Carbon (from la, carbo "coal") is a with the C and 6. It is lic and —making four s available to form s. It belongs to group 14 of the periodic table. Carbon makes up only about 0.025 percent of Earth's crust. Three occur naturally, and being stable, while is a , decaying with a of about 5,730 years. Carbon is one of the . Carbon is the 15th , and the after , , and . Carbon's abundance, its unique diversity of s, and its unusual ability to form s at the temperatures commonly encountered on enables this element to serve as a common element of . It is the second most abundant element in the by mass (about 18.5%) after oxygen. The atoms of carbon can bond together in diverse ways, resulting in various . The best known allotropes are , , and . The of carbon vary widely with the allotropic form. For example, graphite is and black while diamond is highly . Graphite is soft enough to form a streak on paper (hence its name, from the verb "γράφειν" which means "to write"), while diamond is the naturally occurring material known. Graphite is a good while diamond has a low . Under normal conditions, diamond, s, and have the highest of . All carbon allotropes are solids under normal conditions, with graphite being the most form at standard temperature and pressure. They are chemically resistant and require high temperature to react even with oxygen. The most common of carbon in s is +4, while +2 is found in and complexes. The largest sources of inorganic carbon are s, s and , but significant quantities occur in organic deposits of , , , and s. Carbon forms a vast number of , more than any other element, with almost ten million compounds described to date, and yet that number is but a fraction of the number of theoretically possible compounds under standard conditions. For this reason, carbon has often been referred to as the "king of the elements".


The include , one of the softest known substances, and , the hardest naturally occurring substance. It readily with other small s, including other carbon atoms, and is capable of forming multiple stable bonds with suitable multivalent atoms. Carbon is known to form almost ten million compounds, a large majority of all . Carbon also has the highest point of all elements. At it has no melting point, as its is at and , so it sublimes at about . Graphite is much more reactive than diamond at standard conditions, despite being more thermodynamically stable, as its delocalised is much more vulnerable to attack. For example, graphite can be oxidised by hot concentrated at standard conditions to , C6(CO2H)6, which preserves the hexagonal units of graphite while breaking up the larger structure.Greenwood and Earnshaw, pp. 289–292. Carbon sublimes in a carbon arc, which has a temperature of about 5800 K (5,530 °C or 9,980 °F). Thus, irrespective of its allotropic form, carbon remains solid at higher temperatures than the highest-melting-point metals such as or . Although thermodynamically prone to , carbon resists oxidation more effectively than elements such as and , which are weaker reducing agents at room temperature. Carbon is the sixth element, with a ground-state of 1s22s22p2, of which the four outer electrons are s. Its first four ionisation energies, 1086.5, 2352.6, 4620.5 and 6222.7 kJ/mol, are much higher than those of the heavier group-14 elements. The electronegativity of carbon is 2.5, significantly higher than the heavier group-14 elements (1.8–1.9), but close to most of the nearby nonmetals, as well as some of the second- and third-row s. Carbon's are normally taken as 77.2 pm (C−C), 66.7 pm (C=C) and 60.3 pm (C≡C), although these may vary depending on coordination number and what the carbon is bonded to. In general, covalent radius decreases with lower coordination number and higher bond order.Greenwood and Earnshaw, pp. 276–8. Carbon compounds form the basis of all known life on , and the provides some of the energy produced by the and other s. Although it forms an extraordinary variety of compounds, most forms of carbon are comparatively unreactive under normal conditions. At standard temperature and pressure, it resists all but the strongest oxidizers. It does not react with , , or any . At elevated temperatures, carbon reacts with to form and will rob oxygen from metal oxides to leave the elemental metal. This is used in the iron and steel industry to iron and to control the carbon content of : : + 4 C → 3 Fe + 4 CO Carbon monoxide can be recycled to smelt even more iron: : + 4 CO → 3 Fe + 4 with to form and with steam in the coal-gas reaction: :C + HO → CO + H. Carbon combines with some metals at high temperatures to form metallic carbides, such as the iron carbide in steel and , widely used as an and for making hard tips for cutting tools. The system of carbon allotropes spans a range of extremes:


