Intermolecular force

Intermolecular forces are electromagnetic forces which act between molecules or between widely separated regions of a macromolecule. Listed in order of decreasing strength, these forces are:

Description and strength

These are fundamentally electrostatic interactions (ionic interactions, hydrogen bond, dipole-dipole interactions) or electrodynamic interactions (van der Waals/London forces). Electrostatic interactions are classically described by Coulomb's law; the basic difference between them is the strength of their charge. Ionic interactions are the strongest with integer level charges, hydrogen bonds have partial charges that are about an order of magnitude weaker, and dipole-dipole interactions also come from partial charges another order of magnitude weaker.

A very approximate strength order would be:
Bond type	Relative strength
Ionic bonds	1000
Hydrogen bonds	100
Dipole-dipole	10
London Forces	1

Ionic interactions

These are interactions that occur between charged species (ions). Like charges repel, while opposite charges attract. These bonds form when the electronegativities between two atoms are large enough that one steals an electron from the other. The now oppositely charged ions are attracted. Ionic compounds have high melting and boiling points due to the large amount of heat required to break the forces between the charged ions. When molten they are also good conductors of heat and electricity, due to free or delocalised electrons.

Hydrogen bonding

Hydrogen bonding occurs when a hydrogen atom is covalently bonded to a small highly electronegative atom such as nitrogen, oxygen, or fluorine. The result is a dipolar molecule. The hydrogen atom has a partial positive charge δ+ (see unit: Debye) and can interact with another highly electronegative atom in an adjacent molecule (again N, O, or F). This results in a stabilizing interaction that binds the two molecules together. An important example is water: δ+ δ- δ+ H O - H \ / O | | | H / δ- δ+ H δ+

Hydrogen bonds are found throughout nature. They give water its unique properties that are so important to life on earth. Hydrogen bonds between hydrogen atoms and nitrogen atoms of adjacent base pairs provide the intermolecular force that bind together the two strands in a molecule of DNA.

The critical difference between hydrogen bonding and dipole-dipole interactions is that the hydrogen is partially transferred to the second molecule - the second molecule's lone pair of electrons forms a covalent bond and the pair becomes somewhat like:

H₂O⁺-H ^-O-H

The effect is twofold: The bonding is stronger and is directional. The directional nature of hydrogen bonding requires the two molecules to adopt a specific relative geometry.

Dipole-dipole interactions

Dipole-dipole interactions, also called Keesom interactions after Willem Hendrik Keesom who produced the first mathematical description in 1921, are the forces that occur between two molecules with permanent dipoles (spatially oriented δ+ within a molecule). These work in a similar manner to ionic interactions, but are weaker because only partial charges are involved. An example of this can be seen in hydrochloric acid:

(+)(-) (+)(-) H--Cl

H--Cl

London dispersion forces

Also called London forces, instantaneous dipole (or multipole) effects (spatially variable δ+) or Van der Waals forces, these involve the attraction between temporarily induced dipoles in nonpolar molecules (often disappear within an instant). This polarization can be induced either by a polar molecule or by the repulsion of negatively charged electron clouds in nonpolar molecules. An example of the former is chlorine dissolving in water:

(+)(-)(+) (-) (+) Dipole H-O-H

-Cl-Cl Dipole

An example of the second scenario is found in molecular chlorine:

(+) (-) (+) (-) Dipole Cl-Cl

--Cl-Cl Dipole

London Dispersion forces exist between all atoms. London forces are the only reason for rare-gas atoms to condense at low temperature.

Software for calculation of intermolecular forces

Quantum 3.2