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carbon [Lat.,=charcoal], nonmetallic chemical element; symbol C; at. no. 6; interval in which at. wt. ranges 12.0096–12.0116; m.p. about 3,550℃; graphite sublimes about 3,375℃; b.p. 4,827℃; sp. gr. 1.8–2.1 (amorphous), 1.9–2.3 (graphite), 3.15–3.53 (diamond); valence +2, +3, +4, or −4.
Properties and Isotopes
Carbon is found free in nature in at least four distinct forms (see allotropy). One form, graphite, is a very soft, dark gray or black, lustrous material with either a hexagonal or rhombohedral crystalline structure. Diamond, a second crystalline form, is the hardest substance known. In a third form, the so-called amorphous carbon, the element occurs partly free and partly combined with other elements; charcoal, coal, coke, lampblack, peat, and lignite are some sources of amorphous carbon. A fourth form contains the fullerenes, stable molecules consisting of carbon atoms that arrange themselves into 12 pentagonal faces and any number greater than 1 of hexagonal faces. The most prominent of the fullerenes is buckminsterfullerene, a spheroidal molecule, resembling a soccer ball, consisting of 60 carbon atoms. A fifth form, “white” carbon, is believed to exist. Carbon has the capacity to act chemically both as a metal and as a nonmetal. It is a constituent of all organic matter.
Carbon has 13 known isotopes, which have from 2 to 14 neutrons in the nucleus and mass numbers from 8 to 20. Carbon-12 was chosen by IUPAC in 1961 as the basis for atomic weights; it is assigned an atomic mass of exactly 12 atomic mass units. Carbon-13 absorbs radio waves and is used in nuclear magnetic resonance spectrometry to study organic compounds. Carbon-14, which has a half-life of 5,730 years, is a naturally occurring isotope that can also be produced in a nuclear reactor. It is used extensively as a research tool in tracer studies; a compound synthesized with carbon-14 is said to be “tagged” and can be traced through a chemical or biochemical reaction. Carbon-14 has been used in the study of such problems as utilization of foods in animal nutrition, catalytic petroleum processes, photosynthesis, and the mechanism of aging in steel. It is also used for determining the age of archaeological specimens (see dating).
Natural Occurrence and Uses
See P. L. Walker, Jr., and P. A. Thrower, ed., Chemistry and Physics of Carbon (11 vol., 1966–74); H. O. Pierson, Handbook of Carbon, Graphite, Diamond, and Fullerenes: Properties, Processing, and Applications (1993).
C, a chemical element of group IV of Mendeleev’s periodic system. Atomic number, 6; atomic weight, 12.011. The two known stable isotopes of carbon are 12C (98.892 percent) and13C (1.108 percent). The most important radioisotope is 14C, with a half-life of 5.6 × 103 years. Small amounts of 14C (approximately 2 × 10–10 percent by weight) are constantly being formed in the upper layers of the atmosphere by the action of cosmic neutron radiation on 14N. The age of biological remains is determined by the specific activity of 14C. This isotope is also commonly used as an isotope tracer.
Historical survey. Carbon has been known since antiquity. Charcoal was used to reduce metals from ores, and diamonds were valued as gems. The use of graphite in the production of crucibles and pencils began much later.
In 1778, K. Scheele discovered that carbon dioxide is liberated when graphite is heated with saltpeter, the same result as that obtained by heating coal with saltpeter. The chemical composition of diamond was established by A. Lavoisier, who in 1772 studied the combustion of diamond in air, and by S. Tennant, who in 1797 showed that equal amounts of diamond and coal yield, upon oxidation, equal amounts of carbon dioxide. Carbon was recognized as a chemical element in 1789 by Lavoisier. Its name comes from the Latin word carbo, meaning coal.
