When a pure element can exist in different structural forms, those forms are known as allotropes. Although they are composed of the exact same atoms, the different arrangements of these atoms lead to dramatically varied physical properties, such as hardness, density, and electrical conductivity. Carbon is unique among elements because its ability to form four stable chemical bonds allows it to organize into various three-dimensional structures. The distinct ways carbon atoms link together create the most well-known allotropes, which range from the hardest natural substance to a soft, slippery solid.
Diamond: The Hardest Allotrope
Diamond is defined by its highly ordered, three-dimensional lattice structure. In this arrangement, every single carbon atom is covalently bonded to four neighboring carbon atoms, forming a repeating tetrahedral unit. This uniform, giant covalent structure extends throughout the entire crystal, making it incredibly strong. The density of diamond is high, typically around 3.51 grams per cubic centimeter, which is a direct result of this tightly packed atomic framework.
The extreme hardness of diamond, which ranks as a 10 on the Mohs scale, comes from the difficulty of breaking these numerous strong bonds simultaneously. Diamond is a poor conductor of electricity because all four valence electrons of each carbon atom are immobilized in the strong covalent bonds. With no free electrons available to move through the structure, electrical current cannot flow easily. These properties make diamond invaluable for industrial applications like high-precision cutting and polishing tools, in addition to its use in jewelry.
Graphite: The Layered Allotrope
In stark contrast to diamond, graphite possesses a layered structure that dictates its soft nature and electrical conductivity. Carbon atoms in graphite are arranged in flat, two-dimensional sheets of hexagonal rings, where each atom is bonded to only three neighbors. These individual sheets are known as graphene, and the strong covalent bonds within them are actually stronger than those in diamond.
The sheets, however, are stacked and held together only by weak intermolecular forces, specifically Van der Waals forces, which are easily overcome. This weak interlayer attraction allows the sheets to slide past one another with minimal effort, making graphite feel soft and slippery. This property is why graphite is commonly used as a lubricant and in the “lead” of pencils. The fourth valence electron from each carbon atom is delocalized and free to move throughout the entire layer, making graphite an excellent electrical conductor.
Fullerenes: The Cage-Like Allotrope
Fullerenes represent a third, structurally distinct family of carbon allotropes discovered relatively recently in 1985. The most widely recognized member is Buckminsterfullerene, or C60, a molecule consisting of exactly 60 carbon atoms. This structure forms a hollow, closed cage that perfectly resembles a soccer ball or a geodesic dome.
The C60 molecule is composed of 12 pentagonal rings and 20 hexagonal rings fused together, creating a highly symmetrical, spherical shape often nicknamed a “buckyball.” Unlike the macroscopic lattice structures of diamond and graphite, fullerenes are discrete, individual molecules. Their unique, hollow structure allows them to be dissolved in organic solvents, which is not possible for the other two allotropes. Fullerenes are the subject of research for potential applications in electronics, medicine, and as high-temperature superconductors when combined with certain elements.