carbon

Dictionary: car·bon   (kär'bən)

1. (Symbol C) A naturally abundant nonmetallic element that occurs in many inorganic and in all organic compounds, exists freely as graphite and diamond and as a constituent of coal, limestone, and petroleum, and is capable of chemical self-bonding to form an enormous number of chemically, biologically, and commercially important molecules. Atomic number 6; atomic weight 12.011; sublimation point above 3,500°C; boiling point 4,827°C; specific gravity of amorphous carbon 1.8 to 2.1, of diamond 3.15 to 3.53, of graphite 1.9 to 2.3; valence 2, 3, 4.
2. 1. A sheet of carbon paper.
   2. A carbon copy.
3. Electricity.
   1. Either of two rods through which current flows to form an arc, as in lighting or welding.
   2. A carbonaceous electrode in an electric cell.

[French carbone, from Latin carbō, carbōn-, a coal, charcoal.]

Concept

The phrase "carbon-based life forms," often used in science-fiction books and movies by aliens to describe the creatures of Earth, is something of a cliché. It is also a redundancy when applied to creatures on Earth, the only planet known to support life: all living things contain carbon. Carbon is also in plenty of things that were once living, which makes it useful for dating the remains of past settlements on Earth. Of even greater usefulness is petroleum, a substance containing carbon-based forms that died long ago, became fossilized, and ultimately changed chemically into fuels. Then again, not all materials containing carbon were once living creatures; yet because carbon is a common denominator to all living things on Earth, the branch of study known as organic chemistry is devoted to the study of compounds containing carbon. Among the most important organic compounds are the many carboxylic acids that are vital to life, but carbon is also present in numerous important inorganic compounds—most notably carbon dioxide and carbon monoxide.

How It Works

The Basics of Carbon

Carbon's name comes from the Latin word carbo, or charcoal—which, indeed, is almost pure carbon. Its chemical symbol is C, and it has an atomic number of 6, meaning that there are six protons in its nucleus. Its two stable isotopes are ¹²C, which constitutes 98.9% of all carbon found in nature, and ¹³C, which accounts for the other 1.1%.

The mass of the ¹²C atom is the basis for the atomic mass unit (amu), by which mass figures for all other elements are measured: the amu is defined as exactly 1/12 the mass of a single ¹²C atom. The difference in mass between ¹²C and ¹³C, which is heavier because of its extra neutron, account for the fact that the atomic mass of carbon is 12.01 amu: were it not for the small quantities of ¹³C present in a sample of carbon, the mass would be exactly 12.00 amu.

Where Carbon Is Found

Carbon makes up only a small portion of the known elemental mass in Earth's crust, oceans, and atmosphere—just 0.08%, or 1/1250 of the whole—yet it is the fourteenth most abundant element on the planet. In the human body, carbon is second only to oxygen in abundance, and accounts for 18% of the body's mass. Thus if a person weighs 100 lb (45.3 kg), she is carrying around 18 lb (8.2 kg) of carbon—the same material from which diamonds are made!

Present in the inorganic rocks of the ground and in the living creatures above it, carbon is everywhere. Combined with other elements, it forms carbonates, most notably calcium carbonate (CaCO₃), which appears in the form of limestone, marble, and chalk. In combination with hydrogen, it creates hydrocarbons, present in deposits of fossil fuels: natural gas, petroleum, and coal. In the environment, carbon—in the form of carbon dioxide (CO₂)—is taken in by plants, which undergo the process of photosynthesis and release oxygen to animals. Animals breathe in oxygen and release carbon dioxide to the atmosphere.

Carbon and Bonding

Located in Group 4 of the periodic table of elements (Group 14 in the IUPAC system), carbon has a valence electron configuration of 2s²2p²; likewise, all the members of Group 4—sometimes known as the "carbon family"—have configurations of ns²np², where n is the number of the period or row that the element occupies on the table.

There are two elements noted for their ability to form long strings of atoms and seemingly endless varieties of molecules: one is carbon, and the other is silicon, directly below it on the periodic table. Silicon, found in virtually all types of rocks except the calcium carbonates (mentioned above), is to the inorganic world what carbon is to the organic. Yet silicon atoms are about one and a half times as large as those of carbon; thus not even silicon can compete with carbon's ability to form a seemingly limitless array of molecules in various shapes and sizes, and having various chemical properties.

Basics of Chemical Bonding

Carbon is further distinguished by its high value of electronegativity, the relative ability of an atom to attract valence electrons. Electronegativity increases with an increase in group number, and decreases with an increase in period number. In other words, the elements with the highest electronegativity values lie in the upper right-hand corner of the periodic table.

Actually, the previous statement requires one significant qualification: the extreme right-hand side of the periodic table is occupied by elements with negligible electronegativity values. These are the noble gases, which have eight valence electrons each. Eight, as it turns out, is the "magic number" for chemical bonding: most elements follow what is known as the octet rule, meaning that when one element bonds to another, the two atoms have eight valence electrons.

If the two atoms have an electric charge and thus are ions, they form strong ionic bonds. Ionic bonding occurs when a metal bonds with a nonmetal. The other principal type of bond is a covalent bond, in which two uncharged atoms share eight valence electrons. If the electronegativity values of the two elements involved are equal, they share the electrons equally; but if one element has a higher electronegativity value, the electrons will be more drawn to that element.

Electronegativity of Carbon

To return to electronegativity and the periodic table, let us ignore the noble gases, which are the chemical equivalent of snobs. (Hence the term "noble," meaning that they are set apart.) To the left of the noble gases are the halogens, a wildly gregarious bunch—none more so than the element that occupies the top of the column, fluorine. With an electronegativity value of 4.0, fluorine is the most reactive of all elements, and the only one capable of bonding even to a few of the noble gases.

