Carbon in its elemental forms (charcoal, soot, diamond, graphite) has been known since prehistoric times. Humans used charcoal as fuel and to create pigments in cave paintings over 30,000 years ago. Diamonds were prized as precious gems since antiquity. However, it was not until 1772 that Antoine Lavoisier (1743-1794) demonstrated that diamond is a form of pure carbon by burning it and observing that only carbon dioxide was produced. In 1779, the Swedish chemist Carl Wilhelm Scheele (1742-1786) showed that graphite was also pure carbon. The name carbon derives from the Latin carbo meaning "coal". In 1985, a new allotropic form of carbon was discovered: fullerenes (C₆₀), followed by carbon nanotubes in 1991 and graphene in 2004, revolutionizing materials science.
Carbon (symbol C, atomic number 6) is a non-metal of group 14 in the periodic table, consisting of six protons, usually six neutrons (for the most common isotope), and six electrons. The stable isotopes are carbon-12 \(\,^{12}\mathrm{C}\) (≈ 98.93%) and carbon-13 \(\,^{13}\mathrm{C}\) (≈ 1.07%). Carbon-14 \(\,^{14}\mathrm{C}\) is radioactive with a half-life of 5,730 years, used for archaeological dating.
Carbon exhibits several allotropic forms with radically different properties. Diamond is transparent, extremely hard (10 on the Mohs scale), an electrical insulator, with a tetrahedral structure where each atom is bonded to four others (sp³ hybridization). Graphite is opaque, black, soft, electrically conductive, with a layered planar structure where each atom is bonded to three others (sp² hybridization). Other forms include fullerenes (spherical or ellipsoidal structures), carbon nanotubes (rolled graphene sheets), and graphene (a single atomic layer of graphite). Amorphous carbon exists in forms such as charcoal, carbon black, and soot.
Density: graphite ≈ 2.26 g/cm³, diamond ≈ 3.51 g/cm³. Melting point (graphite, under pressure): ≈ 3823 K (3550 °C). Sublimation point (graphite, atmospheric pressure): ≈ 3915 K (3642 °C).
| Isotope / Notation | Protons (Z) | Neutrons (N) | Atomic mass (u) | Natural abundance | Half-life / Stability | Decay / Remarks |
|---|---|---|---|---|---|---|
| Carbon-11 — \(\,^{11}\mathrm{C}\,\) | 6 | 5 | 11.011433 u | Unnatural | 20.334 minutes | Radioactive β\(^+\) decay to \(\,^{11}\mathrm{B}\); used in positron emission tomography (PET). |
| Carbon-12 — \(\,^{12}\mathrm{C}\,\) | 6 | 6 | 12.000000 u (by definition) | ≈ 98.93 % | Stable | Reference isotope for the atomic mass scale; basis of all organic chemistry. |
| Carbon-13 — \(\,^{13}\mathrm{C}\,\) | 6 | 7 | 13.003355 u | ≈ 1.07 % | Stable | Used in carbon-13 NMR spectroscopy and as a tracer in biochemistry and geochemistry. |
| Carbon-14 — \(\,^{14}\mathrm{C}\,\) | 6 | 8 | 14.003241 u | Traces (cosmogenic) | 5,730 years | Radioactive β\(^-\) decay to \(\,^{14}\mathrm{N}\); produced by cosmic rays; used for carbon-14 dating. |
| Carbon-15 — \(\,^{15}\mathrm{C}\,\) | 6 | 9 | 15.010599 u | Unnatural | 2.449 s | Radioactive β\(^-\); artificially produced in particle accelerators. |
| Other isotopes — \(\,^{8}\mathrm{C}-\,^{10}\mathrm{C},\,^{16}\mathrm{C}-\,^{22}\mathrm{C}\) | 6 | 2-4, 10-16 | — (resonances) | Unnatural | \(10^{-21}\) — 0.747 s | Unstable states observed in nuclear physics; some have neutron halo structures. |
N.B. :
Electron shells: How electrons organize around the nucleus.
Carbon has 6 electrons distributed across two electron shells. Its full electronic configuration is: 1s² 2s² 2p², or simplified as: [He] 2s² 2p². This configuration can also be written as: K(2) L(4).
K Shell (n=1): Contains 2 electrons in the 1s sub-shell. This inner shell is complete and highly stable.
L Shell (n=2): Contains 4 electrons distributed as 2s² 2p². The 2s orbitals are complete, while the 2p orbitals contain only 2 electrons out of 6 possible. Thus, 4 electrons are missing to reach the stable neon configuration with 8 electrons (octet).
The 4 electrons in the outer shell (2s² 2p²) are the valence electrons of carbon. This configuration explains its chemical properties:
By gaining 4 electrons, carbon theoretically forms the C⁴⁻ ion (oxidation state -4), a very rare state observed only in certain metallic carbides like Al₄C₃, adopting the stable neon configuration [Ne].
By losing 4 electrons, carbon would form the C⁴⁺ ion (oxidation state +4), another rare ionic state but observed in carbon dioxide CO₂ in covalent form.
Carbon can exhibit intermediate oxidation states: -3, -2, -1, 0, +1, +2, +3, and +4, with a preference for covalent bonds rather than ionic ones.
