Carbon

• 1. Introduction

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  Carbon is the element with the symbol C and atomic number 6, and it hides in plain sight inside almost everything alive. It makes up only about 0.025 percent of Earth's crust, yet it is the second most abundant element in the human body by mass, at about 18.5 percent, after oxygen. The same element appears as the softest material and the hardest, as a black smudge and a transparent gem. How does one element wear so many faces? Why does it sit at the center of all known life, and how did humans come to recognize that coal, charcoal, graphite, and diamond are all the same thing? The answers run from the interiors of dying stars to the tattoos on a 5,200-year-old corpse.

• 2. Many Faces

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  Graphite is soft enough to leave a streak on paper, which is why its name comes from the Greek verb meaning to write. Each of its atoms bonds to three others in flat sheets of fused hexagonal rings, and those sheets slip easily past one another through weak van der Waals forces. A loose outer electron from every atom forms a roaming pi-cloud, so graphite conducts electricity, but only within each sheet. That same delocalization makes graphite the more stable form at room temperature. Diamond reverses nearly every one of those traits. At very high pressures carbon packs into diamond, with nearly twice the density of graphite, each atom bonded tetrahedrally to four others in a three-dimensional network of puckered six-membered rings. The bonds are so strong that diamond is the hardest naturally occurring substance measured by resistance to scratching, yet it is an excellent electrical insulator. Despite the slogan that diamonds are forever, they are thermodynamically unstable and should slowly become graphite, held back only by a high activation energy barrier. At very high temperatures the transition speeds up, and a diamond can burn up in a house fire. Fullerenes, discovered in 1985, bend carbon into closed shapes by mixing pentagons and other rings into the hexagonal pattern, warping the sheets into spheres, ellipses, and cylinders. The names fullerene and buckyball honor Richard Buckminster Fuller, the popularizer of geodesic domes that the molecules resemble. Carbon nanotubes form hollow cylinders, while nanobuds, first reported in 2007, graft buckyballs onto a nanotube's outer wall. Graphene, a single hexagonal sheet, appeared as of 2009 to be the strongest material ever tested, with proposed uses ranging from a space elevator to storing hydrogen for car engines. The oddities do not stop there. Carbon nanofoam, found in 1997, is ferromagnetic and among the lightest known solids, around 2 kilograms per cubic metre. Linear acetylenic carbon, or carbyne, is a chain of alternating single and triple bonds whose Young's modulus is 40 times that of diamond. In 2015, a team at North Carolina State University fired a high-energy, low-duration laser pulse at amorphous carbon dust to make Q-carbon, reported to be ferromagnetic, fluorescent, and harder than diamond.

• 3. The Reluctant Element

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  At standard temperature and pressure, most forms of carbon resist all but the strongest oxidizers, ignoring sulfuric acid, hydrochloric acid, chlorine, and alkalis alike. Graphite is the exception that proves the rule. Despite being the more stable form, its exposed delocalized pi system makes it more reactive than diamond, and hot concentrated nitric acid can oxidize it to mellitic acid while preserving its hexagonal units. Heat changes everything. At elevated temperatures carbon reacts with oxygen to form carbon oxides, and it will strip oxygen from metal oxides to leave the bare metal. That exothermic reaction is the backbone of the iron and steel industry, used to smelt iron and to control the carbon content of steel. Carbon also joins with sulfur to make carbon disulfide and reacts with steam in the coal-gas reaction used in coal gasification. Carbon's behavior in its rarer phases is far less studied than its solids. It has the highest sublimation point of all elements and no melting point at atmospheric pressure, because its triple point sits at 10.8 megapascals and 4,600 kelvins, so it sublimes near 3,900 kelvins. In the vapor phase some carbon appears as highly reactive diatomic dicarbon, which glows green when excited. The liquid phase is a dark, mobile, reflective fluid that exists only above 4,000 kelvins and under pressures over 100 atmospheres.

• 4. Clocks And Halos

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  Carbon-12 makes up 98.93 percent of the carbon on Earth, while carbon-13 accounts for the remaining 1.07 percent. Both are stable, and biological materials carry slightly more carbon-13 because biochemical reactions discriminate against it. In 1961, the International Union of Pure and Applied Chemistry adopted carbon-12 as the basis for atomic weights, and carbon-13 is the isotope read in nuclear magnetic resonance experiments. Carbon-14 turns the element into a clock. Cosmic rays striking nitrogen in the upper atmosphere create this radioisotope, found in trace amounts of about 1 part per trillion. It decays with a half-life of 5,700 years, so it is virtually absent in ancient rocks. While an organism lives, its carbon-14 stays nearly constant, then declines predictably after death. Radiocarbon dating, invented in 1949, uses this decline to date carbonaceous materials up to about 40,000 years old. The carbon family has stranger members at the edges. There are 15 known isotopes, and the shortest-lived decays through proton emission with a half-life of 3.5 seconds. One exotic isotope shows a nuclear halo, meaning its radius is appreciably larger than a sphere of constant density would predict.