So why is fluorine—capable of forming multitudinous bonds—not as chemically significant as carbon? There are a number of answers, but a simple one is this: because fluorine is too strong, and tends to "overwhelm" other elements, precluding the possibility of forming long chains, it is less chemically significant than carbon. Carbon, on the other hand, has an electronegativity value of 2.5, which places it well behind fluorine. Yet it is still at sixth place (in a tie with iodine and sulfur) on the periodic table, behind only fluorine; oxygen (3.5); nitrogen and chlorine (3.0); and bromine (2.8). In addition, with four valence electrons, carbon is ideally suited to find other elements (or other carbon atoms) for forming covalent bonds according to the octet rule.

Multiple Bonds

Normally, an element does not necessarily have the ability to bond with as many other elements as it has valence electrons, but carbon—with its four valence electrons—happens to be tetravalent, or capable of bonding to four other atoms at once. Additionally, carbon is capable of forming not only a single bond, but also a double bond, or even a triple bond, with other elements.

Suppose a carbon atom bonds to two oxygen atoms to form carbon dioxide. Let us imagine these three atoms side by side, with the oxygen in the middle. (This, in fact, is how these bonds are depicted in the Couper and Lewis systems of chemical symbolism, discussed in the Chemical Bonding essay.) We know that the carbon has four valence electrons, that the oxygens have six, and that the goal is for each atom to have eight valence electrons—some of which it will share covalently.

Two of the valence electrons from the carbon bond with two valence electrons each from the oxygen atoms on either side. This means that the carbon is doubly bonded to each of the oxygen atoms. Therefore, the two oxygens each have four other unbonded valence electrons, which might bond to another atom. It is theoretically possible, also, for the carbon to form a triple bond with one of the oxygens by sharing three of its valence electrons. It would then have one electron free to share with the other oxygen.

Real-Life Applications

Organic Chemistry

We have stated that carbon forms tetravalent bonds, and makes multiple bonds with a single atom. In addition, we have mentioned the fact that carbon forms long chains of atoms and varieties of shapes. But how does it do these things, and why? These are good questions, but not ones we will attempt to answer here. In fact, an entire branch of chemistry is devoted to answering these theoretical questions, as well as to determining solutions to a host of other, more practical problems.

Organic chemistry is the study of carbon, its compounds, and their properties. (There are carbon-containing compounds that are not considered organic, however. Among these are oxides such as carbon dioxide and monoxide; as well as carbonates, most notably calcium carbonate.) At one time, chemists thought that "organic" was synonymous with "living," and even as recently as the early nineteenth century, they believed that organic substances contained a supernatural "life force." Then, in 1828, German chemist Friedrich Wöhler (1800-1882) cracked the code that distinguished the living from the nonliving, and the organic from the inorganic.

Wöhler took a sample of ammonium cyanate (NH₄OCN), and by heating it, converted it into urea (H₂N-CO-NH₂), a waste product in the urine of mammals. In other words, he had turned an inorganic material into a organic one, and he did so, as he observed, "without benefit of a kidney, a bladder, or a dog." It was almost as though he had created life. In fact, what Wöhler had glimpsed—and what other scientists who followed came to understand, was this: what separates the organic from the inorganic is the manner in which the carbon chains are arranged.

Ammonium cyanate and urea have exactly the same numbers and proportions of atoms, yet they are different compounds. They are thus isomers: substances which have the same formula, but are different chemically. In urea, the carbon forms an organic chain, and in ammonium cyanate, it does not. Thus, to reduce the specifics of organic chemistry even further, it can be said that this area of the field constitutes the study of carbon chains, and ways to rearrange them in order to create new substances.

Rubber, vitamins, cloth, and paper are all organically based compounds we encounter in our daily lives. In each case, the material comes from something that once was living, but what truly makes these substance organic in nature is the common denominator of carbon, as well as the specific arrangements of the atoms. We have organic chemistry to thank for any number of things: aspirins and all manner of other drugs; preservatives that keep food from spoiling; perfumes and toiletries; dyes and flavorings, and so on.

Allotropes of Carbon

Graphite

Carbon has several allotropes—different versions of the same element, distinguished by molecular structure. The first of these is graphite, a soft material with an unusual crystalline structure. Graphite is essentially a series of one-atom-thick sheets of carbon, bonded together in a hexagonal pattern, but with only very weak attractions between adjacent sheets. A piece of graphite is thus like a big, thick stack of carbon paper: on the one hand, the stack is heavy, but the sheets are likely to slide against one another.

Actually, people born after about 1980 may have little experience with carbon paper, which was gradually phased out as photocopiers became cheaper and more readily available. Today, carbon paper is most often encountered when signing a credit-card receipt: the signature goes through the graphite-based backing of the receipt, onto a customer copy.

In such a situation, one might notice that the copied image of the signature looks as though it were signed in pencil. This is not surprising, considering that pencil "lead" is, in fact, a mixture of graphite, clay, and wax. In ancient times, people did indeed use lead—the heaviest member of Group 4, the "carbon family"—for writing, because it left gray marks on a surface. Lead, of course, is poisonous, and is not used today in pencils or in most applications that would involve prolonged exposure of humans to the element. Nonetheless, people still use the word "lead" in reference to pencils, much as they still refer to a galvanized steel roof with a zinc coating as a "tin roof."

In graphite the atoms of each "sheet" are tightly bonded in a hexagonal, or six-sided, pattern, but the attractions between the sheets are not very strong. This makes it highly useful as a lubricant for locks, where oil would tend to be messy. A good conductor of electricity, graphite is also utilized for making high-temperature electrolysis cells. In addition, the fact that graphite resists temperatures of up to about 6,332°F (3,500°C) makes it useful in electric motors and generators.