is a very short-lived species and, therefore, carbon is stabilized in various multi-atomic structures with diverse molecular configurations called . The three relatively well-known allotropes of carbon are , , and . Once considered exotic, s are nowadays commonly synthesized and used in research; they include s, s, s and . Several other exotic allotropes have also been discovered, such as , , and (carbyne). is a two-dimensional sheet of carbon with the atoms arranged in a hexagonal lattice. As of 2009, graphene appears to be the strongest material ever tested. The process of separating it from will require some further technological development before it is economical for industrial processes. If successful, graphene could be used in the construction of a . It could also be used to safely store hydrogen for use in a hydrogen based engine in cars. The form is an assortment of carbon atoms in a non-crystalline, irregular, glassy state, not held in a crystalline macrostructure. It is present as a powder, and is the main constituent of substances such as , () and . At normal pressures, carbon takes the form of graphite, in which each atom is bonded trigonally to three others in a plane composed of fused al rings, just like those in s. The resulting network is 2-dimensional, and the resulting flat sheets are stacked and loosely bonded through weak s. This gives graphite its softness and its properties (the sheets slip easily past one another). Because of the delocalization of one of the outer electrons of each atom to form a , graphite conducts , but only in the plane of each sheet. This results in a lower bulk for carbon than for most s. The delocalization also accounts for the energetic stability of graphite over diamond at room temperature. At very high pressures, carbon forms the more compact allotrope, , having nearly twice the density of graphite. Here, each atom is bonded to four others, forming a 3-dimensional network of puckered six-membered rings of atoms. Diamond has the same as and , and because of the strength of the carbon-carbon , it is the hardest naturally occurring substance measured by . Contrary to the popular belief that ''"diamonds are forever"'', they are thermodynamically unstable (Δf''G''°(diamond, 298 K) = 2.9 kJ/mol) under normal conditions (298 K, 105 Pa) and transform into . Due to a high activation energy barrier, the transition into graphite is so slow at normal temperature that it is unnoticeable. The bottom left corner of the phase diagram for carbon has not been scrutinized experimentally. Although a computational study employing methods reached the conclusion that as and , diamond becomes more stable than graphite by approximately 1.1 kJ/mol, more recent and definitive experimental and computational studies show that graphite is more stable than diamond for , without applied pressure, by 2.7 kJ/mol at ''T'' = 0 K and 3.2 kJ/mol at ''T'' = 298.15 K. Under some conditions, carbon crystallizes as , a lattice with all atoms covalently bonded and properties similar to those of diamond. s are a synthetic crystalline formation with a graphite-like structure, but in place of flat only, some of the cells of which fullerenes are formed may be pentagons, nonplanar hexagons, or even heptagons of carbon atoms. The sheets are thus warped into spheres, ellipses, or cylinders. The properties of fullerenes (split into buckyballs, buckytubes, and nanobuds) have not yet been fully analyzed and represent an intense area of research in . The names ''fullerene'' and ''buckyball'' are given after , popularizer of s, which resemble the structure of fullerenes. The buckyballs are fairly large molecules formed completely of carbon bonded trigonally, forming s (the best-known and simplest is the soccerball-shaped C ). Carbon nanotubes (buckytubes) are structurally similar to buckyballs, except that each atom is bonded trigonally in a curved sheet that forms a hollow . Nanobuds were first reported in 2007 and are hybrid buckytube/buckyball materials (buckyballs are covalently bonded to the outer wall of a nanotube) that combine the properties of both in a single structure. Of the other discovered allotropes, is a allotrope discovered in 1997. It consists of a low-density cluster-assembly of carbon atoms strung together in a loose three-dimensional web, in which the atoms are bonded trigonally in six- and seven-membered rings. It is among the lightest known solids, with a density of about 2 kg/m. Similarly, contains a high proportion of closed , but contrary to normal graphite, the graphitic layers are not stacked like pages in a book, but have a more random arrangement. has the chemical structure −(C:::C)''n''−. Carbon in this modification is linear with ''sp'' , and is a with alternating single and triple bonds. This carbyne is of considerable interest to as its is 40 times that of the hardest known material – diamond. In 2015, a team at the announced the development of another allotrope they have dubbed , created by a high energy low duration laser pulse on amorphous carbon dust. Q-carbon is reported to exhibit ferromagnetism, , and a hardness superior to diamonds. In the vapor phase, some of the carbon is in the form of (). When excited, this gas glows green.


