Intermolecular Forces

Ionic and covalent bonds hold the atoms within a molecule together, and they are strong — typically hundreds of kilojoules per mole. Intermolecular forces, by contrast, are the much weaker attractive forces that act between separate, complete molecules (or, for noble gases, between separate atoms), typically only a few to a few tens of kilojoules per mole. They are collectively often called van der Waals forces, and although weak individually, they are entirely responsible for determining whether a substance is a gas, liquid or solid at a given temperature, and for trends in melting point, boiling point and viscosity across a series of related compounds.

Several distinct types of intermolecular force are recognised, in roughly increasing order of the polarity of the species involved:

• London dispersion forces operate between all molecules and atoms, polar or non-polar alike, and are the only intermolecular force available to non-polar species. They arise because the electron cloud of any atom or molecule is, at any given instant, never perfectly symmetric — small, momentary ("instantaneous") fluctuations in electron distribution create a temporary, fleeting dipole, which in turn induces a matching temporary dipole in a neighbouring molecule, producing a brief mutual attraction. Individually these attractions are extremely weak and short-lived, but they act continuously and, summed over a whole substance, are far from negligible. Dispersion forces grow stronger with increasing molecular size and the number of electrons present, because a larger, more diffuse electron cloud is more easily distorted (more polarisable). Liquid argon and gaseous methane, $\text{CH}_4$, being non-polar, are held together (and condense at all) by dispersion forces alone.
• Dipole–dipole forces operate between polar molecules — molecules that possess a permanent bond dipole (or net molecular dipole) even without any external influence. The partially positive end of one polar molecule is attracted to the partially negative end of a neighbouring polar molecule. Hydrogen chloride gas, $\text{HCl}$, is a polar molecule (chlorine being considerably more electronegative than hydrogen), so in addition to the dispersion forces present in every substance, $\text{HCl}$ molecules also experience dipole–dipole attraction to one another — a good example of a real substance experiencing more than one type of intermolecular force simultaneously. …