Diamond

The second allotrope of carbon is also crystalline in structure. This is diamond, most familiar in the form of jewelry, but in fact widely applied for a number of other purposes. According to the Moh scale, which measures the hardness of minerals, diamond is a 10—in other words, the hardest type of material. It is used for making drills that bore through solid rock; likewise, small diamonds are used in dentists' drills for boring through the ultra-hard enamel on teeth.

Neither diamonds nor graphite are, in the strictest sense of the term, formed of molecules. Their arrangement is definite, as with a molecule, but their size is not: they simply form repeating patterns that seem to stretch on forever. Whereas graphite is in the form of sheets, a diamond is basically a huge "molecule" composed of carbon atoms strung together by covalent bonds. The size of this "molecule" corresponds to the size of the diamond: a diamond of 1 carat, for instance, contains about 10²² (10,000,000,000,000,000,000,000 or 10 billion billion) carbon atoms.

The diamonds used in industry look quite different from the ones that appear in jewelry. Industrial diamonds are small, dark, and cloudy in appearance, and though they have the same chemical properties as gem-quality diamonds, they are cut with functionality (rather than beauty) in mind. A diamond is hard, but brittle: in other words, it can be broken, but it is very difficult to scratch or cut a diamond—except with another diamond.

The cutting of fine diamonds for jewelry is an art, exemplified in the alluring qualities of such famous gems as the jewels in the British Crown or the infamous Hope Diamond in Washington, D.C.'s Smithsonian Institution. Such diamonds—as well as the diamonds on an engagement ring—are cut to refract or bend light rays, and to disperse the colors of visible light.

Buckminsterfullerene

Until 1985, carbon was believed to exist in only two crystalline forms, graphite and diamond. In that year, however, chemists at Rice University in Houston, Texas, and at the University of Sussex in England, discovered a third variety of carbon—and later jointly received a Nobel Prize for their work. This "new" carbon molecule composed of 60 bonded atoms in the shape of what is called a "hollow truncated icosahedron." In plain language, this is rather like a soccer ball, with interlocking pentagons and hexagons. However, because the surface of each geometric shape is flat, the "ball" itself is not a perfect sphere. Rather, it describes the shape of a geodesic dome, a design created by American engineer and philosopher R. Buckminster Fuller (1895-1983).

There are other varieties of buckminsterfullerene molecules, known as fullerenes. However, the 60-atom shape, designated as ⁶⁰C, is the most common of all fullerenes, the result of condensing carbon slowly at high temperatures. Fullerenes potentially have a number of applications, particularly because they exhibit a whole range of electrical properties: some are insulators, while some are conductors, semiconductors, and even superconductors. Due to the high cost of producing fullerenes artificially, however, the ways in which they are applied remain rather limited.

Amorphous Carbon

There is a fourth way in which carbon appears, distinguished from the other three in that it is amorphous, as opposed to crystalline, in structure. An example of amorphous carbon is carbon black, obtained from smoky flames and used in ink, or for blacking rubber tires.

Though it retains some of the microscopic structures of the plant cells in the wood from which it is made, charcoal—wood or other plant material that has been heated without enough air present to make it burn—is mostly amorphous carbon. One form of charcoal is activated charcoal, in which steam is used to remove the sticky products of wood decomposition. What remains are porous grains of pure carbon with enormous microscopic surface areas. These are used in water purifiers and gas masks.

Coal and coke are particularly significant varieties of amorphous carbon. Formed by the decay of fossils, coal was one of the first "fossil fuels" (for example, petroleum) used to provide heat and power for industrial societies. Indeed, when the words "industrial revolution" are mentioned, many people picture tall black smokestacks belching smoke from coal fires. Fortunately—from an environmental standpoint—coal is not nearly so widely used today, and when it is (as for instance in electric power plants), the methods for burning it are much more efficient than those applied in the nineteenth century.

Actually, much of what those smokestacks of yesteryear burned was coke, a refined version of coal that contains almost pure carbon. Produced by heating soft coal in the absence of air, coke has a much greater heat value than coal, and is still widely used as a reducing agent in the production of steel and other alloys.

Carbon Dioxide

Carbon forms many millions of compounds, some families of which will be discussed below. Two others, formed by the bonding of carbon atoms with oxygen atoms, are of particular significance. In carbon dioxide, a single carbon joins with two oxygens to produce a gas essential to plant life. In carbon monoxide (CO), a single oxygen joins the carbon, creating a toxic—but nonetheless important—compound.

The first gas to be distinguished from ordinary air, carbon dioxide is an essential component in the natural balance between plant and animal life. Animals, including humans, produce carbon dioxide by breathing, and humans further produce it by burning wood and other fuels. Plants use carbon dioxide when they store energy in the form of food, and they release oxygen to be used by animals.

Discovery

Flemish chemist and physicist Johannes van Helmont (1579-1644) discovered in 1630 that air was not, as had been thought up to that time, a single element: it contained a second substance, produced in the burning of wood, which he called "gas sylvestre." Thus he is recognized as the first scientist to note the existence of carbon dioxide.

More than a century later, in 1756, Scottish chemist Joseph Black (1728-1799) showed that carbon dioxide—which he called "fixed air"—combines with other chemicals to form compounds. This and other determinations Black made concerning carbon dioxide led to enormous progress in the discovery of gases by various chemists of the late eighteenth century.

By that time, chemists had begun to arrive at a greater degree of understanding with regard to the relationship between plant life and carbon dioxide. Up until that time, it had been believed that plants purify the air by day, and poison it at night. Carbon dioxide and its role in the connection between animal and plant life provided a much more sophisticated explanation as to the ways plants "breathe."

