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Dipole-dipole interaction

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Dipole-dipole interaction

Intermolecular forces are forces of attraction or repulsion which act between neighboring particles (atoms, molecules or ions). They are weak compared to the intramolecular forces, the forces which keep a molecule together. For example, the covalent bond present within HCl molecules is much stronger than the forces present between the neighboring molecules, which exist when the molecules are sufficiently close to each other.

The first reference to the nature of microscopic forces is found in Alexis Clairault's work Theorie de la Figure de la Terre.[1] Other scientists who have contributed to the investigation of microscopic forces include Laplace, Gauss, Maxwell and Boltzmann.

Attractive intermolecular forces:

Dipole-dipole interactions

Dipole-dipole interactions are electrostatic interactions between permanent dipoles in molecules. These interactions tend to align the molecules to increase attraction (reducing potential energy). An example of a dipole-dipole interaction can be seen in hydrogen chloride (HCl): the positive end of a polar molecule will attract the negative end of the other molecule and influence its position. Polar molecules have a net attraction between them. Examples of polar molecules include hydrogen chloride (HCl) and chloroform (CHCl3).

Often molecules contain dipolar groups, but have no overall dipole moment. This occurs if there is symmetry within the molecule that causes the dipoles to cancel each other out. This occurs in molecules such as tetrachloromethane. Note that the dipole-dipole interaction between two individual atoms is usually zero, since atoms rarely carry a permanent dipole. See atomic dipoles.

Ion-dipole and ion-induced dipole forces

Ion-dipole and ion-induced dipole forces are similar to dipole-dipole and induced-dipole interactions but involve ions, instead of only polar and non-polar molecules. Ion-dipole and ion-induced dipole forces are stronger than dipole-dipole interactions because the charge of any ion is much greater than the charge of a dipole moment. Ion-dipole bonding is stronger than hydrogen bonding.

An ion-dipole force consists of an ion and a polar molecule interacting. They align so that the positive and negative groups are next to one another, allowing for maximum attraction.

An ion-induced dipole force consists of an ion and a non-polar molecule interacting. Like a dipole-induced dipole force, the charge of the ion causes distortion of the electron cloud on the non-polar molecule.[2]

Hydrogen bonding

Main article: Hydrogen bond

A hydrogen bond is the attraction between the lone pair of an electronegative atom and a hydrogen atom that is bonded to either nitrogen, oxygen, or fluorine.[3] The hydrogen bond is often described as a strong electrostatic dipole-dipole interaction. However, it also has some features of covalent bonding: it is directional, stronger than a van der Waals interaction, produces interatomic distances shorter than the sum of van der Waals radius, and usually involves a limited number of interaction partners, which can be interpreted as a kind of valence.

Intermolecular hydrogen bonding is responsible for the high boiling point of water (100 °C) compared to the other group 16 hydrides, which have no hydrogen bonds. Intramolecular hydrogen bonding is partly responsible for the secondary, tertiary, and quaternary structures of proteins and nucleic acids. It also plays an important role in the structure of polymers, both synthetic and natural.

Van der Waals forces

Main article: Van der Waals force

Keesom (permanent dipole) force

Keesom interactions (named after Willem Hendrik Keesom) are attractive interactions of dipoles that are Boltzmann-averaged over different rotational orientations of the dipoles. It is assumed that the molecules are constantly rotating and never get locked into place. This is a good assumption, but at some point molecules do get locked into place. The energy of a Keesom interaction depends on the inverse sixth power of the distance, unlike the interaction energy of two spatially fixed dipoles, which depends on the inverse third power of the distance. The Keesom interaction can only occur among molecules that possess permanent dipole moments aka two polar molecules. Also Keesom interactions are very weak Van der Waals interactions and do not occur in aqueous solutions that contain electrolytes. The angle averaged interaction is given by the following equation:

\frac{-2m_1^2m_2^2}{48\pi^2\varepsilon_o^2\varepsilon_r^2k_bTr^6}=V

Where m = charge per length, \varepsilon_o = permitivity of free space, \varepsilon_r = dielectric constant of surrounding material, T = temperature, k_b = Boltzmann constant, and r = distance between molecules.

Debye (induced dipole) force

The induced dipole forces appear from the induction (also known as polarization), which is the attractive interaction between a permanent multipole on one molecule with an induced (by the former di/multi-pole) multipole on another.[4][5][6][7] This interaction is called the Debye force, named after Peter J.W. Debye.

One example of an induction-interaction between permanent dipole and induced dipole is the interaction between HCl and Ar. In this system, Ar experiences a dipole as its electrons are attracted (to the H side of HCl) or repelled (from the Cl side) by HCl.[4][6] The angle averaged interaction is given by the following equation.

\frac{-m_1^2\alpha_2}{16\pi^2\varepsilon_o^2\varepsilon_r^2r^6}=V

Where \alpha= polarizability

This kind of interaction can be expected between any polar molecule and non-polar/symmetrical molecule. The induction-interaction force is far weaker than dipole-dipole interaction, but stronger than the London dispersion force.

London dispersion force

Otherwise known as quantum-induced instantaneous polarization or instantaneous dipole-induced dipole force, the London dispersion force is caused by correlated movements of the electrons in interacting molecules. Electrons that belong to different molecules start "fleeing" and avoiding each other at the short intermolecular distances, which is frequently described as formation of "instantaneous dipoles" that attract each other.

Relative strength of forces

Bond type Dissociation energy (kcal/mol),[8][9]
Ionic 400
Hydrogen bonds 12–16
Dipole–dipole 0.5–2
London (van der Waals) Forces <1

Note: this comparison is only approximate – the actual relative strengths will vary depending on the molecules involved.

Quantum mechanical theories

Main article: Quantum mechanical explanation of intermolecular interactions

Intermolecular forces observed between atoms and molecules can be described phenomenologically as occurring between permanent and instantaneous dipoles, as outlined above. Alternatively, one may seek a fundamental, unifying theory that is able to explain the various types of interactions such as hydrogen bonding, van der Waals forces and dipole-dipole interactions. Typically, this is done by applying the ideas of quantum mechanics to molecules, and Rayleigh–Schrödinger perturbation theory has been especially effective in this regard. When applied to existing quantum chemistry methods, such a quantum mechanical explanation of intermolecular interactions, this provides an array of approximate methods that can be used to analyze intermolecular interactions.

See also

References

External links

Software for calculation of intermolecular forces
  • Quantum 3.2
  • SAPT: An ab initio quantumchemical package.
ar:ترابط بين جزئي خارجي

zh:分子间作用力

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