Covalent solids

• Type of physical science: Chemistry
• Field of study: Chemistry of solids

The different types of attractive forces that hold crystals together can be used to classify solids. Those substances categorized as covalent solids are held together by shared pairs of electrons and exhibit characteristics that make them desirable materials in the manufacture of a variety of articles, including properties such as hardness, resistance to high temperature, and lack of reactivity.

Overview

A covalent solid is a rigid, concentrated form of matter in which the dominant cohesive forces arise from the sharing of pairs of electrons between atoms. The solid state of matter is distinguished from the liquid and gaseous states by the tendency of a solid to retain its shape and volume independently of the container in which it is held. Though it is true that some solids do show changes of shape on prolonged standing—glasses, for example—such solids are often thought of as undercooled liquids; they are liquids in which cooling results in an extremely high viscosity or resistance to flow. The characteristic rigidity of a solid is a manifestation of restricted motion at the molecular level: In a solid, the attractive forces are strong enough to overcome the random thermal motions of the atoms, molecules, or ions that make up the solid, strong enough to hold those units in a relatively fixed location. To be sure, molecular motion persists in the solid state, but this motion consists for the most part of vibration of units around a mean location rather than translation through space.

In a crystalline solid, the spatial points around which structural units vibrate form a regular, three-dimensional, repeated pattern known as a lattice. The points themselves are referred to as “lattice sites” and can be occupied by atoms, molecules, or ions. In a crystal, the attractive forces that operate between units at different lattice sites hold the solid together.

Because these attractive forces range from very weak to extremely strong, different solids exhibit different tendencies to break apart, melt, or dissolve. In fact, many of the observable physical properties of a solid can be explained rather easily in terms of the forces that exist at the molecular level, and one commonly used classification scheme is based on identifying the units that occupy lattice sites and the forces that hold them in place.

In this classification scheme, crystalline solids are identified as molecular, ionic, metallic, or covalent. In a molecular solid, entire molecules occupy lattice sites, and the forces of attraction include one or more of a variety of rather weak forces collectively referred to as van der Waals forces. Ionic solids are composed of positive and negative ions occupying different lattice sites of the crystal. The mutual force of attraction that exists between oppositely charged entities holds ionic crystals together. In a metal, an array of positively charged ions—each ion an individual metal atom stripped of its outermost electrons—is immersed in a mobile electron gas.

The presence of this negatively charged gas penetrating the array of positive atomic cores gives rise to a cohesive force referred to as a metallic bond. Finally, in a covalent solid, the lattice sites are occupied by atoms, and these atoms are held in place by rather strong forces of attraction known as covalent bonds.

Covalent chemical bonds are forces of attraction that result from the sharing of pairs of electrons between atoms. Although covalent bonds range in stability, they are often stronger than metallic or ionic bonds and are always much stronger than van der Waals forces. They also exhibit directional properties; that is, bonding from some central atom will occur only in certain allowed directions. Thus, carbon commonly bonds to four atoms to form a three-dimensional structure known as a tetrahedron. Similarly, silicon and other members of the carbon family form tetrahedral molecules. When carbon bonds to three atoms, the resulting structure is flat, a triangle with carbon at the center. In general, the structure adopted around some central atom will depend upon the atom and the number of other atoms to which it is joined. Though shared by some types of van der Waals bonds, the directional properties of covalent bonds are not found in metallic bonds or ionic bonds.

Because these bonds are strong, attractive forces, covalent solids exhibit properties that reflect great cohesion in the solid state. Such solids tend to have very high melting points because a great amount of thermal energy must be deposited in the solid to break chemical bonds and free atoms from their lattice sites. These solids tend to be insoluble in common solvents for similar reasons. Any energy that might result from solvent-solute association is far less than that needed to free solute atoms from the crystalline lattice. Covalent solids tend to be hard materials because of the strength of the crystalline lattice. Finally, most of these solids are electrical insulators because their atoms do not have electrons available to carry electrical current; in most covalent solids, electrons are tied up in forming chemical bonds.

Because covalent bonds exhibit directional properties, certain geometries are forced on the lattices of covalent solids. Thus, solids composed of carbon possess structures consistent with atoms in a tetrahedral or trigonal planar geometry. One common crystalline form of carbon consists of tetrahedrons linked one to another in a three-dimensional network. Another consists of a two-dimensional array of linked hexagons in which each vertex is occupied by an atom linked to three other atoms. This arrangement is consistent with the flat triangular geometry adopted by carbon when bonded to three atoms.