Uses

Around the same time that Black made his observations on carbon dioxide, English chemist Joseph Priestley (1733-1804) became the first scientist to put the chemical to use. Dissolving it in water, he created carbonated water, which today is used in making soft drinks. Not only does the gas add bubbles to drinks, it also acts as a preservative.

Though the natural uses of carbon dioxide are by far the most important, the compound has numerous industrial and commercial applications. Used in fire extinguishers, carbon dioxide is ideal for controlling electrical and oil fires, which cannot be put out with water. Heavier than air, carbon dioxide blankets the flames and smothers them.

In the solid form of dry ice, carbon dioxide is used for chilling perishable food during transport. It is also one of the only compounds that experiences sublimation, or the instantaneous transformation of a solid to a gas without passing through an intermediate liquid state, at conditions of ordinary pressure and temperature. Dry ice has often been used in movies to generate "mists" or "smoke" in a particular scene.

Carbon Monoxide

During the late eighteenth century, Priestley discovered a carbon-oxygen compound different from carbon dioxide: carbon monoxide. Scientists had actually known of this toxic gas, released in the incomplete combustion of wood, from the Middle Ages onward, but Priestley was the first to identify it scientifically.

Industry uses carbon monoxide in a number of ways. By blowing air across very hot coke, the result is producer gas, which, along with water gas (made by passing hot steam over coal) is an important fuel. Producer gas constitutes a 6:1:18 mixture of carbon monoxide, carbon dioxide, and nitrogen, while water gas is 40% carbon monoxide, 50% hydrogen, and 10% carbon dioxide and other gases.

Not only are producer and water gas used for fuel, they are also applied as reducing agents. Thus, when carbon monoxide is passed over hot iron oxides, the oxides are reduced to metallic iron, while the carbon monoxide is oxidized to form carbon dioxide. Carbon monoxide is also used in reactions with metals such as nickel, iron, and cobalt to form some types of carbonyls.

Carbon monoxide—produced by burning petroleum in automobiles, as well as by the combustion of wood, coal, and other carbon-containing fuels—is extremely hazardous to human health. It bonds with iron in hemoglobin, the substance in red blood cells that transports oxygen throughout the body, and in effect fools the body into thinking that it is receiving oxygenated hemoglobin, or oxyhemoglobin. Upon reaching the cells, carbon monoxide has much less tendency than oxygen to break down, and therefore it continues to circulate throughout the body. Low concentrations can cause nausea, vomiting, and other effects, while prolonged exposure to high concentrations can result in death.

Carbon and the Environment

Carbon is released into the atmosphere by one of three means: cellular respiration; the burning of fossil fuels; and the eruption of volcanoes. When plants take in carbon dioxide from the atmosphere, they combine this with water and manufacture organic compounds using energy they have trapped from sunlight by means of photosynthesis—the conversion of light to chemical energy through biological means. As a by-product of photosynthesis, plants release oxygen into the atmosphere.

In the process of undergoing photosynthesis, plants produce carbohydrates, which are various compounds of carbon, hydrogen, and oxygen essential to life. The other two fundamental components of a diet are fats and proteins, both carbon-based as well. Animals eat the plants, or eat other animals that eat the plants, and thus incorporate the fats, proteins, and sugars (a form of carbohydrate) from the plants into their bodies. Cellular respiration is the process whereby these nutrients are broken down to create carbon dioxide.

Photosynthesis and cellular respiration are thus linked in what is known as the carbon cycle. Cellular respiration also releases carbon into the atmosphere through the action of decomposers—bacteria and fungi that feed on the remains of plants and animals. The decomposers extract the energy in the chemical bonds of the decomposing matter, thus releasing more carbon dioxide into the atmosphere.

When creatures die and are buried in such a way that they cannot be reached by decomposers—for instance, at the bottom of the ocean, or beneath layers of rock—the carbon in their bodies is eventually converted to fossil fuels, including petroleum, natural gas, and coal. The burning of fossil fuels releases carbon (both monoxide and dioxide) into the atmosphere.

Because the rate of such burning has increased dramatically since the late nineteenth century, this has raised fears that carbon dioxide in the atmosphere may create a greenhouse effect, leading to global warming. On the other hand, volcanoes release tons of carbon into the atmosphere regardless of whether humans burn fossil fuels or not.

Radiocarbon Dating

Radiocarbon dating is used to date the age of charcoal, wood, and other biological materials. When an organism is alive, it incorporates a certain ratio of carbon-12 in proportion to the amount of the radioisotope (that is, radioactive isotope) carbon-14 that it receives from the atmosphere. As soon as the organism dies, however, it stops incorporating new carbon, and the ratio between carbon-12 and carbon-14 will begin to change as the carbon-14 decays to form nitrogen-14.

Carbon-14 has a half-life of 5,730 years, meaning that it takes that long for half the isotopes in a sample to decay to nitrogen-14. Therefore a scientist can use the ratios of carbon-12, carbon-14, and nitrogen-14 to guess the age of an organic sample. The problem with radiocarbon dating, however, is that there is a good likelihood the sample can become contaminated by additional carbon from the soil. Furthermore, it cannot be said with certainty that the ratio of carbon-12 to carbon-14 in the atmosphere has been constant throughout time.

Where to Learn More

Blashfield, Jean F. Carbon. Austin, TX: Raintree Steck-Vaughn, 1999.

"Carbon." Xrefer (Web site). <http://www.xrefer.com/entry/639742> (May 30, 2001).

"Diamonds." American Museum of Natural History (Web site). <http://www.amnh.org/exhibitions/diamonds/structure.html> (May 30, 2001).

Knapp, Brian J. Carbon Chemistry. Illustrated by David Woodroffe. Danbury, CT: Grolier Educational, 1998.