Carbon is the in the by mass after hydrogen, helium, and oxygen. In July 2020, astronomers reported evidence that carbon was formed mainly in , particularly those bigger than two solar masses. Carbon is abundant in the , s, s, and in the of most s. Some s contain microscopic diamonds that were formed when the was still a . Microscopic diamonds may also be formed by the intense pressure and high temperature at the sites of meteorite impacts. In 2014 announced
greatly upgraded database
for tracking (PAHs) in the . More than 20% of the carbon in the universe may be associated with PAHs, complex compounds of carbon and hydrogen without oxygen. These compounds figure in the where they are hypothesized to have a role in and formation of . PAHs seem to have been formed "a couple of billion years" after the , are widespread throughout the universe, and are associated with and . It has been estimated that the solid earth as a whole contains 730 of carbon, with 2000 ppm in the core and 120 ppm in the combined mantle and crust. Since the mass of the earth is , this would imply 4360 million s of carbon. This is much more than the amount of carbon in the oceans or atmosphere (below). In combination with in , carbon is found in the Earth's atmosphere (approximately 900 gigatonnes of carbon — each ppm corresponds to 2.13 Gt) and dissolved in all water bodies (approximately 36,000 gigatonnes of carbon). Carbon in the has been estimated at 550 gigatonnes but with a large uncertainty, due mostly to a huge uncertainty in the amount of terrestrial deep . (such as , , and ) contain carbon as well. amount to around 900 gigatonnes with perhaps 18,000 Gt of resources. are around 150 gigatonnes. Proven sources of natural gas are about (containing about 105 gigatonnes of carbon), but studies estimate another of "unconventional" deposits such as , representing about 540 gigatonnes of carbon. Carbon is also found in in polar regions and under the seas. Various estimates put this carbon between 500, 2500 , or 3,000 Gt. In the past, quantities of hydrocarbons were greater. According to one source, in the period from 1751 to 2008 about 347 gigatonnes of carbon were released as carbon dioxide to the atmosphere from burning of fossil fuels. Another source puts the amount added to the atmosphere for the period since 1750 at 879 Gt, and the total going to the atmosphere, sea, and land (such as ) at almost 2,000 Gt. Carbon is a constituent (about 12% by mass) of the very large masses of rock (, , and so on). is very rich in carbon ( contains 92–98%) and is the largest commercial source of mineral carbon, accounting for 4,000 gigatonnes or 80% of . As for individual carbon allotropes, graphite is found in large quantities in the (mostly in and ), , , , and . Natural diamonds occur in the rock , found in ancient "necks", or "pipes". Most diamond deposits are in , notably in , , , the , and . Diamond deposits have also been found in , , the Russian , , and in Northern and Western . Diamonds are now also being recovered from the ocean floor off the . Diamonds are found naturally, but about 30% of all industrial diamonds used in the U.S. are now manufactured. Carbon-14 is formed in upper layers of the troposphere and the stratosphere at altitudes of 9–15 km by a reaction that is precipitated by s. s are produced that collide with the nuclei of nitrogen-14, forming carbon-14 and a proton. As such, of atmospheric carbon dioxide contains carbon-14. Carbon-rich asteroids are relatively preponderant in the outer parts of the in our . These asteroids have not yet been directly sampled by scientists. The asteroids can be used in hypothetical , which may be possible in the future, but is currently technologically impossible.


