The Nature of Molecular Bonds
Recall that a molecule is defined as a discrete aggregate of atoms bound together sufficiently tightly by directed covalent forces to allow it to retain its individuality when the substance is dissolved, melted, or vaporized. The two words italicized in the preceding sentence are important. Covalent bonding implies that the forces acting between atoms within the molecule are much stronger than those acting between molecules; the directional property of covalent bonding confers on each molecule a distinctive shape which affects a number of its properties.
Liquids and solids composed of molecules are held together by van der Waals forces, and many of their properties reflect this weak binding. Molecular solids tend to be soft or deformable, have low melting points, and are often sufficiently volatile to evaporate directly into the gas phase. This latter property often gives such solids a distinctive odor. Whereas the characteristic melting point of metals and ionic solids is ~1000 °C, most molecular solids melt well below ~300 °C. Thus, many corresponding substances are either liquid (ice) or gaseous (oxygen) at room temperature.
Molecular solids also have relatively low density and hardness. The elements involved are light, and the intermolecular bonds are relatively long and are thus weak. Because of the charge neutrality of the constituent molecules, and because of the long distance between them, molecular solids are electrical insulators.
Because dispersion forces and the other van der Waals forces increase with the number of atoms, larger molecules are generally less volatile, and have higher melting points, than do the smaller ones. Also, as one moves down a column in the periodic table, the outer electrons are more loosely bound to the nucleus, increasing the polarisability of the atom and thus its susceptibility to van der Waals-type interactions. This effect is particularly apparent in the progression of the boiling points of the successively heavier noble gas elements.
The term "molecular solid" may refer not to a certain chemical composition, but to a specific form of a material. For example, solid phosphorus can crystallize in different allotropes called "white", "red" and "black" phosphorus. White phosphorus forms molecular crystals composed of tetrahedral P4 molecules. A molecular solid, white phosphorus has a relatively low density of 1.82 g/cm3 and melting point of 44.1 °C; it is a soft material which can be cut with a knife. Heating at ambient pressure to 250 °C or exposing to sunlight converts white phosphorus to red phosphorus, in which the P4 tetrahedra are no longer isolated, but are connected by covalent bonds into polymer-like chains. Heating white phosphorus under high (GPa) pressures converts it to black phosphorus, which has a layered, graphite-like structure. When white phosphorus is converted to the covalent red phosphorus, the density increases to 2.2–2.4 g/cm3 and melting point to 590 °C; when white phosphorus is transformed into the (also covalent) black phosphorus, the density becomes 2.69–3.8 g/cm3 and melting temperature ~200 °C.
Both red and black phosphorus forms are significantly harder than white phosphorus, and whereas white phosphorus is an insulator, the black allotrope, which consists of layers extending over the whole crystal, does conduct electricity. The structural transitions in phosphorus are reversible: upon releasing high pressure, black phosphorus gradually converts into the red allotrope, and by vaporizing red phosphorus at 490 °C in inert atmosphere and condensing the vapor, covalent red phosphorus can be transformed back into the white molecular solid. Similarly, yellow arsenic is a molecular solid composed of As4 units; it is metastable and gradually transforms into gray arsenic upon heating or illumination. Some forms of sulfur and selenium are composed of S8 (or Se8) units and are molecular solids at ambient conditions, but they can convert into covalent allotropes having atomic chains extending all through the crystal.
Changes in the chemical composition can have even stronger effects on the bonding in solids. For example, whereas both hydrogen and lithium belong to the first group of the periodic table, LiCl is ionic and HCl is a molecular solid. Several classes of molecular solids can be distinguished. The vast majority of molecular solids can be attributed to organic compounds containing carbon and hydrogen, such as hydrocarbons (CnHm). Spherical molecules consisting of different number of carbon atoms, that is fullerenes, are another important class. Less numerous, yet distinctive molecular solids are halogens (e.g. Cl2) and their compounds with hydrogen (HCl), as well as light chalcogens (O2) and pnictogens (N2). Conductivity of molecular solids can be illustrated on example of fullerene. Its solid is an insulator because all valence electrons of carbon atoms are involved into the covalent bonds within the individual carbon molecules. However, inserting (intercalating) alkali metal atoms between the fullerene molecules provides extra electrons, which can be easily ionized from the metal atoms and make material conductive and even superconductive. (Figure 2)