Loudon, G. Marc. Organic Chemistry. Menlo Park, CA: Benjamin/Cummings, 1988.

"Organic Chemistry" (Web site). <http://edie.cprost.sfu.ca/~rhlogan/organic.html> (May 30, 2001).

"Organic Chemistry." Frostburg State University Chemistry Helper (Web site). <http://www.chemhelper.com/> (May 30, 2001).

Sparrow, Giles. Carbon. New York: Benchmark Books, 1999.

Stille, Darlene. The Respiratory System. New York: Children's Press, 1997.

A chemical element, C, with an atomic number of 6 and an atomic weight of 12.01115. Carbon is unique in chemistry because it forms a vast number of compounds, larger than the sum total of all other elements combined. By far the largest group of these compounds are those composed of carbon and hydrogen. It has been estimated that there are at least 1,000,000 known organic compounds, and this number is increasing rapidly each year. Although the classification is not rigorous, carbon forms another series of compounds, classified as inorganic, comprising a much smaller number than the organic compounds. See also Organic chemistry; Periodic table.

Elemental carbon exists in two well-defined crystalline allotropic forms, diamond and graphite. Other forms, which are poorly developed in crystallinity, are charcoal, coke, and carbon black. Chemically pure carbon is prepared by the thermal decomposition of sugar (sucrose) in the absence of air. The physical and chemical properties of carbon are dependent on the crystal structure of the element. The density varies from 2.25 g/cm³ (1.30 oz/in.³) for graphite to 3.51 g/cm³ (2.03 oz/in.³) for diamond. For graphite, the melting point is 3500°C (6332°F) and the extrapolated boiling point is 4830°C (8726°F). Elemental carbon is a fairly inert substance. It is insoluble in water, dilute acids and bases, and organic solvents. At elevated temperatures, it combines with oxygen to form carbon monoxide or carbon dioxide. With hot oxidizing agents, such as nitric acid and potassium nitrate, mellitic acid, C₆(CO₂H)₆, is obtained. Of the halogens, only fluorine reacts with elemental carbon. A number of metals combine with the element at elevated temperatures to form carbides.

Carbon forms three gaseous compounds with oxygen: carbon monoxide, CO; carbon dioxide, CO₂; and carbon suboxide, C₃O₂. The first two oxides are the more important from an industrial standpoint. Carbon forms compounds with the halogens which have the general formula CX₄, where X is fluorine, chlorine, bromine, or iodine. At room temperature, carbon tetrafluoride is a gas, carbon tetrachloride is a liquid, and the other two compounds are solids. Mixed carbon tetrahalides are also known. Perhaps the most important of them is dichlorodifluoromethane, CCl₂F₂, commonly called Freon. See also Carbon dioxide; Halogenated hydrocarbon.

Carbon and its compounds are found widely distributed in nature. It is estimated that carbon makes up 0.032% of the Earth's crust. Free carbon is found in large deposits as coal, an amorphous form of the element which contains additional complex carbon-hydrogen-nitrogen compounds. Pure crystalline carbon is found as graphite and as diamonds.

Extensive amounts of carbon are found in the form of its compounds. In the atmosphere, carbon is present in amounts of up to 0.03% by volume as carbon dioxide. Various minerals such as limestone, dolomite, marble, and chalk all contain carbon in the form of carbonate. All plant and animal life is composed of complex organic compounds containing carbon combined with hydrogen, oxygen, nitrogen, and other elements. The remains of past plant and animal life are found as deposits of petroleum, asphalt, and bitumen. Deposits of natural gas contain compounds that are composed of carbon and hydrogen.

The free element has many uses, ranging from ornamental applications of the diamond in jewelry to the black-colored pigment of carbon black in automobile tires and printing inks. Another form of carbon, graphite, is used for high-temperature crucibles, arc-light and dry-cell electrodes, lead pencils, and as a lubricant. Charcoal, an amorphous form of carbon, is used as an absorbent for gases and as a decolorizing agent. See also Charcoal; Graphite.

The compounds of carbon find many uses. Carbon dioxide is used for the carbonation of beverages, for fire extinguishers, and in the solid state as a refrigerant. Carbon monoxide finds use as a reducing agent for many metallurgical processes. Carbon tetrachloride and carbon disulfide are important solvents for industrial uses. Freon is used in refrigeration devices. Calcium carbide is used to prepare acetylene, which is used for the welding and cutting of metals as well as for the preparation of other organic compounds. Other metal carbides find important uses as refractories and metal cutters.

A nonmetallic tetravalent element that occurs in pure form in diamonds and graphite. It occurs as a component of all living tissue. Most of the study of organic chemistry focuses on the vast number of carbon compounds.

For more information on carbon, visit Britannica.com.

A chemical element; its symbol is C. The carbon nucleus has six protons and six or more neutrons; six electrons are in orbit around the carbon nucleus. (See hydrocarbons and organic molecules.)

A chemical element, atomic number 6,atomic weight 12.011, symbol C.

• asymmetric c. atom — one bonded to four different atoms. See also isomer.
• c. fiber — made by the pyrolization of polymer fibers at very high temperatures and used in various forms as soft tissue implants, particularly in tendon and ligament repair.
• c. fixation — see dark reaction.

An element with atomic number 6; symbol: C. Carbon is one of the four elements essential for life. (The others are hydrogen, oxygen, and nitrogen.)

For other uses, see Carbon (disambiguation).