s of carbon are that contain six s plus a number of s (varying from 2 to 16). Carbon has two stable, naturally occurring s. The isotope (C) forms 98.93% of the carbon on Earth, while (C) forms the remaining 1.07%. The concentration of C is further increased in biological materials because biochemical reactions discriminate against C. In 1961, the (IUPAC) adopted the isotope as the basis for s. Identification of carbon in (NMR) experiments is done with the isotope C. (C) is a naturally occurring , created in the (lower and upper ) by interaction of with s. It is found in trace amounts on Earth of 1 part per (0.0000000001%) or more, mostly confined to the atmosphere and superficial deposits, particularly of and other organic materials. This isotope decays by 0.158 MeV . Because of its relatively short of 5730 years, C is virtually absent in ancient rocks. The amount of C in the and in living organisms is almost constant, but decreases predictably in their bodies after death. This principle is used in , invented in 1949, which has been used extensively to determine the age of carbonaceous materials with ages up to about 40,000 years. There are 15 known isotopes of carbon and the shortest-lived of these is C which decays through and and has a half-life of 1.98739 × 10 s. The exotic C exhibits a , which means its is appreciably larger than would be expected if the were a of constant .

Formation in stars

Formation of the carbon atomic nucleus occurs within a or star through the . This requires a nearly simultaneous collision of three s ( nuclei), as the products of further reactions of helium with hydrogen or another helium nucleus produce and respectively, both of which are highly unstable and decay almost instantly back into smaller nuclei. The triple-alpha process happens in conditions of temperatures over 100 megakelvins and helium concentration that the rapid expansion and cooling of the early universe prohibited, and therefore no significant carbon was created during the . According to current physical cosmology theory, carbon is formed in the interiors of stars on the . When massive stars die as supernova, the carbon is scattered into space as dust. This dust becomes component material for the formation of the systems with accreted planets. The is one such star system with an abundance of carbon, enabling the existence of life as we know it. The is an additional hydrogen fusion mechanism that powers stars, wherein carbon operates as a . Rotational transitions of various isotopic forms of carbon monoxide (for example, CO, CO, and CO) are detectable in the wavelength range, and are used in the study of in .

Carbon cycle

Under terrestrial conditions, conversion of one element to another is very rare. Therefore, the amount of carbon on Earth is effectively constant. Thus, processes that use carbon must obtain it from somewhere and dispose of it somewhere else. The paths of carbon in the environment form the . For example, plants draw from the atmosphere (or seawater) and build it into biomass, as in the , a process of . Some of this biomass is eaten by animals, while some carbon is exhaled by animals as carbon dioxide. The carbon cycle is considerably more complicated than this short loop; for example, some carbon dioxide is dissolved in the oceans; if bacteria do not consume it, dead plant or animal matter may become or , which releases carbon when burned.


