Intermolecular forces are attractive forces that exist between molecules. They dictate several bulk properties, such as melting points, boiling points, and solubilities (miscibilities) of substances. For example, a high-boiling-point liquid, like water (H2O, b.p. 100 °C), exhibits stronger intermolecular forces compared to a low-boiling-point liquid, like hexane (C6H14, b.p. 68.73 °C). The three kinds of intermolecular interactions include i) ion–dipole forces, ii) dipole–dipole interactions, and iii) van der Waals forces, which include London dispersion forces.
Ion–dipole forces are the electrostatic attractions between an ion and a dipole. They are common in solutions and play an important role in the dissolution of ionic compounds, like KCl, in water. The strength of ion–dipole interactions is directly proportional to i) the charge on the ion and ii) the magnitude of the dipole of polar molecules.
Polar molecules have a partial positive charge on one end and a partial negative charge on the other end of the molecule—a separation of charge called a dipole. The attractive force between two permanent dipoles is called a dipole–dipole attraction—the electrostatic force between the partially positive end of one polar molecule and the partially negative end of another. Hydrogen bonding is a type of dipole–dipole interaction between molecules with hydrogen, bonded to a highly electronegative atom, such as O, N, or F. The resulting partially positively charged H atom on one molecule (the hydrogen bond donor) could interact strongly with a lone pair of electrons of a partially negatively charged O, N, or F atom on adjacent molecules (the hydrogen bond acceptor). Hydrogen bonding increases the boiling point considerably.
The weakest of all forces is the van der Waals forces, which depend on the intermolecular distances between atoms and molecules. London dispersion forces, a subset of van der Waals forces, are experienced as a result of interactions between uncharged atoms/molecules owing to temporary, spontaneous shifts in electron distribution. The strength of these forces appears to increase with increasing molecular weight owing to the increase in surface area. As a result, compounds of higher molecular weights will generally boil at higher temperatures. Of note is that a branched hydrocarbon (neopentane) normally has a smaller surface area than its respective straight-chain (n-pentane) isomer, and therefore, a lower boiling point.
Liquids that can be homogeneously mixed in any proportion are said to be miscible. Miscible liquids have similar polarities. For example, methanol and water are both polar and capable of hydrogen bonding. On mixing, methanol and water interact through intermolecular hydrogen bonds of comparable strength to the methanol–methanol, and water–water interactions; thus, they are miscible. Likewise, nonpolar liquids like hexane and bromine are miscible with each other through dispersion forces. The chemical axiom “like dissolves like” is useful to predict the miscibility of compounds. Two liquids that do not mix to an appreciable extent are called immiscible. For example, nonpolar hexane is immiscible in polar water. Relatively weak attractive forces between the hexane and water do not adequately overcome the stronger hydrogen bonding forces between water molecules.
This text is adapted from Openstax, Chemistry 2e, Section 10.1: Intermolecular Forces, Section 11.3: Solubility, and Chapter 10: Liquids and Solids.