6 Boron (B) ← Carbon → Nitrogen (N) (None) ↑ C ↓ Silicon (Si) Periodic table
General
Name, symbol, number	Carbon, C, 6
Element category	Nonmetal
Group, period, block	14, 2, p
Appearance	Black (Graphite) Clear (Diamond)
Standard atomic weight	12.0107 g·mol−1
Electron configuration	1s2 2s2 2p2 or [He] 2s2 2p2
Electrons per shell	2,4 (Image)
Physical properties
Phase	Solid
Density (near r.t.)	amorphous:[1] 1.8 - 2.1 g·cm−3
Density (near r.t.)	graphite: 2.267 g·cm−3
Density (near r.t.)	diamond: 3.515 g·cm−3
Melting point	(1 atm) 3800 (graphite) K (3527 °C, 6381 °F)
Boiling point	(1 atm) 4300 K (4027 °C, 7281 °F)
Heat of fusion	117 (graphite) kJ·mol−1
Specific heat capacity	(25 °C) 8.517(graphite), 6.155(diamond) J·mol−1·K−1
Oxidation states	4, 3 [2], 2, 1 [3], 0, -1, -2, -3, -4[4]
Electronegativity	2.55 (Pauling scale)
Ionization energies (more)	1st: 1086.5 kJ·mol−1
	2nd: 2352.6 kJ·mol−1
	3rd: 4620.5 kJ·mol−1
Atomic radius	70 pm
Atomic radius (calc.)	67 pm
Covalent radius	77 pm
Van der Waals radius	170 pm
Miscellaneous
Magnetic ordering	diamagnetic[5]
Thermal conductivity	(300 K) 119-165 (graphite) 900-2300 (diamond) W·m−1·K−1
Thermal expansion	(25 °C) 0.8 (diamond) [6] µm·m−1·K−1
Speed of sound (thin rod)	(20 °C) 18350 (diamond) m/s
Young's modulus	1050 (diamond) [6] GPa
Shear modulus	478 (diamond) [6] GPa
Bulk modulus	442 (diamond) [6] GPa
Poisson ratio	0.1 (diamond) [6]
Mohs hardness	1-2 (Graphite) 10 (Diamond)
CAS registry number	7440-44-0
Most stable isotopes
Main article: Isotopes of Carbon iso NA half-life DM DE (MeV) DP 15 12C 98.9% 12C is stable with 6 neutrons 13C 1.1% 13C is stable with 7 neutrons 14C trace 5730 y beta- 0.156 14N
References
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Carbon in the form of graphite

Carbon (pronounced /ˈkɑrbən/) is the chemical element with symbol C and atomic number 6. As a member of group 14 on the periodic table, it is nonmetallic and tetravalent—making four electrons available to form covalent chemical bonds. There are three naturally occurring isotopes, with ¹²C and ¹³C being stable, while ¹⁴C is radioactive, decaying with a half-life of about 5730 years.^[7] Carbon is one of the few elements known since antiquity.^[8][9] The name "carbon" comes from Latin language carbo, coal, and, in some Romance and Slavic languages, the word carbon can refer both to the element and to coal.

There are several allotropes of carbon of which the best known are graphite, diamond, and amorphous carbon.^[10] The physical properties of carbon vary widely with the allotropic form. For example, diamond is highly transparent, while graphite is opaque and black. Diamond is among the hardest materials known, while graphite is soft enough to form a streak on paper (hence its name, from the Greek word "to write"). Diamond has a very low electrical conductivity, while graphite is a very good conductor. Under normal conditions, diamond has the highest thermal conductivity of all known materials. All the allotropic forms are solids under normal conditions but graphite is the most thermodynamically stable.

All forms of carbon are highly stable, requiring high temperature to react even with oxygen. The most common oxidation state of carbon in inorganic compounds is +4, while +2 is found in carbon monoxide and other transition metal carbonyl complexes. The largest sources of inorganic carbon are limestones, dolomites and carbon dioxide, but significant quantities occur in organic deposits of coal, peat, oil and methane clathrates. Carbon forms more compounds than any other element, with almost ten million pure organic compounds described to date, which in turn are a tiny fraction of such compounds that are theoretically possible under standard conditions.^[11]

Carbon is the fourth most abundant element in the universe by mass after hydrogen, helium, and oxygen. It is present in all known lifeforms, and in the human body carbon is the second most abundant element by mass (about 18.5%) after oxygen.^[12] This abundance, together with the unique diversity of organic compounds and their unusual polymer-forming ability at the temperatures commonly encountered on Earth, make this element the chemical basis of all known life.

Contents 1 Characteristics 1.1 Allotropes 1.2 Occurrence 1.3 Isotopes 1.4 Formation in stars 1.5 Carbon cycle 2 Compounds 2.1 Organic compounds 2.2 Inorganic compounds 2.3 Organometallic compounds 3 History and etymology 4 Production 4.1 Graphite 4.2 Diamond 5 Applications 5.1 Diamonds 6 Precautions 7 See also 8 References 9 External links

Characteristics

The different forms or allotropes of carbon (see below) include the hardest naturally occurring substance, diamond, and also one of the softest known substances, graphite. Moreover, it has an affinity for bonding with other small atoms, including other carbon atoms, and is capable of forming multiple stable covalent bonds with such atoms. As a result, carbon is known to form almost ten million different compounds; the large majority of all chemical compounds.^[11] Carbon also has the highest melting and sublimation point of all elements. At atmospheric pressure it has no actual melting point as its triple point is at 10 MPa so it sublimates above 4000 K.^[13][14]

Carbon sublimes in a carbon arc which has a temperature of about 5800K. Thus, irrespective of its allotropic form, carbon remains solid at higher temperatures than the highest melting point metals such as tungsten or rhenium. Although thermodynamically prone to oxidation, carbon resists oxidation more effectively than elements such as iron and copper that are weaker reducing agents at room temperature.

Diamond and graphite are two allotropes of carbon: pure forms of the same element that differ in structure.

