, ISBN 9780081006931
4. ^ "London Dispersion Forces". Retrieved May 24, 2019.
5. ^ Schneider,Hans-Jörg Dispersive Interactions in Solution Complexes Dispersive Interactions in Solution Complexes Acc. Chem. Res 2015, 48 , 1815–1822.[1]
6. ^ R. Eisenschitz & F. London (1930), "Über das Verhältnis der van der Waalsschen Kräfte zu den homöopolaren Bindungskräften", Zeitschrift für Physik, 60 (7–8): 491–527, Bibcode:1930ZPhy...60..491E, doi:10.1007/BF01341258
7. ^ London, F. (1930), "Zur Theorie und Systematik der Molekularkräfte", Zeitschrift für Physik, 63 (3–4): 245, Bibcode:1930ZPhy...63..245L, doi:10.1007/BF01421741 and London, F. (1937), Zeitschrift für Physikalische Chemie, 33: 8–26 Missing or empty |title= (help). English translations in Parr, Robert G. (2000), H. Hettema (ed.), "Quantum Chemistry, Classic Scientific Papers", Physics Today, Singapore: World Scientific, 54 (6): 63, Bibcode:2001PhT....54f..63H, doi:10.1063/1.1387598
8. ^ F. London (1937), "The general theory of molecular forces", Transactions of the Faraday Society, 33: 8–26, doi:10.1039/tf937330008b
9. ^ J. O. Hirschfelder; C. F. Curtiss & R. B. Bird (1954), Molecular Theory of Gases and Liquids, New York: Wiley
10. ^ A. J. Stone (1996), The Theory of Intermolecular Forces, Oxford: Clarendon Press
11. ^ Jacob Israelachvili (1992), Intermolecular and Surface Forces (2nd ed.), Academic Press