Covalent compounds are formed when atoms share electrons, creating distinct, electrically neutral units called molecules, such as water, sugar, and carbon dioxide. Unlike substances like table salt or metals, which possess high melting temperatures, solid covalent compounds melt at relatively low temperatures. This physical characteristic is a direct consequence of how these compounds organize themselves in the solid state. The low energy required to transition from a solid to a liquid points to a fundamental weakness in the attractive forces holding the molecules together.
The Structure of Molecular Solids
Solid covalent compounds exist as molecular solids. This structure is defined by an arrangement of individual, discrete molecules packed closely together in a fixed, repeating pattern. For instance, solid water (ice) is composed of separate H₂O molecules, and dry ice is made of individual CO₂ molecules. These molecules maintain their identity and internal structure even when compressed into a solid form.
This arrangement is markedly different from the continuous, three-dimensional network found in other types of solids, like diamond or quartz. In those solids, all the atoms are linked by strong covalent bonds, essentially forming one giant molecule. In a molecular solid, the atoms within each molecule are bonded strongly, but the attraction between one molecule and its neighboring molecule is comparatively weak.
The Difference Between Intermolecular and Intramolecular Forces
The key to understanding the low melting point of these solids lies in recognizing the two distinct types of forces at play. The first type is the intramolecular force, which refers to the strong covalent bonds that exist within each individual molecule. These bonds involve the sharing of electrons and determine the chemical identity and structure of the molecule, requiring a large amount of energy to break.
The second type is the intermolecular force (IMF), which represents the much weaker attractive forces between separate molecules. When a covalent compound is in its solid state, the molecules are held in place solely by these IMFs. These forces are physical attractions, not chemical bonds, and are generally electrostatic, arising from the interaction between positive and negative regions of adjacent molecules.
Intermolecular forces include London Dispersion Forces and Dipole-Dipole interactions. These forces are generally electrostatic, arising from the interaction between positive and negative regions of adjacent molecules. Regardless of the specific type, these intermolecular forces are significantly weaker than the covalent bonds holding the atoms together inside the molecule. For example, breaking the strong covalent bonds in water requires about 927 kilojoules of energy per mole, while overcoming the weak attractions between water molecules requires only about 41 kilojoules per mole.
Why Little Energy is Needed for Phase Change
Melting is a physical process requiring thermal energy to allow molecules to overcome the forces holding them in a rigid, fixed lattice structure. For a molecular solid to melt, heat energy only needs to disrupt the weak intermolecular forces that exist between the discrete molecules. The heat does not need to be high enough to break the strong covalent bonds holding the atoms within each molecule together.
Because the attractive forces between molecules are so weak, only a small input of heat is necessary to give the molecules enough kinetic energy to move past one another freely, defining a liquid. The molecules themselves remain intact throughout the entire phase change. This explains why molecular solids like ice melt at 0°C, or why wax softens easily when exposed to a small flame.
The low energy threshold to disrupt the weak IMFs results in the low melting points observed for this class of compounds. Common molecular solids such as sucrose (186°C) or naphthalene (80°C) demonstrate this principle. These temperatures are far lower than the melting point of an ionic solid like table salt, which requires 801°C to break its much stronger electrostatic network.