Carbon compounds form the basis of all known so far life on Earth, and the carbon-nitrogen cycle provides some of the energy produced by the Sun and other stars. 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 sulfuric acid, hydrochloric acid, chlorine or any alkalis. At elevated temperatures carbon reacts with oxygen to form carbon oxides, and will reduce such metal oxides as iron oxide to the metal. This exothermic reaction is used in the iron and steel industry to control the carbon content of steel:

Fe₃O₄ + 4C_(s) → 3Fe_(s) + 4CO_(g)

with sulfur to form carbon disulfide and with steam in the coal-gas reaction
C_(s) + H₂O_(g) → CO_(g) + H_2(g).
Carbon combines with some metals at high temperatures to form metallic carbides, such as the iron carbide cementite in steel, and tungsten carbide, widely used as an abrasive and for making hard tips for cutting tools.

As to 2009, graphene appears the strongest material ever tested.^[15] However, the process of separating it from graphite will require some technological development before it is economical enough to be used in industrial processes.^[16]

The system of carbon allotropes spans a range of extremes:

Synthetic nanocrystalline diamond is the hardest materials known.	Graphite is one of the softest materials known.
Diamond is the ultimate abrasive.	Graphite is a very good lubricant.
Diamond is an excellent electrical insulator.	Graphite is a conductor of electricity.
Diamond is the best known naturally occurring thermal conductor	Some forms of graphite are used for thermal insulation (i.e. firebreaks and heat shields)
Diamond is highly transparent.	Graphite is opaque.
Diamond crystallizes in the cubic system.	Graphite crystallizes in the hexagonal system.
Amorphous carbon is completely isotropic.	Carbon nanotubes are among the most anisotropic materials ever produced.

Allotropes

Main article: Allotropes of carbon

Atomic carbon is a very short-lived species and therefore, carbon is stabilized in various multi-atomic structures with different molecular configurations called allotropes. The three relatively well-known allotropes of carbon are amorphous carbon, graphite, and diamond. Once considered exotic, fullerenes are nowadays commonly synthesized and used in research; they include buckyballs,^[17][18] carbon nanotubes,^[19] carbon nanobuds^[20] and nanofibers.^[21][22] Several other exotic allotropes have also been discovered, such as lonsdaleite,^[23] glassy carbon,^[24] carbon nanofoam^[25] and linear acetylenic carbon.^[26]

• The amorphous form, is an assortment of carbon atoms in a non-crystalline, irregular, glassy state, which is essentially graphite but not held in a crystalline macrostructure. It is present as a powder, and is the main constituent of substances such as charcoal, lampblack (soot) and activated carbon.

• 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 hexagonal rings, just like those in aromatic hydrocarbons. The resulting network is 2-dimensional, and the resulting flat sheets are stacked and loosely bonded through weak Van der Waals forces. This gives graphite its softness and its cleaving properties (the sheets slip easily past one another). Because of the delocalization of one of the outer electrons of each atom to form a π-cloud, graphite conducts electricity, but only in the plane of each covalently bonded sheet. This results in a lower bulk electrical conductivity for carbon than for most metals. The delocalization also accounts for the energetic stability of graphite over diamond at room temperature.

Some allotropes of carbon: a) diamond; b) graphite; c) lonsdaleite; d–f) fullerenes (C₆₀, C₅₄₀, C₇₀); g) amorphous carbon; h) carbon nanotube.

• At very high pressures carbon forms the more compact allotrope diamond, having nearly twice the density of graphite. Here, each atom is bonded tetrahedrally to four others, thus making a 3-dimensional network of puckered six-membered rings of atoms. Diamond has the same cubic structure as silicon and germanium and, thanks to the strength of the carbon-carbon bonds is the hardest naturally occurring substance in terms of resistance to scratching. Contrary to the popular belief that "diamonds are forever", they are in fact thermodynamically unstable under normal conditions and transform into graphite.^[10] But due to a high activation energy barrier, the transition into graphite is so extremely slow at room temperature as to be unnoticeable.

• Under some conditions, carbon crystallizes as lonsdaleite. This form has a hexagonal crystal lattice where all atoms are covalently bonded. Therefore, all properties of lonsdaleite are close to those of diamond. ^[23]

• Fullerenes have a graphite-like structure, but instead of purely hexagonal packing, they also contain pentagons (or even heptagons) of carbon atoms, which bend the sheet into spheres, ellipses or cylinders. The properties of fullerenes (split into buckyballs, buckytubes and nanobuds) have not yet been fully analyzed and represents an intense area of research in nanomaterials. The names "fullerene" and "buckyball" are given after Richard Buckminster Fuller, popularizer of geodesic domes, which resemble the structure of fullerenes. The buckyballs are fairly large molecules formed completely of carbon bonded trigonally, forming spheroids (the best-known and simplest is the soccerball-shaped structure C₆₀ buckminsterfullerene).^[17] Carbon nanotubes are structurally similar to buckyballs, except that each atom is bonded trigonally in a curved sheet that forms a hollow cylinder.^[18][19] Nanobuds were first published in 2007 and are hybrid bucky tube/buckyball materials (buckyballs are covalently bonded to the outer wall of a nanotube) that combine the properties of both in a single structure.^[20]

• Of the other discovered allotropes, Carbon nanofoam is a ferromagnetic 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/m³.^[27] Similarly, glassy carbon contains a high proportion of closed porosity.^[24] But unlike normal graphite, the graphitic layers are not stacked like pages in a book, but have a more random arrangement. Linear acetylenic carbon^[26] has the chemical structure^[26] -(C:::C)_n- .Carbon in this modification is linear with sp orbital hybridization, and is a polymer with alternating single and triple bonds. This type of carbyne is of considerable interest to nanotechnology as its Young's modulus is forty times that of the hardest known material - diamond.^[28]

Occurrence

Graphite ore

Raw diamond crystal.