Many common materials are covalent solids. Diamond is a covalent crystalline form of carbon. Its high electrical resistance, exceptional melting point, and extreme hardness are properties understood in terms of its structure: The crystal lattice is composed of a network of linked tetrahedrons held together by very strong attractions. Other gemstones exhibit comparable network structures that vary in specific form depending upon the particular atoms making up the crystal lattice. Sapphire consists of a network based on aluminum and oxygen atoms, as does ruby. The lattice of garnet is built of silicon and oxygen atoms. More complicated networks are found in stones such as emerald and topaz.

Less glamorous solids based on network structures are also common.

Silicon carbide possesses a diamond structure and comparable physical properties. Quartz is a crystalline form of silica (silicon dioxide, a major component of sand) in which silicon atoms in a diamondlike lattice are bridged by intervening oxygen atoms. Quartz is a hard, high-melting insulator. Other common minerals exhibit network lattices as well, but their structures are complicated by the presence of ions. Nevertheless, many of the properties expected of covalent solids are observed in these minerals: hardness, high melting or softening point, and high electrical resistance.

Graphite is a different kind of covalent solid. Like a diamond, this solid is made up of carbon atoms. In graphite, however, each carbon atom is in a trigonal planar geometry and is linked to three other similarly arranged carbon atoms to form sheets of six-membered rings.

These sheets stack one atop another to build up the graphite solid. Because there are no covalent bonds between different sheets, they slide rather freely across one another. Thus, graphite is an excellent lubricant. Because the carbon atoms of graphite use up only three of their four available electrons to form bonds, one electron per atom remains free. Thus, graphite exhibits somewhat surprising behavior for a covalent solid; it is an excellent electrical conductor.

Applications

Covalent solids are typically used alone or in combination with other materials to manufacture products that exhibit one or more of the properties characteristic of covalent solids: hardness, resistance to chemical attack, the ability to withstand elevated temperatures, and low (or sometimes high) electrical conductivity. In everyday use, sandpaper, China, ovenware, ceramic tile, spark plug jackets, and dentures represent merely a few articles made from covalent solids. Less familiar but of great importance to the chemical industry are materials used as catalyst supports, chromatographic stationary phases, and inert electrodes. As a sampling of some of the more familiar products, consider abrasives and ceramic materials.

Abrasives consist of hard, sharp, finely divided particles and are commonly used to cut, grind, sand, or polish. When especially tough materials are to be machined, grinding and cutting tools fashioned from substances such as aluminum oxide (alumina, alundum) or silicon carbide (carborundum) are used. Sandpaper is paper coated with alumina or carborundum. Less expensive grades of sandpaper use naturally occurring materials such as flint (largely quartz) that are also covalent solids. Diamond-tipped drill bits are used to bore hard metals. Larger bits of this type are employed in geological drilling and in petroleum exploration. Diamond dust embedded in wheels is used to cut glass. For polishing, fine particles of synthetic or naturally occurring abrasives are used. Corundum and emery are aluminum oxides. Garnet and flint are quartz-like silicon oxides. Jeweler’s rouge consists of various metal oxides.

While an abrasive such as alundum may bear little resemblance to sapphire, both materials are, in fact, composed of aluminum oxide networks. In sapphire, however, there is one extended network; sapphire is a single crystal. Alundum, on the other hand, consists of innumerable small aluminum oxide crystals fused or compressed together; it is polycrystalline. In the same fashion, quartz is composed of large single crystals of silicon dioxide, while silica abrasives or powders are polycrystalline.

To harness the desirable properties of covalent solids for extensive and varied domestic and industrial uses, finely divided and inexpensive covalent solids are aggregated by heat treatment to form ceramic materials. Ceramics are complex mixtures made by firing one or more clay minerals. Clays themselves are complex mixtures but, to a large extent, can be considered covalent solids. As might be expected, ceramics, like the clays from which they are formed, exhibit hardness (when fired or compacted) and the ability to resist the effects of heat and chemical attack. In comparison to single-crystal covalent solids, however, ceramics manifest these characteristics to a lesser extent, the properties of a particular ceramic depending upon its composition and conditions of preparation. Thus, sapphire is harder, is more chemically resistant, and melts at higher temperatures than ceramics formed by sintering pure aluminum oxide or aluminum oxide clays.