• 5. Forged In Stars

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  Carbon is the fourth most abundant element in the observable universe by mass, after hydrogen, helium, and oxygen. None of it came from the Big Bang. Its nucleus forms inside giant and supergiant stars through the triple-alpha process, a near-simultaneous collision of three helium nuclei, because the intermediate steps produce lithium-5 and beryllium-8, both so unstable they decay almost instantly. The process needs temperatures above 100 megakelvins and a helium concentration that the rapidly cooling early universe could not provide. When massive stars die as supernovae, they scatter carbon into space as dust. That dust becomes raw material for the next generation of star systems and their planets, and most scholars hold that all the carbon in the Solar System and the Milky Way comes from dying stars. Carbon also earns its keep as a catalyst, powering stars through the CNO cycle and providing most of the energy in larger stars such as Sirius. More than 20 percent of the universe's carbon may be locked in polycyclic aromatic hydrocarbons, complex compounds of carbon and hydrogen without oxygen. NASA announced a greatly upgraded database for tracking them in 2014. These PAHs appear to have formed a couple of billion years after the Big Bang and feature in the PAH world hypothesis about the origin of life. Some meteorites even carry microscopic diamonds formed when the Solar System was still a protoplanetary disk.

• 6. The Chemistry Of Life

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  Carbon can link to itself in endless chains, a property called catenation, and its carbon-carbon bonds are strong and stable. A tally of unique compounds shows that more contain carbon than do not, with about two hundred million already described and indexed, still only a fraction of the theoretically possible total. In nearly all stable organic compounds, carbon obeys the octet rule and forms four covalent bonds, which may include double and triple bonds. The simplest organic molecule is the hydrocarbon, hydrogen atoms bonded to a chain of carbon. Swap in heteroatoms like oxygen, nitrogen, sulfur, or phosphorus, and carbon builds the molecules of life. With oxygen and hydrogen it forms sugars, alcohols, fats, and terpenes; with nitrogen, alkaloids; with sulfur, antibiotics and amino acids. Add phosphorus and it forms DNA, RNA, and adenosine triphosphate, the chief energy-transfer molecule in every living cell. Carbon's exclusivity for life is not just chemistry but considered opinion. Norman Horowitz, who headed the Mariner and Viking missions to Mars from 1965 to 1976, judged carbon's unique characteristics so essential that no other element could replace it, even on another planet. The element also bends its own rules in exotic compounds. In 2016 researchers confirmed that the hexamethylbenzene dication holds a carbon atom with six bonds, and in nature the iron-molybdenum cofactor for microbial nitrogen fixation centers on a carbon bonded to six iron atoms.

• 7. From Antiquity To Lavoisier

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  Diamonds were known probably as early as 2500 BCE in China, while charcoal was made then by the same chemistry as today, heating wood in a clay-covered pyramid to exclude air. The English word carbon descends from the Latin carbo for coal and charcoal, and German, Dutch, and Danish all use words that literally mean coal-substance. Carbon counts among the few elements known since antiquity. The modern understanding arrived in a burst of 18th-century experiments. In 1722, Rene Antoine Ferchault de Reaumur showed that iron became steel by absorbing some substance, now known to be carbon. In 1772, Antoine Lavoisier burned charcoal and diamond, found neither produced water, and saw both release the same amount of carbon dioxide per gram, proving diamond is a form of carbon. In 1779, Carl Wilhelm Scheele showed that graphite, long mistaken for a form of lead, was charcoal with a trace of iron. The element finally got its name in 1786, when Claude Louis Berthollet, Gaspard Monge, and C. A. Vandermonde oxidized graphite and proposed calling the element carbone. Lavoisier listed carbon as an element in his 1789 textbook. Two centuries later, the story added a new chapter when Robert Curl, Harold Kroto, and Richard Smalley won the Nobel Prize in Chemistry in 1996 for the fullerene discovered in 1985.

• 8. Mining, Markets, And Hazards

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  Vein graphite, the rarest and highest quality natural type, is only commercially mined in Sri Lanka. World production of natural graphite reached 1.1 million tonnes in 2010, with China contributing 800,000 tonnes, India 130,000, Brazil 76,000, North Korea 30,000, and Canada 25,000. The United States mined no natural graphite, but produced 118,000 tonnes of synthetic graphite valued at an estimated 998 million dollars in 2009. Long before that, the pure deposits at Borrowdale in Cumberland, England, were sawn into strips and encased in wood to make pencils until the 19th century. Diamonds tell a story of shifting geography. India led world production from roughly the 9th century BC to the mid-18th century AD, until Brazil yielded the first non-Indian diamonds in 1725. Primary deposits opened only in the 1870s after the South African fields were found, and over 4.5 billion carats have been mined since. By 2005 Russia produced almost one-fifth of global output, while Australia's Argyle mine became the single largest source, producing 14 million carats in 2018. Synthetic diamonds, invented in the 1950s, now reach 3 billion carats, or 600 tonnes, a year, and about 80 percent of mined diamonds are industrial-grade bort. For all its usefulness, carbon carries dangers tied to form and setting. Inhaling coal dust or soot in quantity irritates lung tissue and causes coalworker's pneumoconiosis, and large coal accumulations can spontaneously combust when long-buried deposits meet air. In nuclear reactors where graphite moderates neutrons, stored Wigner energy can release suddenly; annealing to at least 250 degrees Celsius vents it safely, but the procedure went wrong in the Windscale fire. Yet pure carbon itself has extremely low toxicity, which is why Otzi the Iceman's carbon tattoos survived for 5,200 years after his death.