Interactions between Particles

Since chemical species in condensed phases interact strongly, solids and liquids are more complicated than gases, where intermolecular forces are usually negligible. The interactions of atoms, molecules, and ions are electrostatic in origin, but there are several types of interaction with quite different energies: Ion-ion interactions involve by far the largest energies. When like charges interact (a repulsive interaction), the energy is positive. When unlike charges interact (an attractive interaction), the energy is negative. These interactions are most important in ionic crystals.

Ion-dipole interactions are next in order of interaction energy. Such interactions are exemplified by the solvation of ions such as Na⁺ and Cl^- by water. Water has a large dipole moment and H₂O molecules are strongly attracted to Na+ with the negative end (O) of the water dipoles oriented toward the cation. Similarly Cl- is solvated with the positive end of the water dipole (H) oriented toward the anion.  Dipole-dipole interactions are responsible for the cohesion of dipolar liquids. For example, methyl ether, CH₃OCH₃, with a dipole moment, boils at -25°C, but propane, CH₃CH₂CH₃, with about the same molar mass but negligible dipole moment, boils at - 45°C.

Dipole–induced dipole interactions are also important. Nonpolar molecules tend to have their electron clouds attracted (or repelled) by a nearby dipolar molecule oriented with its positive (or negative) end toward the nonpolar species. This induced dipole interacts, on the average, with the dipolar molecule as if it were a permanent dipole.

This effect is responsible for the solubility of nonpolar gases such as O₂, N₂, or CO₂ in water. Induced dipole–induced dipole interactions, also called London forces, result when a nonpolar molecule undergoes a distortion, perhaps as a result of a collision, which results in a momentary separation of the centers of positive and negative charge. The resulting dipole induces a dipole in a neighboring molecule, resulting in an attractive interaction.

Both dipole–induced dipole and induced dipole-induced dipole interactions generally increase with number of electrons, particularly in atoms on the outer surface of the molecule. For example, the boiling points of the inert gases—He (4 K), Ne (27 K), Ar (88 K), Kr (120 K), and Xe (166 K)—increase with the number of polarizable electrons. UF₆ (bp 56°C) is more volatile than SbCl₅ (bp 79°C); uranium has more electrons than antimony, but the halogens are on the surface of the molecules, and there are more polarizable electrons in the five Cl atoms than in the six F atoms.  Dipole-dipole, dipole–induced dipole, and induced dipole–induced dipole interactions are often referred to collectively as van der Waals interactions.

Some dipolar molecules exhibit much greater intermolecular interactions than might have been expected from their dipole moments alone. This effect is illustrated by the data of Table 1.1.

Table 1.1. Boiling Points of Some Simple Hydrides

The heavier hydrides behave as expected, i.e., an increase in boiling point with number of polarizable electrons. However, NH₃, H₂O, and HF have anomalously high boiling points. This effect is due to hydrogen bonding, the tendency of hydrogen atoms bound to electronegative atoms to interact with unshared electron pairs on adjacent molecules.

Hydrogen bonding is expected for hydrides with a strongly electronegative central atom—N, O, and F—with unshared pairs of electrons which can interact with the positive hydrogen of an adjacent molecule. Water is by far the best former of hydrogen bonds since it has two unshared pairs and two H's. NH₃ (one unshared pair and three H's) and HF (three unshared pairs and one H) have only half the potential hydrogen bonds of H₂O. 

Hydrogen bonds are not limited to NH₃, H₂O, and HF. Molecules containing  NH₂, -OH, or -F groups are capable of hydrogen bonding. Hydrogen bonds are of great importance in many chemical and biochemical systems. The structure of ice depends on hydrogen bonds; the base pairing scheme of DNA vital to the genetic code relies on hydrogen bonds; the structures of proteins and many synthetic polymers are dictated by hydrogen bond formation.