The ability to control the properties of ceramics in this fashion provides great flexibility to manufacturers. By varying the clays used and the temperatures and numbers of firings, they can modify the porosity, translucence, softening temperature, and electrical resistance of the finished product. Applications mirror this variability. Consider only one type of ceramic, sintered alumina; typical applications are found as burner elements, high-temperature sample containers (crucibles and ignition boats), acid-resistant valves and pumps, spark-plug jacketing, printed circuit substrates, and parts for microwave generators. In these uses, the properties of chemical inertness, stability with respect to extreme temperature, and high electrical resistance are apparent, but other physical properties are more varied: The ceramic of a spark plug feels smooth and is nonporous; the ceramic of a crucible feels rougher and, depending upon how it was produced, may actually be pervious to gases.

Less familiar, more specialized uses for covalent solids are found in the chemical industry itself, where these solids are used as catalyst supports, as chromatographic stationary phases, and as electrodes.

In many synthetic processes, precious metal solid-phase catalysts are employed.

Platinum and palladium are common examples. Since such materials show catalytic activity only at their outer surfaces, it is economical to prepare these materials in a fashion that maximizes surface area in relation to the amount of catalyst. One such method of preparation involves dispersing the active catalytic material on the surface of an inert substance called a support. To prevent sintering of catalyst particles (and subsequent decrease in surface area and catalytic activity), supports are often required to withstand high temperatures. Among materials commonly used as supports are finely divided samples of alumina, silica (silicon dioxide), and graphite.

A similar function is performed by a chromatographic stationary phase. In the process of chromatography, components of a mixture in a mobile fluid phase are passed through a column containing a solid of high surface area. Insofar as the components exhibit different attractions to the solid, they are slowed to different extents in their passage through the column.

Thus, a separation can be affected. Because of their relative chemical inertness, the same covalent solids that have been employed as catalytic supports find use as chromatographic column packings, either as the solid phase that interacts directly with the components of the mixture or as an inert material upon which another phase, better suited to effect a particular separation, is deposited. Among the more common chromatographic solid phases are various types of aluminas and silicas.

One final interesting example of a use to which a covalent solid has been put is provided by graphite serving as an inert electrode. Unlike many covalent solids, graphite is an excellent conductor of electricity, behaving in this respect almost like a metal. In contrast to most metals, however, graphite resists chemical attack by many acids, alkalis, and salt solutions.

Thus, graphite finds use where a metal electrode cannot be employed. In the electrolytic production of aluminum by the Hall process, graphite electrodes immersed in a hot molten salt solution are used to effect the transformation of corundum to the metal. This interesting process illustrates a covalent solid (graphite) used to transform another covalent solid (aluminum oxide or corundum) into a metal of great economic importance (aluminum).

Thus, the uses of covalent solids, though varied, are typically in applications that capitalize on their hardness, resistance to chemical attack, and ability to withstand elevated temperatures, properties that result from their crystalline structure and cohesive forces.

Advances in the field have led to the development of engineered covalent network materials such as covalent organic frameworks (COFs) and related structures, which are used in gas storage, catalysis, and separation technologies.

Context

Describing a solid as covalent implies a classification by which the material is distinguished from ionic, metallic, and molecular crystals. While the empirical basis of this classification is rooted in properties that were surely known in prehistoric times (hardness, ease of melting, luster, and the ability to conduct heat), it was only as an understanding of the nature of chemical bonding emerged in the twentieth century that the underlying reasons for the differing properties of classes of solids became clear.

The idea of a chemical bond holding atoms together in a structured molecule grew during the latter part of the nineteenth century as a result of the work of chemists, including Friedrich August Kekule, Joseph-Achille Le Bel, and Jacobus van ‘t Hoff. From their independent efforts, an insight into the directional properties of the bonds of carbon was achieved. Knowledge of what constituted such a bond, however, did not come until later.

By the beginning of the twentieth century, the existence of the electron had been demonstrated and some of its properties investigated by Sir Joseph Thomson. Shortly thereafter, several chemists began to speculate about the possible role of electrons in bonding, and as the properties of a nuclear planetary model of the atom emerged from the work of Ernest Rutherford and later Niels Bohr, attempts to explain bonding in terms of electrons increased. Between the years 1916 and 1919, Gilbert Newton Lewis, Walther Kossel, and Irving Langmuir, in separate efforts, developed the idea of a chemical bond as consisting of a shared pair of electrons.

