(physical chemistry) The force between two molecules; it is that negative gradient of the potential energy between the interacting molecules, if energy is a function of the distance between the centers of the molecules.
intermolecular forces
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(′in·tər·mə′lek·yə·lər ′fōrs)
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| Sci-Tech Encyclopedia: Intermolecular forces |
Attractive or repulsive interactions that occur between all atoms and molecules. Intermolecular forces become significant at molecular separations of about 1 nanometer or less, but are much weaker than the forces associated with chemical bonding. They are important, however, because they are responsible for many of the physical properties of solids, liquids, and gases. These forces are also largely responsible for the three-dimensional arrangements of biological molecules and polymers.
Intermolecular forces can be classified into several types, of which two are universal. The attractive force known as dispersion arises from the quantum-mechanical fluctuation of the electron density around the nucleus of each atom. At distances greater than 1 nm or so, the electrons of each atom move independently of the other, and the charge distribution is spherically symmetric. At shorter distances, an instantaneous fluctuation of the charge density in one atom can affect the other. If the electrons of one atom move briefly to the side nearer the other, the electrons of the other atom are repelled to the far side. In this configuration, both atoms have a small dipole moment, and they attract each other electrostatically. At another moment, the electrons may move the other way, but their motions are correlated so that an attractive force is maintained on average. Molecular orbital theory shows that the electrons of each atom are slightly more likely to be on the side nearer to the other atom, so that each atomic nucleus is attracted by its own electrons in the direction of the other atom.
At small separations the electron clouds can overlap, and repulsive forces arise. These forces are described as exchange-repulsion, and are a consequence of the Pauli exclusion principle, a quantum-mechanical effect which prevents electrons from occupying the same region of space simultaneously. To accommodate it, electrons are squeezed out from the region between the nuclei, which repel each other as a result. Each element can be assigned, approximately, a characteristic van der Waals radius; that is, when atoms in different molecules approach more closely than the sum of their radii, the repulsion ennergy increases sharply. It is this effect that gives molecules their characteristic shape, leading to steric effects in chemical reactions. See also Steric effect (chemistry).
The other important source of intermolecular forces is the electrostatic interaction. When molecules are formed from atoms, electrons flow from electropositive atoms to electronegative ones, so that the atoms become somewhat positively or negatively charged. In addition, the charge distribution of each atom may be distorted by the process of bond formation, leading to atomic dipole and quadrupole moments. The electrostatic interaction between these is an important source of intermolecular forces, especially in polar molecules, but also in molecules that are not normally thought of as highly polar. The electrostatic field of a molecule may cause polarization of its neighbors, and this leads to a further induction contribution to the intermolecular interaction. An induction interaction can often polarize both molecules in such a way as to favor interactions with further molecules, leading to a cooperative network of intermolecular attractions. This effect is important in the network structure of water and ice. See also Water desalination.
Intermolecular forces are responsible for many of the bulk properties of matter in all its phases. A realistic description of the relationship between pressure, volume, and temperature of a gas must include the effects of attractive and repulsive forces between molecules. The viscosity, diffusion, and surface tension of liquids are examples of physical properties which depend strongly on intermolecular forces. Intermolecular forces are also responsible for the ordered arrangement of molecules in solids, and account for their elasticity and properties (such as the velocity of sound in materials).
| Columbia Encyclopedia: intermolecular forces |
intermolecular forces, forces that are exerted by molecules on each other and that, in general, affect the macroscopic properties of the material of which the molecules are a part. Such forces may be either attractive or repulsive in nature. They are conveniently divided into two classes: short-range forces, which operate when the centers of the molecules are separated by 3 angstroms or less, and long-range forces, which operate at greater distances. Generally, if molecules do not tend to interact chemically, the short-range forces between them are repulsive. These forces arise from interactions of the electrons associated with the molecules and are also known as exchange forces. Molecules that interact chemically have attractive exchange forces; these are also known as valence forces. Mechanical rigidity of molecules and effects such as limited compressibility of matter arise from repulsive exchange forces. Long-range forces, or van der Waals forces as they are also called, are attractive and account for a wide range of physical phenomena, such as friction, surface tension, adhesion and cohesion of liquids and solids, viscosity, and the discrepancies between the actual behavior of gases and that predicted by the ideal gas law. Van der Waals forces arise in a number of ways, one being the tendency of electrically polarized molecules to become aligned. Quantum theory indicates also that in some cases the electrostatic fields associated with electrons in neighboring molecules constrain the electrons to move more or less in phase.
