A solid is a state of matter defined by a fixed shape and a definite volume, resulting from the strong forces that hold its constituent particles in place. Scientists classify these crystalline solids into four main categories based on the specific type of attractive force that binds the atoms, ions, or molecules together. These forces determine the material’s physical properties, such as its hardness, melting temperature, and ability to conduct electricity. The four distinct types of crystalline solids are ionic, covalent network, molecular, and metallic solids.
Ionic Solids
Ionic solids are structured assemblies of positively and negatively charged ions held together by electrostatic attractions, known as ionic bonds. These forces are non-directional and extend equally in all directions, causing the ions to arrange themselves into a highly organized, repeating three-dimensional pattern called a crystal lattice. This arrangement creates a structure that is extremely stable and difficult to break apart.
Ionic solids exhibit very high melting points, often well above 800°C, with some compounds like magnesium oxide melting at nearly 2800°C. While the structure is rigid and hard, the material is also brittle. If the crystal lattice is struck, a slight shift can bring ions of the same charge close together, creating a strong repulsive force that causes the solid to fracture immediately.
In their solid state, ionic compounds are poor conductors of electricity because the charged ions are locked into fixed positions within the lattice and cannot move freely. However, if the solid is melted or dissolved in water, the ions become mobile and are able to carry an electrical charge. Common examples of this class of solid include sodium chloride (table salt) and calcium fluoride.
Covalent Network Solids
Covalent network solids, sometimes called atomic solids, are distinguished by a continuous, extended network of covalent bonds linking all the atoms throughout the entire sample. These covalent bonds are exceptionally strong and highly directional, creating a rigid and extensive structure.
To melt or break a covalent network solid, these numerous covalent bonds must be broken, a process requiring vast amounts of energy. This results in the highest melting points of all solid types; for example, diamond, a form of pure carbon, melts at over 3500°C. Their extensive bonding also makes them extremely hard and insoluble in most solvents.
Diamond serves as the classic example, where each carbon atom is bonded tetrahedrally to four neighbors, creating one of the hardest substances known. Another familiar example is silicon dioxide, the primary component of quartz and sand, which melts at approximately 1650°C. The electrons in these solids are localized between the bonded atoms, making most covalent network solids poor electrical conductors.
Molecular Solids
Molecular solids are composed of discrete, individual molecules held together by relatively weak intermolecular forces, rather than strong chemical bonds. These weak attractions include London dispersion forces, dipole-dipole interactions, and hydrogen bonds.
Molecular solids are typically soft and possess very low melting points, often melting well below 300°C. The energy required to overcome these weak intermolecular forces is minimal, which is why substances like solid water (ice) melt at 0°C, and solid carbon dioxide (dry ice) sublimes, or turns directly into a gas, at -78.5°C.
Because the valence electrons are tightly held within the individual molecules and are not free to move throughout the structure, molecular solids are poor conductors of electricity. The melting point for these solids often correlates with the size of the molecule, as larger molecules have stronger London dispersion forces. Examples include ice, solid CO2, and sugar (sucrose).
Metallic Solids
Metallic solids consist of a lattice of positive metal ions surrounded by a “sea” of delocalized valence electrons. These valence electrons are free to move throughout the entire crystal. This collective sharing of electrons is known as metallic bonding.
The free movement of electrons allows metallic solids to be excellent conductors of both heat and electricity. Furthermore, the non-directional nature of the metallic bond allows the layers of positive ions to slide past one another without fracturing the material.
This structural flexibility makes metals both malleable, meaning they can be hammered into thin sheets, and ductile, meaning they can be drawn into wires. The strength of the metallic bond varies widely, leading to a broad range of melting points, from mercury being a liquid at room temperature to tungsten melting near 3400°C. Typical examples of metallic solids include copper, gold, iron, and aluminum.