Distribution in nature. The average content of carbon in the earth’s crust is 2.3 × 10–2 percent by weight (1 × 10–2 percent in ultrabasic rock, 1 × 10–2 percent in basic rock, 2 × 10–2 percent in intermediate rock, and 3 × 10–2 percent in acidic rock). Carbon accumulates in the upper portion of the earth’s crust (biosphere). The content of carbon is 18 percent in living matter, 50 percent in wood, 80 percent in coal, 85 percent in petroleum, and 96 percent in anthracite. A significant portion of the carbon in the lithosphere is concentrated in limestones and dolomites.
There are 112 minerals of carbon as such. The number of organic compounds of carbon, namely, hydrocarbons and hydrocarbon derivatives, is exceedingly large.
The accumulation of carbon in the earth’s crust results in the accumulation of many other elements that are sorbed by organic matter and precipitated in the form of, for example, insoluble carbonates. Carbon dioxide and carbonic acid have important geochemical roles in the earth’s crust. Enormous amounts of carbon dioxide are liberated as a result of volcanic activity; historically, this activity was the principal source of carbon in the biosphere.
Mankind has extracted carbon from the earth in quantities that are extremely large in relation to the element’s average content in the crust. The carbon is extracted in the form of coal, petroleum, and natural gas, substances that are major sources of energy.
The carbon cycle has enormous geochemical significance. The biological role of the element is discussed below.
Carbon is also widely distributed in space; it is the sun’s fourth most abundant element, after hydrogen, helium, and oxygen.
Physical and chemical properties. The four known crystalline modifications of carbon are graphite, diamond, artificial graphite, and lonsdaleite. Graphite is a gray-black, opaque, scaly, very soft substance that has a metallic luster and is greasy to the touch. It consists of hexagonal crystals, with a = 2.462 angstroms (Å) and c = 6.701 Å. At room temperature and normal pressure (0.1 meganewton per sq m [MN/m2], or 1 kilogram-force per sq cm [kgf/cm2]), graphite is thermodynamically stable.
Diamond is a very hard crystalline substance. Its crystals have a cubic face-centered lattice, with a = 3.560 Å. At room temperature and normal pressure, diamond is metastable. Diamond converts into graphite to a marked extent at temperatures above 1400°C in a vacuum or an inert atmosphere. Graphite sublimes at a temperature of approximately 3700°C at atmospheric pressure. Liquid carbon can be obtained at pressures greater than 10.5 MN/m2 (105 kgf/cm2) and temperatures above 3700°C. Solid carbon can also have a disordered structure. This amorphous form, which includes coke, carbon black, and charcoal, is not an independent modification of carbon, and its structure is basically that of finely crystalline graphite. Heating several varieties of amorphous carbon above 1500°-1600°C without access to the air induces conversion into graphite. The physical properties of amorphous carbon depend to a great extent on the degree of dispersion of the particles and on the presence of impurities. The density, heat capacity, heat conductivity, and electric conductivity of amorphous carbon are always higher than those of graphite. Artificial graphite is a finely crystalline black powder with a density of 1.9–2 g/cm3; it is made up of long chains of carbon atoms arranged in parallel fashion. Lonsdaleite was discovered in meteorites and can be obtained artificially; its structure and properties have not yet been finally established.
The configuration of the outer electron subshells of the carbon atom is 2s22p2. Carbon characteristically forms four covalent bonds as a result of the excitation of the outer shell to the 2sp3 state. Thus, carbon is equally capable of accepting and donating electrons. The chemical bonds of carbon may involve sp3, sp2, or sp hybrid orbitals, which correspond to coordination numbers of 4, 3, and 2. The number of valence electrons of carbon and the number of valence orbitals are identical, which is one of the reasons for the stability of the bonds between carbon atoms.
The unique capacity of carbon atoms to form bonds with one another resulting in long stable chains and rings accounts for the enormous number and great variety of carbon compounds, the study of which is the province of organic chemistry.
