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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.

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London dispersion forces (Instantaneous dipole/ induced dipole)

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

Dipole-dipole-interaction-in-HCl-2D.png

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

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:

Hydrogen-bonding-in-water-2D.png

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 (kcal)[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

See also

References

  1. ^ "London Dispersion Forces". http://www.chem.purdue.edu/gchelp/liquids/disperse.html. Retrieved 2009-09-20.  
  2. ^ Hydrogen Bonding.
  3. ^ Volland, Dr. Walt. ""Intermolecular" Forces". http://www.800mainstreet.com/08/0008-0012-interforce.html. Retrieved 2009-09-20.  
  4. ^ Organic Chemistry: Structure and Reactivity by Seyhan Ege, pp30-33, 67

External links

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Software for calculation of intermolecular forces


Simple English

In physics, chemistry, and biology, intermolecular forces are forces that act between stable molecules or between functional groups of macromolecules.

These forces are generally much weaker than the chemical bonding forces. Their bonding energies are less than a few kcal/mol. But they are responsible for many different physical, chemical, and biological phenomena. In general one distinguishes short and long range intermolecular forces.

==== Hydrogen bonding ====

Hydrogen bonding is an intermolecular interaction with a hydrogen atom being present in the intermolecular bond.

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