Covalent network solids (CNS) are materials defined by a vast, continuous network of atoms held together by strong covalent bonds. Unlike molecular solids, which are composed of individual molecules, a CNS is essentially one giant macromolecule. Electrical conductivity is determined by whether this extensive bonding framework allows for the movement of charge carriers. Most covalent network solids function as electrical insulators due to the rigid nature of their atomic structure and the localization of their valence electrons.
Defining Covalent Network Solids
Covalent network solids are characterized by atoms covalently bonded throughout the entire crystal, creating an extended three-dimensional lattice. The strength of these bonds gives the materials distinctive properties, such as extreme hardness and high melting points. Since the bonding is continuous and uniform, there are no distinct, separable molecules within the structure. The chemical formula for a CNS, such as silicon dioxide (\(\text{SiO}_2\)) or diamond (C), represents only the simple ratio of atoms. These strong, directional covalent bonds form the backbone of the solid.
The Mechanism of Non-Conductivity
The non-conductivity of most covalent network solids lies in the complete utilization and localization of valence electrons. Electrical conduction requires mobile charge carriers, typically delocalized electrons or free ions. In common CNS examples, such as diamond or quartz, every valence electron is tightly held within a specific covalent bond between two atoms.
For instance, in diamond, each carbon atom forms four single covalent bonds with its neighbors in a tetrahedral arrangement. This \(\text{sp}^3\) hybridization means all four valence electrons are fully involved in forming strong sigma (\(\sigma\)) bonds. There are no delocalized electrons available to move freely under the influence of an applied electric field. The energy gap between the filled valence band and the empty conduction band is too large for electrons to conduct current.
The Notable Exception of Graphite
The most prominent exception to the rule of non-conductivity is graphite, which demonstrates excellent electrical conductivity. This difference is due to a variation in its atomic structure and bonding arrangement compared to materials like diamond. Graphite consists of layers of carbon atoms arranged in hexagonal rings, where each carbon atom is covalently bonded to only three neighbors.
This leaves one valence electron per carbon atom that is not utilized in the primary sigma bond framework. This fourth electron occupies a p-orbital perpendicular to the plane of the carbon layers. These p-orbitals overlap across the entire layer, forming a delocalized pi (\(\pi\)) electron system. These delocalized electrons are free to move within the two-dimensional sheets, allowing graphite to conduct electricity parallel to its layers. Conduction perpendicular to the layers is poor because the sheets are only held together by weak van der Waals forces.