Oxidation states of —4, +2 and +4 are observed in carbon compounds. The element’s atomic radius is 0.77 A, and covalent radii of 0.77 Å, 0.67 Å, and 0.60 Å correspond to, respectively, single, double, and triple bonds. The ionic radius of C4- is 2.60 Å, while that of C4+ is 0.20 Å. Under ordinary conditions, carbon is chemically inert, but at high temperatures it combines with many elements and displays strong reducing properties. The chemical reactivity of carbon decreases in the order amorphous carbon > graphite > diamond; the reactions of these three forms of carbon with atmospheric oxygen (combustion) occur, respectively, at temperatures above 300°-500°C, 600°-700°C, and 850°-1000°C and result in the formation of carbon dioxide (CO2) and carbon monoxide (CO).
Carbon dioxide dissolves in water to form carbonic acid. In 1906, O. Diels obtained carbon suboxide (C302). All forms of carbon are stable toward alkalies and acids; they are slowly oxidized only by very strong oxidizing agents, such as chromium mixtures or mixtures of concentrated HN03 and KC103. Amorphous carbon reacts with fluorine at room temperature, while graphite and diamond react with this element only upon heating. The direct combination of carbon with chlorine occurs in an electric arc. Since carbon does not react with bromine and iodine, many carbon halides are obtained indirectly. Carbonyl chloride (COCl2, phosgene) is the best known of the carbon oxyhalides, which have the general formula COX2, × being the halogen. While hydrogen does not react with diamond, it does react with graphite and amorphous carbon at high temperatures in the presence of catalysts (Ni, Pt); the main product is methane (CH4) at 600°-1000°C and acetylene (C2H2) at 1500°-2000°C. Other hydrocarbons, for example, ethane (C2H6) and benzene (C6H6), may also be present in the products.
Sulfur begins to react with amorphous carbon and graphite at 700°-800°C; with diamond, however, the temperature must be 900°-1000°C. In all cases, the product is carbon disulfide (CS2). Other sulfur-containing carbon compounds, such as carbon monosulfide (CS), carbon thiosuboxide (C3S2), carbonyl sulfide (COS), and thiophosgene (CSC12), are obtained indirectly. Thio-carbonates, the salts of the weak acid thiocarbonic acid, are formed from the reaction of CS2 with metal sulfides. The reaction of carbon with nitrogen to yield cyanogen [(CN)2] accompanies an electric discharge between carbon electrodes in a nitrogen atmosphere. Of the nitrogen-containing carbon compounds, hydrogen cyanide (HCN) and its numerous derivatives, namely, cyanides, cyanogen halides, and nitriles, have the greatest practical importance.
At temperatures above 1000°C, carbon reacts with many metals to yield carbides. Upon heating, all forms of carbon reduce metal oxides, forming free metals (Zn, Cd, Cu, Pb) or carbides (CaC2, Mo2C, WC, TaC). Carbon reacts at temperatures above 600°-800°C with steam and carbon dioxide. A distinguishing feature of graphite is the capacity to react with alkali metals and halides upon moderate heating (to 300°-400oC) to yield inclusion compounds of the type C8M, C24M, or C8X, where × is a halogen and M is a metal. Inclusion compounds of graphite and HN03, H2SO4, and FeCl3 are known, for example, graphite bisulfate (C24SO4H2). All forms of carbon are insoluble in ordinary inorganic and organic solvents but soluble in some molten metals, among them, Fe, Ni, and Co.
Economic importance. The importance of carbon in the economy derives from the element’s presence in fossil fuels, the primary source for more than 90 percent of the energy consumed by the world. Fossil fuels will retain their dominant position in the next few decades despite the rapid development of nuclear power. Only approximately 10 percent of the fossil fuels extracted is used as a raw material in industrial organic synthesis, petrochemical synthesis, and the production of plastics.
B. A. POPOVKIN
In organisms. Carbon is the most important biological element, providing the basis for life on the earth and serving as a structural unit for the enormous number of organic compounds that make up an organism and participate in the organism’s vital activities. Examples are provided by biopolymers and by the many low-molecular-weight biologically active substances, such as vitamins, hormones, and mediators. A significant part of the energy required by an organism is produced by the oxidation of carbon in the cells. The origin of life on the earth is considered by modern science to have been a complex process of evolution of carbon compounds.