The oxidation state 0 corresponds to elemental carbon, which exists in several allotropic forms: diamond, graphite, graphene, fullerenes, and carbon nanotubes.
The electronic configuration of carbon, with 4 electrons in its valence shell, places it in group 14 of the periodic table and gives it a unique position in chemistry. This structure grants it exceptional characteristic properties: the ability to form four stable covalent bonds by sharing its valence electrons, the aptitude to form single, double, or triple bonds with itself and other elements, and the possibility to create linear, branched, or cyclic chains of virtually unlimited length (catenation). Carbon rarely forms ions because the energies required to gain or lose 4 electrons are too high. It therefore favors covalent bonds where electrons are shared. This unique ability to form complex and varied molecular structures makes carbon the fundamental element of organic chemistry and life.
The importance of carbon is absolutely fundamental: it is the basic element of all organic chemistry and life on Earth. All living organisms are built around carbon-based molecules (carbohydrates, lipids, proteins, nucleic acids). Carbon forms more compounds than all other elements combined, with millions of different organic molecules known. Its allotropic forms have remarkable properties: diamond is the hardest natural material, graphite is an excellent conductor and lubricant, and graphene has exceptional electronic properties. The carbon cycle regulates Earth's climate via atmospheric CO₂. Industrially, carbon is used in the form of coal, coke, carbon black, carbon fibers, and in countless applications ranging from fuels to high-tech materials.
Carbon has four valence electrons, allowing it to form four stable covalent bonds. This tetravalence and carbon's ability to form single, double, and triple bonds with itself and other elements (notably hydrogen, oxygen, nitrogen, sulfur, phosphorus) are the basis for the extraordinary diversity of organic chemistry. Carbon can form linear, branched, cyclic chains, and complex three-dimensional structures, creating millions of different organic compounds.
Elemental carbon is relatively inert at room temperature but reacts with oxygen at high temperatures to form carbon dioxide (CO₂) or carbon monoxide (CO) depending on conditions. It reacts with metals to form carbides, with hydrogen to produce hydrocarbons, and with halogens to give carbon tetrahalides. Carbon can exist in several hybridization states (sp, sp², sp³), determining the geometry and properties of its compounds.
Carbon is the central element of biochemistry. All known living organisms are based on carbon-containing organic molecules: carbohydrates, lipids, proteins, nucleic acids. The unique ability of carbon to form stable and complex macromolecules makes it the fundamental element of life as we know it.
Carbon is the fourth most abundant element in the observable universe (after hydrogen, helium, and oxygen) and plays an absolutely central role in stellar and galactic evolution. Unlike hydrogen and helium, which come from the Big Bang, carbon is entirely produced by stellar nucleosynthesis in the cores of massive stars.
The formation of carbon in stars occurs through the triple-alpha process: three helium-4 nuclei fuse to form carbon-12 at temperatures above 100 million kelvin. This reaction is made possible by the existence of a particular excited state of carbon-12 (Hoyle state, predicted in 1953 by Fred Hoyle), which allows overcoming the "beryllium-8 gap". Without this remarkable quantum coincidence, carbon could not form efficiently, and carbon-based life would probably not exist. This observation was one of the first arguments for the anthropic principle in cosmology.
In evolved massive stars, carbon serves as fuel for subsequent fusion reactions (carbon burning) at temperatures of about 600 million kelvin, producing neon, sodium, and magnesium. Carbon is also an essential catalyst in the CNO cycle (carbon-nitrogen-oxygen), a hydrogen-to-helium fusion process that dominates in stars more massive than the Sun.
Stars at the end of their lives enrich the interstellar medium with carbon through stellar winds and supernova explosions. AGB stars (asymptotic giant branch) are particularly effective at producing and dispersing carbon, creating carbon stars where carbon is more abundant than oxygen in the stellar atmosphere. This stellar carbon forms carbon dust in the interstellar medium, which plays a crucial role in the formation of new stars and planets.
In the interstellar medium, carbon exists in several forms: atomic (C, C⁺), molecular (CO, C₂, carbon chains), as graphite grains, and polycyclic aromatic hydrocarbons (PAHs). Carbon monoxide (CO) is the second most abundant molecular compound after H₂ and serves as the main tracer for mapping cold molecular clouds where stars form.
The isotopic ratio ¹²C/¹³C in stars and planetary objects provides valuable information about nucleosynthesis processes, stellar mixing, and galactic chemical evolution. Variations in this ratio observed in primitive meteorites and interstellar dust reveal the diversity of stellar sources that contributed to the formation of our solar system.
N.B.:
The cosmic carbon cycle illustrates the deep interconnection between stars, the interstellar medium, and planetary formation. Carbon formed in stars is dispersed into space, incorporates into molecular clouds, participates in the formation of new generations of stars and planetary systems, and ultimately allows the emergence of life on planets like Earth. On our planet, carbon cycles between the atmosphere (CO₂), oceans (dissolved carbonates), the biosphere (living organic matter), and the lithosphere (carbonate rocks, fossil fuels) in a complex geochemical cycle. Humanity is currently disrupting this cycle by massively releasing fossil carbon into the atmosphere, causing global climate warming. Understanding the carbon cycle at all scales, from cosmic to planetary, is therefore essential not only for fundamental science but also for the future of our civilization.