Organic compounds

Carbon can form very long chains of interconnecting s, a property that is called . Carbon-carbon bonds are strong and stable. Through catenation, carbon forms a countless number of compounds. A tally of unique compounds shows that more contain carbon than do not. A similar claim can be made for hydrogen because most organic compounds contain hydrogen chemically bonded to carbon or another common element like oxygen or nitrogen. The simplest form of an organic molecule is the —a large family of s that are composed of atoms bonded to a chain of carbon atoms. A hydrocarbon backbone can be substituted by other atoms, known as s. Common heteroatoms that appear in organic compounds include oxygen, nitrogen, sulfur, phosphorus, and the nonradioactive halogens, as well as the metals lithium and magnesium. Organic compounds containing bonds to metal are known as organometallic compounds (''see below''). Certain groupings of atoms, often including heteroatoms, recur in large numbers of organic compounds. These collections, known as , confer common reactivity patterns and allow for the systematic study and categorization of organic compounds. Chain length, shape and functional groups all affect the properties of organic molecules. In most stable compounds of carbon (and nearly all stable ''organic'' compounds), carbon obeys the and is ''tetravalent'', meaning that a carbon atom forms a total of four covalent bonds (which may include double and triple bonds). Exceptions include a small number of stabilized ''carbocations'' (three bonds, positive charge), ''radicals'' (three bonds, neutral), ''carbanions'' (three bonds, negative charge) and ''carbenes'' (two bonds, neutral), although these species are much more likely to be encountered as unstable, reactive intermediates. Carbon occurs in all known life and is the basis of . When united with , it forms various hydrocarbons that are important to industry as s, s, s, as chemical feedstock for the manufacture of s and s, and as s. When combined with oxygen and hydrogen, carbon can form many groups of important biological compounds including s, s, s, s, s, and aromatic s, s and s. With it forms s, and with the addition of sulfur also it forms s, s, and products. With the addition of phosphorus to these other elements, it forms and , the chemical-code carriers of life, and (ATP), the most important energy-transfer molecule in all living cells.

Inorganic compounds

Commonly carbon-containing compounds which are associated with minerals or which do not contain bonds to the other carbon atoms, halogens, or hydrogen, are treated separately from classical ; the definition is not rigid, and the classification of some compounds can vary from author to author (see reference articles above). Among these are the simple oxides of carbon. The most prominent oxide is (). This was once the principal constituent of the , but is a minor component of the today. Dissolved in , it forms (), but as most compounds with multiple single-bonded oxygens on a single carbon it is unstable. Through this intermediate, though, resonance-stabilized s are produced. Some important minerals are carbonates, notably . () is similar. Nevertheless, due to its physical properties and its association with organic synthesis, carbon disulfide is sometimes classified as an ''organic'' solvent. The other common oxide is (CO). It is formed by incomplete combustion, and is a colorless, odorless gas. The molecules each contain a triple bond and are fairly , resulting in a tendency to bind permanently to hemoglobin molecules, displacing oxygen, which has a lower binding affinity. (CN), has a similar structure, but behaves much like a ion (). For example, it can form the nitride molecule ((CN)), similar to diatomic halides. Likewise, the heavier analog of cyanide, (CP), is also considered inorganic, though most simple derivatives are highly unstable. Other uncommon oxides are (), the unstable (CO), (CO), (CO), (CO), and (CO). However, mellitic anhydride is the triple acyl anhydride of mellitic acid; moreover, it contains a benzene ring. Thus, many chemists consider it to be organic. With reactive s, such as , carbon forms either s (C) or s () to form alloys with high melting points. These anions are also associated with and , both very weak s. With an electronegativity of 2.5, carbon prefers to form s. A few carbides are covalent lattices, like (SiC), which resembles . Nevertheless, even the most polar and salt-like of carbides are not completely ionic compounds.Greenwood and Earnshaw, pp. 297–301

