Carbon is capable of forming many allotropes (structurally different forms of the same element) due to its valency (tetravalent). Well-known forms of carbon include diamond and graphite. In recent decades, many more allotropes have been discovered and researched, including ball shapes such as buckminsterfullerene and sheets such as graphene. Larger-scale structures of carbon include nanotubes, nanobuds carbon quantum dots and nanoribbons. Other unusual forms of carbon exist at very high temperatures or extreme pressures. Around 500 hypothetical 3‑periodic allotropes of carbon are known , according to the Samara Carbon Allotrope Database (SACADA).
Atomic and diatomic carbon
Under certain conditions, carbon can be found in its atomic form. It can be formed by vaporizing graphite, by passing large electric currents to form a carbon arc under very low pressure. It is extremely reactive, but it is an intermediate product used in the creation of carbenes.
Diatomic carbon can also be found under certain conditions. It is often detected via spectroscopy in extraterrestrial bodies, including comets and certain stars. Diamond
Diamond is a well-known allotrope of carbon. The hardness, extremely high refractive index, and high dispersion of light make diamond useful for industrial applications and for jewelry. Diamond is the hardest known natural mineral. This makes it an excellent abrasive and makes it hold polish and luster extremely well. No known naturally occurring substance can cut or scratch diamond, except another diamond. In diamond form, carbon is one of the costliest elements.
The crystal structure of diamond is a face-centered cubic lattice having eight atoms per unit cell to form a diamond cubic structure. Each carbon atom is covalently bonded to four other carbons in a tetrahedral geometry. These tetrahedrons together form a 3-dimensional network of six-membered carbon rings in the chair conformation, allowing for zero bond angle strain. The bonding occurs through sp3 hybridized orbitals to give a C-C bond length of 154 pm. This network of unstrained covalent bonds makes diamond extremely strong. Diamond is thermodynamically less stable than graphite at pressures below .
The dominant industrial use of diamond is cutting, drilling (drill bits), grinding (diamond edged cutters), and polishing. Most uses of diamonds in these technologies do not require large diamonds, and most diamonds that are not gem-quality can find an industrial use. Diamonds are embedded in drill tips and saw blades or ground into a powder for use in grinding and polishing applications (due to its extraordinary hardness). Specialized applications include use in laboratories as containment for high pressure experiments (see diamond anvil), high-performance bearings, and specialized windows of technical apparatuses.
The market for industrial-grade diamonds operates much differently from its gem-grade counterpart. Industrial diamonds are valued mostly for their hardness and heat conductivity, making many of the gemological characteristics of diamond, including clarity and color, mostly irrelevant. This helps explain why 80% of mined diamonds (equal to about 100 million carats or 20 tonnes annually) are unsuitable for use as gemstones and known as bort, destined for industrial use. In addition to mined diamonds, synthetic diamonds found industrial applications almost immediately after their invention in the 1950s; another 400 million carats (80 tonnes) of synthetic diamonds are produced annually for industrial use, which is nearly four times the mass of natural diamonds mined over the same period.
Continuing advances in the production of synthetic diamond have made future applications feasible, combined with increased quality and quantity of supply available from synthetic diamond manufacturers. Possible uses, covered by research efforts in Japan, Europe, and the United States, include diamond as a semiconductor suitable to build microchips from, or as a heat sink in electronics.
Graphite
Graphite, named by Abraham Gottlob Werner in 1789, from the Greek γράφειν (, "to draw/write", for its use in pencils) is one of the most common allotropes of carbon. Unlike diamond, graphite is an electrical conductor. Thus, it can be used in, for instance, electrical arc lamp electrodes. Likewise, under standard conditions, graphite is the most stable form of carbon. Therefore, it is used in thermochemistry as the standard state for defining the heat of formation of carbon compounds.
Graphite conducts electricity, due to delocalization of the pi bond electrons above and below the planes of the carbon atoms. These electrons are free to move, so are able to conduct electricity. However, the electricity is only conducted along the plane of the layers. In diamond, all four outer electrons of each carbon atom are 'localized' between the atoms in covalent bonding. The movement of electrons is restricted and diamond does not conduct an electric current. In graphite, each carbon atom uses only 3 of its 4 outer energy level electrons in covalently bonding to three other carbon atoms in a plane. Each carbon atom contributes one electron to a delocalized system of electrons that is also a part of the chemical bonding. The delocalized electrons are free to move throughout the plane. For this reason, graphite conducts electricity along the planes of carbon atoms, but does not conduct electricity in a direction at right angles to the plane.
Graphite powder is used as a dry lubricant. Although it might be thought that this industrially important property is due entirely to the loose interlamellar coupling between sheets in the structure, in fact in a vacuum environment (such as in technologies for use in space), graphite was found to be a very poor lubricant. This fact led to the discovery that graphite's lubricity is due to adsorbed air and water between the layers, unlike other layered dry lubricants such as molybdenum disulfide. Recent studies suggest that an effect called superlubricity can also account for this effect.
