Covalent bonds are formed by the sharing of electrons between atoms, typically nonmetals, creating a strong link that holds the atoms together within a structure. The relationship between these bonds and a substance’s melting point is not straightforward; the answer depends entirely on how the covalently bonded atoms are arranged in the solid state. Some covalent substances melt at extremely low temperatures, while others possess some of the highest melting points known in chemistry. This wide range of properties highlights that the overall structure of the material is far more important than the mere existence of covalent bonds when determining the energy required for a phase transition.
Understanding the Energy Required for Melting
Melting, the transition from a solid to a liquid, requires an input of energy sufficient to overcome the forces holding the solid structure rigid. Covalent compounds involve two distinct types of attractive forces that must be considered when predicting a melting point. The first is the intramolecular force, which is the strong covalent bond existing within the molecule, holding the atoms together. These bonds are very strong, generally requiring a large amount of energy to break. The second is the intermolecular force, which is the much weaker attraction that exists between separate, distinct molecules in the solid state. Whether a covalent substance has a high or low melting point is determined by which of these two force types must be broken to achieve the liquid state.
Discrete Covalent Molecules and Low Melting Points
Many substances composed of covalent bonds exist as discrete, individual molecules, such as methane (\(\text{CH}_4\)) or carbon dioxide (\(\text{CO}_2\)). In the solid state, these separate molecules are not chemically bonded to one another; instead, they are simply packed close together and held in place by weak intermolecular forces. When heat is supplied to a solid composed of these discrete molecules, the energy is used to increase the kinetic energy of the molecules, causing them to move apart. Only the weak intermolecular forces are overcome during this process, allowing the molecules to separate and flow as a liquid while the strong covalent bonds within each molecule remain completely intact. Because only a small amount of energy is needed to break these weak attractions, these substances melt or sublime at very low temperatures, with \(\text{CO}_2\) subliming at \(-78.5^\circ \text{C}\) at atmospheric pressure.
Giant Covalent Networks and High Melting Points
In stark contrast to discrete molecules, giant covalent network solids are materials where every atom is bonded to its neighbors by a continuous network of strong covalent bonds. These structures are not composed of individual molecules; rather, the entire solid crystal acts as a single, immense molecule. Examples of these network solids include diamond, which is pure carbon, and silicon dioxide (\(\text{SiO}_2\)), the main component of quartz and sand.
To melt a giant covalent network solid, the energy supplied must be sufficient to break a vast number of these extremely strong covalent bonds throughout the entire three-dimensional lattice. Breaking these chemical bonds requires a massive energy input, which translates directly into extremely high melting points. For instance, diamond, where each carbon atom is tetrahedrally bonded to four others, has a melting point estimated to be around \(3550^\circ \text{C}\). Similarly, silicon dioxide forms a strong three-dimensional network, requiring temperatures near \(1700^\circ \text{C}\) to melt.