Organometallic compounds

Organometallic compounds by definition contain at least one carbon-metal covalent bond. A wide range of such compounds exist; major classes include simple alkyl-metal compounds (for example, ), η-alkene compounds (for example, ), and η-allyl compounds (for example, ); s containing cyclopentadienyl ligands (for example, ); and es. Many s and exist (for example, and ); some workers consider metal carbonyl and cyanide complexes without other carbon ligands to be purely inorganic, and not organometallic. However, most organometallic chemists consider metal complexes with any carbon ligand, even 'inorganic carbon' (e.g., carbonyls, cyanides, and certain types of carbides and acetylides) to be organometallic in nature. Metal complexes containing organic ligands without a carbon-metal covalent bond (e.g., metal carboxylates) are termed ''metalorganic'' compounds. While carbon is understood to strongly prefer formation of four covalent bonds, other exotic bonding schemes are also known. s are highly stable dodecahedral derivatives of the 12H12sup>2- unit, with one BH replaced with a CH+. Thus, the carbon is bonded to five boron atoms and one hydrogen atom. The cation PhPAu)Ccontains an octahedral carbon bound to six phosphine-gold fragments. This phenomenon has been attributed to the of the gold ligands, which provide additional stabilization of an otherwise labile species. In nature, the iron-molybdenum cofactor () responsible for microbial likewise has an octahedral carbon center (formally a carbide, C(-IV)) bonded to six iron atoms. In 2016, it was confirmed that, in line with earlier theoretical predictions, the contains a carbon atom with six bonds. More specifically, the dication could be described structurally by the formulation eC(η5-C5Me5)sup>2+, making it an "organic " in which a MeC3+ fragment is bonded to a η5-C5Me5 fragment through all five of the carbons of the ring. It is important to note that in the cases above, each of the bonds to carbon contain less than two formal electron pairs. Thus, the formal electron count of these species does not exceed an octet. This makes them hypercoordinate but not hypervalent. Even in cases of alleged 10-C-5 species (that is, a carbon with five ligands and a formal electron count of ten), as reported by Akiba and co-workers, electronic structure calculations conclude that the electron population around carbon is still less than eight, as is true for other compounds featuring four-electron ing.

History and etymology

The name ''carbon'' comes from the ''carbo'' for coal and charcoal, whence also comes the ''charbon'', meaning charcoal. In , and , the names for carbon are ''Kohlenstoff'', ''koolstof'' and ''kulstof'' respectively, all literally meaning -substance. Carbon was discovered in prehistory and was known in the forms of and to the earliest s. Diamonds were known probably as early as 2500 BCE in China, while carbon in the form of was made around Roman times by the same chemistry as it is today, by heating wood in a covered with to exclude air. In 1722, demonstrated that iron was transformed into steel through the absorption of some substance, now known to be carbon. In 1772, showed that diamonds are a form of carbon; when he burned samples of charcoal and diamond and found that neither produced any water and that both released the same amount of per . In 1779, showed that graphite, which had been thought of as a form of , was instead identical with charcoal but with a small admixture of iron, and that it gave "aerial acid" (his name for carbon dioxide) when oxidized with nitric acid. In 1786, the French scientists , and C. A. Vandermonde confirmed that graphite was mostly carbon by oxidizing it in oxygen in much the same way Lavoisier had done with diamond. Some iron again was left, which the French scientists thought was necessary to the graphite structure. In their publication they proposed the name ''carbone'' (Latin ''carbonum'') for the element in graphite which was given off as a gas upon burning graphite. Antoine Lavoisier then listed carbon as an in his 1789 textbook. A new of carbon, , that was discovered in 1985 includes d forms such as s and . Their discoverers – , and  – received the in Chemistry in 1996. The resulting renewed interest in new forms lead to the discovery of further exotic allotropes, including , and the realization that "" is not strictly .