When a large number of crystallographic defects (physical) bind these planes together, graphite loses its lubrication properties and becomes pyrolytic carbon, a useful material in blood-contacting implants such as prosthetic heart valves.
Graphite is the most stable allotrope of carbon. Contrary to popular belief, high-purity graphite does not readily burn, even at elevated temperatures. For this reason, it is used in nuclear reactors and for high-temperature crucibles for melting metals. At very high temperatures and pressures (roughly 2000 °C and 5 GPa), it can be transformed into diamond.
Natural and crystalline graphites are not often used in pure form as structural materials due to their shear-planes, brittleness and inconsistent mechanical properties.
In its pure glassy (isotropic) synthetic forms, pyrolytic graphite and carbon fiber graphite are extremely strong, heat-resistant (to 3000 °C) materials, used in reentry shields for missile nosecones, solid rocket engines, high temperature reactors, brake shoes and electric motor brushes.
Intumescent or expandable graphites are used in fire seals, fitted around the perimeter of a fire door. During a fire the graphite intumesces (expands and chars) to resist fire penetration and prevent the spread of fumes. A typical start expansion temperature (SET) is between 150 and 300 °C.
Graphite's specific gravity is 2.3, which makes it less dense than diamond.
Graphite is slightly more reactive than diamond. This is because the reactants are able to penetrate between the hexagonal layers of carbon atoms in graphite. It is unaffected by ordinary solvents, dilute acids, or fused alkalis. However, chromic acid oxidizes it to carbon dioxide.
Graphene
A single layer of graphite is called graphene and has extraordinary electrical, thermal, and physical properties. It can be produced by epitaxy on an insulating or conducting substrate or by mechanical exfoliation (repeated peeling) from graphite. Its applications may include replacing silicon in high-performance electronic devices. With two layers stacked, bilayer graphene results with different properties.
Lonsdaleite (hexagonal diamond) Lonsdaleite is an allotrope sometimes called "hexagonal diamond", formed from graphite present in meteorites upon their impact on the earth. The great heat and pressure of the impact transforms the graphite into a denser form similar to diamond but retaining graphite's hexagonal crystal lattice. "Hexagonal diamond" has also been synthesized in the laboratory, by compressing and heating graphite either in a static press or using explosives. It can also be produced by the thermal decomposition of a polymer, poly(hydridocarbyne), at atmospheric pressure, under inert gas atmosphere (e.g. argon, nitrogen), starting at temperature .
Graphenylene Graphenylene is a single layer carbon material with biphenylene-like subunits as basis in its hexagonal lattice structure. It is also known as biphenylene-carbon.
Carbophene Carbophene is a 2 dimensional covalent organic framework. 4-6 carbophene has been synthesized from 1-3-5 trihydroxybenzene. It consists of 4-carbon and 6-carbon rings in 1:1 ratio. The angles between the three σ-bonds of the orbitals are approximately 120°, 90°, and 150°.
AA'-graphite AA'-graphite is an allotrope of carbon similar to graphite, but where the layers are positioned differently to each other as compared to the order in graphite.
Diamane Diamane is a 2D form of diamond. It can be made via high pressures, but without that pressure, the material reverts to graphene. Another technique is to add hydrogen atoms, but those bonds are weak. Using fluorine (xenon-difluoride) instead brings the layers closer together, strengthening the bonds. This is called f-diamane.
Amorphous carbon
Amorphous carbon is the name used for carbon that does not have any crystalline structure. As with all glassy materials, some short-range order can be observed, but there is no long-range pattern of atomic positions. While entirely amorphous carbon can be produced, most amorphous carbon contains microscopic crystals of graphite-like, or even diamond-like carbon.
Coal and soot or carbon black are informally called amorphous carbon. However, they are products of pyrolysis (the process of decomposing a substance by the action of heat), which does not produce true amorphous carbon under normal conditions.
Nanocarbons
Buckminsterfullerenes
The buckminsterfullerenes, or usually just fullerenes or buckyballs for short, were discovered in 1985 by a team of scientists from Rice University and the University of Sussex, three of whom were awarded the 1996 Nobel Prize in Chemistry. They are named for the resemblance to the geodesic structures devised by Richard Buckminster "Bucky" Fuller. Fullerenes are positively curved molecules of varying sizes composed entirely of carbon, which take the form of a hollow sphere, ellipsoid, or tube (the C60 version has the same form as a traditional stitched soccer ball).
As of the early twenty-first century, the chemical and physical properties of fullerenes are still under heavy study, in both pure and applied research labs. In April 2003, fullerenes were under study for potential medicinal use — binding specific antibiotics to the structure to target resistant bacteria and even target certain cancer cells such as melanoma.
Nanotubes
Carbon nanotubes, also called buckytubes, are cylindrical carbon molecules with novel properties that make them potentially useful in a wide variety of applications (e.g., nano-electronics, optics, materials applications, etc.). They exhibit extraordinary strength, unique electrical properties, and are efficient conductors of heat. Non-carbon nanotubes have also been synthesized.