Langmuir termed such a bond “covalent.” With the emergence of quantum mechanics in the late 1920s, the concept of covalent bonding was placed on a firm theoretical footing. In 1927, the quantum mechanical explanation of covalent bonding was developed in a paper by Fritz London and Walter Heitler. Subsequently, valence bond theory, as the quantum theory of covalent bonding came to be called, was extended by the efforts of many investigators.

About the same time that Rutherford and Bohr were developing theories of atomic structure, an understanding of the solid state was growing. In 1912, Max von Laue suggested that X-rays, discovered about twenty-seven years earlier by Wilhelm Röntgen, could be used to investigate the crystal lattices of solids. Within a year, there was experimental verification of von Laue’s prediction, and soon afterward, William Henry Bragg and his son, Lawrence Bragg, solved the mathematical problem of extracting structural information from X-ray data, determining the structure of sodium chloride in the process. Within a short time, structures of many crystals had been determined, including those of the covalent solids diamond and graphite. Later developments have rendered structural determination by X-ray crystallography routine for most solids.

As knowledge of the solid state grew in the late twentieth and twenty-first centuries, the inadequacy of classifying crystals based on the attractive forces within the crystal became apparent. One difficulty arises from the observation that many chemical bonds exhibit both ionic and covalent character; that is, there exist bonds in which electron sharing is unequal enough to approach electron transfer. Further complicating matters is the tendency of metallic character to enter bonding in the solid state. Thus, tin, in the same chemical family as carbon, forms a metallic crystal rather than a covalent crystal. Another difficulty arises in classifying solids in which more than one type of chemical bonding is important. Typical are minerals in which positive ions are interspersed among sheets or networks of covalently bonded nonmetallic atoms that collectively bear negative charges.

It is clear that the concept of a covalent crystal is a useful idea. It provides a basis for simply rationalizing the observed properties of many common and exotic solids. Nevertheless, as the need for materials with characteristics narrowly tailored for specific applications grows, it is necessary to move to an understanding of the detailed structure of specific solids and beyond an oversimplified classification as ionic, metallic, molecular, or covalent.

Principal terms

ATOM: the smallest part of an element; atoms are the fundamental building blocks of chemistry

CRYSTAL: a solid in which external surfaces are flat and make definite angles with one another; crystals exhibit long-range order at the molecular level

ELECTRON: a fundamental particle and the smallest unit of negative charge; electrons are found within atoms

ION: a charged particle; starting from a neutral molecule, positive ions result from removing one or more electrons, negative ions from adding one or more electrons

LATTICE: a repeating spatial pattern; the long-range order of crystals consists of structural units in a three-dimensional lattice

MOLECULE: the smallest part of an element or compound that commonly exists; molecules of compounds consist of two or more atoms bonded together

THERMAL ENERGY: energy of molecular motion possessed by objects at temperatures above absolute zero; any material containing heat is composed of units that move or vibrate randomly

Bibliography

“Covalent Solids.” University of Texas at Austin, ch301.cm.utexas.edu/imfs/#solids/solid-props-all.php. Accessed 22 Apr. 2026.

Galwey, Andrew K. Chemistry of Solids. Chapman and Hall, 1967.

Jolly, William L. Modern Inorganic Chemistry. McGraw-Hill, 1984.

Krock, Richard H., and Merrill L. Ebner. Ceramics, Plastics, and Metals. D. C. Heath, 1965.

Mackay, K. M., et al. Introduction to Modern Inorganic Chemistry. 6th ed., CRC Press, 2017.

Mortimer, Charles E. Chemistry. 6th ed., Wadsworth, 1986.

Norton, F. H. Fine Ceramics: Technology and Applications. McGraw-Hill, 1970.

Tilley, R. J. D. Understanding Solids: The Science of Materials. 3rd ed., John Wiley & Sons Ltd., 2021.

Tolédano, Jean-Claude. Physics of Electrons in Solids. World Scientific Publishing Company, 2021.

Wells, A. F. Structural Inorganic Chemistry. 5th ed., Clarendon Press, 2012.

Yu, Baoqiu, et al. “Single-Crystal 2D Covalent Organic Frameworks for High-Capacity Methane Storage.” Nature Communications, vol. 17, 2026, article 2740, doi:10.1038/s41467-026-69614-7. Accessed 22 Apr. 2026.