Commercially viable natural deposits of graphite occur in many parts of the world, but the most important sources economically are in , , and . Graphite deposits are of origin, found in association with , and s in schists, es and metamorphosed s and as or , sometimes of a metre or more in thickness. Deposits of graphite in , , were at first of sufficient size and purity that, until the 19th century, s were made simply by sawing blocks of natural graphite into strips before encasing the strips in wood. Today, smaller deposits of graphite are obtained by crushing the parent rock and floating the lighter graphite out on water.USGS Minerals Yearbook: Graphite, 2009
and Graphite: Mineral Commodity Summaries 2011
There are three types of natural graphite—amorphous, flake or crystalline flake, and vein or lump. Amorphous graphite is the lowest quality and most abundant. Contrary to science, in industry "amorphous" refers to very small crystal size rather than complete lack of crystal structure. Amorphous is used for lower value graphite products and is the lowest priced graphite. Large amorphous graphite deposits are found in China, Europe, Mexico and the United States. Flake graphite is less common and of higher quality than amorphous; it occurs as separate plates that crystallized in metamorphic rock. Flake graphite can be four times the price of amorphous. Good quality flakes can be processed into for many uses, such as s. The foremost deposits are found in Austria, Brazil, Canada, China, Germany and Madagascar. Vein or lump graphite is the rarest, most valuable, and highest quality type of natural graphite. It occurs in veins along intrusive contacts in solid lumps, and it is only commercially mined in Sri Lanka. According to the , world production of natural graphite was 1.1 million tonnes in 2010, to which China contributed 800,000 t, India 130,000 t, Brazil 76,000 t, North Korea 30,000 t and Canada 25,000 t. No natural graphite was reported mined in the United States, but 118,000 t of synthetic graphite with an estimated value of $998 million was produced in 2009.


The diamond supply chain is controlled by a limited number of powerful businesses, and is also highly concentrated in a small number of locations around the world (see figure). Only a very small fraction of the diamond ore consists of actual diamonds. The ore is crushed, during which care has to be taken in order to prevent larger diamonds from being destroyed in this process and subsequently the particles are sorted by density. Today, diamonds are located in the diamond-rich density fraction with the help of , after which the final sorting steps are done by hand. Before the use of s became commonplace, the separation was done with grease belts; diamonds have a stronger tendency to stick to grease than the other minerals in the ore. Historically diamonds were known to be found only in alluvial deposits in . discussion on alluvial diamonds in India and elsewhere as well as earliest finds India led the world in diamond production from the time of their discovery in approximately the 9th century BC Ball was a Geologist in British service. Chapter I, Page 1 to the mid-18th century AD, but the commercial potential of these sources had been exhausted by the late 18th century and at that time India was eclipsed by Brazil where the first non-Indian diamonds were found in 1725. Diamond production of primary deposits (kimberlites and lamproites) only started in the 1870s after the discovery of the diamond fields in South Africa. Production has increased over time and now an accumulated total of 4.5 billion carats have been mined since that date. About 20% of that amount has been mined in the last 5 years alone, and during the last ten years 9 new mines have started production while 4 more are waiting to be opened soon. Most of these mines are located in Canada, Zimbabwe, Angola, and one in Russia. In the United States, diamonds have been found in , and . In 2004, a startling discovery of a microscopic diamond in the United States led to the January 2008 bulk-sampling of in a remote part of . Today, most commercially viable diamond deposits are in , , and the . In 2005, Russia produced almost one-fifth of the global diamond output, reports the . Australia has the richest diamantiferous pipe with production reaching peak levels of per year in the 1990s. There are also commercial deposits being actively mined in the of , (mostly in ; for example, and ), Brazil, and in Northern and Western .


Carbon is essential to all known living systems, and without it life as we know it could not exist (see ). The major economic use of carbon other than food and wood is in the form of hydrocarbons, most notably the gas and (petroleum). is in by the to produce , , and other products. is a natural, carbon-containing polymer produced by plants in the form of , , , and . is used primarily for maintaining structure in plants. Commercially valuable carbon polymers of animal origin include , and . are made from synthetic carbon polymers, often with oxygen and nitrogen atoms included at regular intervals in the main polymer chain. The raw materials for many of these synthetic substances come from crude oil. The uses of carbon and its compounds are extremely varied. It can form with , of which the most common is . is combined with s to form the 'lead' used in s used for and . It is also used as a and a , as a molding material in manufacture, in for dry and in and , in for and as a in . is used as a drawing material in work, barbecue , , and in many other applications. Wood, coal and oil are used as for production of energy and . Gem quality is used in jewelry, and s are used in drilling, cutting and polishing tools for machining metals and stone. Plastics are made from fossil hydrocarbons, and , made by of synthetic s is used to reinforce plastics to form advanced, lightweight . is made by pyrolysis of extruded and stretched filaments of (PAN) and other organic substances. The crystallographic structure and mechanical properties of the fiber depend on the type of starting material, and on the subsequent processing. Carbon fibers made from PAN have structure resembling narrow filaments of graphite, but thermal processing may re-order the structure into a continuous rolled sheet. The result is fibers with higher than steel. is used as the black in , artist's oil paint and water colours, , automotive finishes, and . is also used as a in products such as tyres and in compounds. is used as an and in material in applications as diverse as , , and s, and in medicine to toxins, poisons, or gases from the . Carbon is used in at high temperatures. is used to reduce iron ore into iron (smelting). of steel is achieved by heating finished steel components in carbon powder. s of , , and , are among the hardest known materials, and are used as in cutting and grinding tools. Carbon compounds make up most of the materials used in clothing, such as natural and synthetic and , and almost all of the interior surfaces in the other than glass, stone and metal.