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In physics, chemistry, and biology, intermolecular forces are forces that act between stable molecules or between functional groups of macromolecules. Intermolecular forces include momentary attractions between molecules, diatomic free elements, and individual atoms. These forces, most notably London Dispersion forces, dipole-dipole interactions and hydrogen bonding, are significantly weaker than either ionic or covalent bonding, but still have a noticeable chemical effect (see hydrogen bonding in water). Intermolecular forces are due to differences in charge density in molecules.
Contents |
London dispersion forces (Instantaneous dipole/ induced dipole)
Main article: London dispersion force
The London dispersion force (one of the three types of van der Waals forces) is caused by instantaneous changes in the dipole of atoms, caused by the location of the electrons in the atoms' orbitals. The probability of an electron in an atom is given by the Schrödinger equation. When an electron is on one side of the nucleus, this side becomes slightly negative (indicated by δ-); this in turn repels electrons in neighbouring atoms, making these regions slightly positive (δ+). This induced dipole causes a brief electrostatic attraction between the two molecules. The electron immediately moves to another point and the electrostatic attraction is broken.
London Dispersion forces are typically very weak (see the comparison below) because the attractions are so quickly broken, and the charges involved are so small.[1]
Dipole-dipole interactions
Dipole-Dipole interactions, also called Keesom interactions after Willem Hendrik Keesom, are caused by permanent dipoles in molecules. When one atom is covalently bonded to another with a significantly different electronegativity, the electronegative atom draws the electrons in the bond nearer to itself, becoming slightly negative. Conversely, the other atom becomes slightly positive. Electrostatic forces are generated between the opposing charges and the molecules align themselves to increase the attraction (reducing potential energy).
An example of dipole-dipole interactions can be seen in hydrochloric acid
This is not an example of hydrogen bonding (see below) because the chlorine atom is not electronegative enough.
Note that almost always the dipole-dipole interaction between two atoms is zero, because atoms rarely carry a permanent dipole, see atomic dipoles.
Often, molecules can have dipoles within them, but have no overall dipole moment. This occurs if there is symmetry within the molecule, causing each of the dipoles to cancel each other out. This occurs in molecules such as tetrachloromethane.
Hydrogen bonding
Main article: Hydrogen bond
Hydrogen bonds are a stronger form of dipole-dipole interactions, caused by highly electronegative atoms. They only occur between hydrogen and oxygen, fluorine or nitrogen,[2] and are the strongest intermolecular force. The high electronegativities of F, O and N create highly polar bonds with hydrogen, which leads to strong bonding between hydrogen atoms on one molecule and the lone pairs of F, O or N atoms on adjacent molecules. The high boiling point of water is an effect of the extensive hydrogen bonding between the molecules:
For quite some time it was believed that hydrogen bonding required an explanation that was different from the other intermolecular interactions. However, reliable computer calculations that became possible during the 1980s have shown that only the four effects listed above play a role, with the dipole-dipole interaction being particularly important. Since the four effects account completely for the bonding in small dimers like the water dimer, for which highly accurate calculations are feasible, it is now generally believed that no other bonding effects are operative.
Hydrogen bonds are found throughout nature. In water the dynamics of these bonds produce unique properties essential to all known life-forms. Hydrogen bonds, between hydrogen atoms and nitrogen atoms, of adjacent DNA base pairs generate intermolecular forces that improve binding between the strands of the molecule. Hydrophobic effects between the double-stranded DNA and the solute nucleoplasm prevail in sustaining the double-helix structure of DNA.
Relative strength of forces
| Bond type | Dissociation energy (kJ)[3],[4] |
|---|---|
| Covalent | 400 |
| Hydrogen bonds | 12-16 |
| Dipole-dipole | 2.0 - 0.5 |
| 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
 |
This section requires expansion. |
See also
- Coomber's relationship
- Hydrophobic effect
- Intramolecular force
- Molecular solid
- Polymer
- Quantum chemistry computer programs
- Software for molecular mechanics modeling
References
- ^ "London Dispersion Forces". http://www.chem.purdue.edu/gchelp/liquids/disperse.html. Retrieved 2009-09-20.
- ^ Hydrogen Bonding.
- ^ Volland, Dr. Walt. ""Intermolecular" Forces". http://www.800mainstreet.com/08/0008-0012-interforce.html. Retrieved 2009-09-20.
- ^ Organic Chemistry: Structure and Reactivity by Seyhan Ege, pp30-33, 67
External links
Software for calculation of intermolecular forces
- Quantum 3.2
- SAPT: An ab initio quantumchemical package.
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