The unique role of carbon in living things derives from the element’s properties, the aggregate of which is not possessed by any other element of the periodic system. Strong chemical bonds are formed between carbon atoms, as well as between carbon and other elements, but these bonds (single, double, triple) may be cleaved under relatively mild physiological conditions. The capacity of carbon to form four equivalent valence bonds with other carbon atoms permits the construction of different types of carbon skeletons, including linear, branched-chain, and cyclic structures. It is significant that just three elements, namely, carbon, oxygen, and hydrogen, constitute 98 percent of the total mass of living organisms. This fact makes possible a certain economy in living things; with the practically limitless structural diversity of carbon compounds, the existence of just a few types of chemical bonds permits a significant reduction in the number of enzymes required for the breakdown and synthesis of organic compounds. The special features of the carbon atom’s structure account for the various types of isomerism in organic compounds. (The tendency toward optical isomerism proved to be decisive in the biochemical evolution of amino acids, carbohydrates, and certain alkaloids.)
According to the generally accepted hypothesis of A. I. Opa-rin, the first organic compounds on earth were abiological. The sources of carbon were the methane (CH4) and hydrogen cyanide (HCN) present in the earth’s primordial atmosphere. With the appearance of life, the only source of inorganic carbon, through which all the organic matter in the biosphere was formed, was carbon dioxide (CO2); this compound was present in the atmosphere and was dissolved in the earth’s waters in HCO3– form.
The most powerful mechanism for the assimilation of carbon (in the form of CO2) is photosynthesis, which is everywhere carried out by green plants. (Approximately 100 billion tons of CO2are assimilated annually.) Another method of assimilating CO2, which evolved earlier, is chemosynthesis. Here, chemosynthetic bacteria use, rather than radiant energy from the sun, the energy derived from the oxidation of inorganic compounds. The carbon contained in the food of most animals is in the form of organic compounds. Depending on the method of assimilating organic compounds, a distinction is made between autotrophic and heterotrophic organisms. The possibility of using those microorganisms whose only source of carbon is petroleum hydrocarbons for the biosynthesis of protein and other nutrients is presently the focus of much attention in modern science and technology.
The content of carbon in living organisms, as a percentage of dry mass, is 34.5–40 percent in aquatic plants and animals, 45.4–46.5 percent in terrestrial plants and animals, and 54 percent in bacteria. The vital activities of an organism, especially tissue respiration, are accompanied by the oxidative decomposition of organic compounds and the liberation of CO2 into the external medium. Carbon is also liberated in more complex metabolic end products. As plants and animals die, a portion of the carbon is reconverted into CO2 as a result of decay processes carried out by microorganisms. Thus, a carbon cycle can be observed in nature. A significant portion of carbon is mineralized, forming deposits of, for example, coal, petroleum, and limestone. While CO2functions mainly as a source of carbon, it also, when dissolved in water and biological fluids, helps to maintain the optimal level of acidity for an organism’s vital activities. In the form of CaCO3, carbon forms the outer skeleton of many invertebrates, for example, the shells of mollusks; it is also present in corals and the eggshells of birds. Such carbon compounds as HCN, CO, and CC14 predominated in the earth’s atmosphere in the period before the appearance of life; subsequently, in the course of biological evolution, these compounds become strong antimetabolites.
In addition to the stable isotopes of carbon, radioactive l4C is widespread in nature. Its content in humans is approximately 1 microcurie. Many great advances in the study of metabolism and of the carbon cycle in nature have been linked to the use of carbon isotopes in biological and medical research. Thus, radiocarbon marking has been used to prove the possibility of the fixation of H14CO~3 by plants and animal tissues, establish the sequence of reactions in photosynthesis, study the metabolism of amino acids, and trace the biosynthetic pathways of biologically active compounds. The use of 14C has facilitated advances in molecular biology in the study of the mechanisms of protein biosynthesis and the transfer of genetic information. The evaluation of the age of carbon-containing organic remains through the determination of the specific activity of, 4C is particularly useful in paleontology and archaeology.
N. N. CHERNOV
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