The industry falls into two categories: one dealing with gem-grade diamonds and the other, with industrial-grade diamonds. While a large trade in both types of diamonds exists, the two markets function dramatically differently. Unlike s such as or , gem diamonds do not trade as a : there is a substantial mark-up in the sale of diamonds, and there is not a very active market for resale of diamonds. Industrial diamonds are valued mostly for their hardness and heat conductivity, with the gemological qualities of clarity and color being mostly irrelevant. About 80% of mined diamonds (equal to about 100 million carats or 20 tonnes annually) are unsuitable for use as gemstones are relegated for industrial use (known as '')''. s, invented in the 1950s, found almost immediate industrial applications; 3 billion carats (600 s) of synthetic diamond is produced annually. The dominant industrial use of diamond is in cutting, drilling, grinding, and polishing. Most of these applications do not require large diamonds; in fact, most diamonds of gem-quality except for their small size can be used industrially. Diamonds are embedded in drill tips or saw blades, or ground into a powder for use in grinding and polishing applications. Specialized applications include use in laboratories as containment for (see ), high-performance , and limited use in specialized s. With the continuing advances in the production of synthetic diamonds, new applications are becoming feasible. Garnering much excitement is the possible use of diamond as a suitable for , and because of its exceptional heat conductance property, as a in .


Pure carbon has extremely low to humans and can be handled safely in the form of graphite or charcoal. It is resistant to dissolution or chemical attack, even in the acidic contents of the digestive tract. Consequently, once it enters into the body's tissues it is likely to remain there indefinitely. was probably one of the first pigments to be used for ing, and was found to have carbon tattoos that survived during his life and for 5200 years after his death. Inhalation of coal dust or soot (carbon black) in large quantities can be dangerous, irritating lung tissues and causing the congestive disease, . Diamond dust used as an abrasive can be harmful if ingested or inhaled. Microparticles of carbon are produced in diesel engine exhaust fumes, and may accumulate in the lungs. In these examples, the harm may result from contaminants (e.g., organic chemicals, heavy metals) rather than from the carbon itself. Carbon generally has low toxicity to ; but carbon nanoparticles are deadly to '. Carbon may burn vigorously and brightly in the presence of air at high temperatures. Large accumulations of coal, which have remained inert for hundreds of millions of years in the absence of oxygen, may when exposed to air in coal mine waste tips, ship cargo holds and coal bunkers, and storage dumps. In where graphite is used as a , accumulation of followed by a sudden, spontaneous release may occur. to at least 250 °C can release the energy safely, although in the the procedure went wrong, causing other reactor materials to combust. The great variety of carbon compounds include such lethal poisons as , the from seeds of the ', (CN), and ; and such essentials to life as and .

See also

* * * * * *




External links


at ' (University of Nottingham)
Carbon on Britannica

Carbon—Super Stuff. Animation with sound and interactive 3D-models.
{{Authority control Reducing agents