Carbon is the fourth most abundant chemical element in the universe by mass after hydrogen, helium, and oxygen. Carbon is abundant in the Sun, stars, comets, and in the atmospheres of most planets. Some meteorites contain microscopic diamonds that were formed when the solar system was still a protoplanetary disk. Microscopic diamonds may also be formed by the intense pressure and high temperature at the sites of meteorite impacts.^[29]

"Present day" (1990s) sea surface dissolved inorganic carbon concentration (from the GLODAP climatology)

In combination with oxygen in carbon dioxide, carbon is found in the Earth's atmosphere (in quantities of approximately 810 gigatonnes) and dissolved in all water bodies (approximately 36,000 gigatons). Around 1,900 gigatons are present in the biosphere. Hydrocarbons (such as coal, petroleum, and natural gas) contain carbon as well—coal "reserves" (not "resources") amount to around 900 gigatons, and oil reserves around 150 gigatons. With smaller amounts of calcium, magnesium, and iron, carbon is a major component in very large masses of carbonate rock (limestone, dolomite, marble etc.).

Coal is a significant commercial source of mineral carbon; anthracite containing 92–98% carbon^[30] and the largest source (4,000 Gt, or 80% of coal, gas and oil reserves) of carbon in a form suitable for use as fuel.^[31]

Graphite is found in large quantities in New York and Texas, the United States, Russia, Mexico, Greenland, and India.

Natural diamonds occur in the rock kimberlite, found in ancient volcanic "necks," or "pipes". Most diamond deposits are in Africa, notably in South Africa, Namibia, Botswana, the Republic of the Congo, and Sierra Leone. There are also deposits in Arkansas, Canada, the Russian Arctic, Brazil and in Northern and Western Australia.

Diamonds are now also being recovered from the ocean floor off the Cape of Good Hope. However, though diamonds are found naturally, about 30% of all industrial diamonds used in the U.S. are now made synthetically.

According to studies from the Massachusetts Institute of Technology, an estimate of the global carbon budget is:^{[citation needed]}

Biosphere, oceans, atmosphere
0.45 x 1018 kilograms (3.7 x 1018 moles)
Crust
Organic carbon	13.2 x 1018 kg
Carbonates	62.4 x 1018 kg
Mantle
1200 x 1018 kg

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 cosmic rays. Thermal neutrons are produced that collide with the nuclei of nitrogen-14, forming carbon-14 and a proton.

Isotopes

Main article: Isotopes of carbon

Isotopes of carbon are atomic nuclei that contain six protons plus a number of neutrons (varying from 2 to 16). Carbon has two stable, naturally occurring isotopes.^[7] The isotope carbon-12 (¹²C) forms 98.93% of the carbon on Earth, while carbon-13 (¹³C) forms the remaining 1.07%.^[7] The concentration of ¹²C is further increased in biological materials because biochemical reactions discriminate against ¹³C.^[32] In 1961 the International Union of Pure and Applied Chemistry (IUPAC) adopted the isotope carbon-12 as the basis for atomic weights.^[33] Identification of carbon in NMR experiments is done with the isotope ¹³C.

Carbon-14 (¹⁴C) is a naturally occurring radioisotope which occurs in trace amounts on Earth of up to 1 part per trillion (0.0000000001%), mostly confined to the atmosphere and superficial deposits, particularly of peat and other organic materials.^[34] This isotope decays by 0.158 MeV β^- emission. Because of its relatively short half-life of 5730 years, ¹⁴C is virtually absent in ancient rocks, but is created in the upper atmosphere (lower stratosphere and upper troposphere) by interaction of nitrogen with cosmic rays.^[35] The abundance of ¹⁴C in the atmosphere and in living organisms is almost constant, but decreases predictably in their bodies after death. This principle is used in radiocarbon dating, invented in 1949, which has been used extensively to determine the age of carbonaceous materials with ages up to about 40,000 years.^[36][37]

There are 15 known isotopes of carbon and the shortest-lived of these is ⁸C which decays through proton emission and alpha decay and has a half-life of 1.98739x10^-21 s.^[38] The exotic ¹⁹C exhibits a nuclear halo, which means its radius is appreciably larger than would be expected if the nucleus was a sphere of constant density.^[39]

Formation in stars

Main articles: Triple-alpha process and CNO cycle

Formation of the carbon atomic nucleus requires a nearly simultaneous triple collision of alpha particles (helium nuclei) within the core of a giant or supergiant star. This happens in conditions of temperature and helium concentration that the rapid expansion and cooling of the early universe prohibited, and therefore no significant carbon was created during the Big Bang. Instead, the interiors of stars in the horizontal branch transform three helium nuclei into carbon by means of this triple-alpha process. In order to be available for formation of life as we know it, this carbon must then later be scattered into space as dust, in supernova explosions, as part of the material which later forms second, third-generation star systems which have planets accreted from such dust. The Solar System is one such third-generation star system.

One of the fusion mechanisms powering stars is the carbon-nitrogen cycle.

Rotational transitions of various isotopic forms of carbon monoxide (e.g. ¹²CO, ¹³CO, and C¹⁸O) are detectable in the submillimeter regime, and are used in the study of newly forming stars in molecular clouds.

Carbon cycle

Main article: Carbon cycle

Diagram of the carbon cycle. The black numbers indicate how much carbon is stored in various reservoirs, in billions of tons ("GtC" stands for gigatons of carbon; figures are circa 2004). The purple numbers indicate how much carbon moves between reservoirs each year. The sediments, as defined in this diagram, do not include the ~70 million GtC of carbonate rock and kerogen.

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 somewhere and dispose of it somewhere else. The paths that carbon follows in the environment make up the carbon cycle. For example, plants draw carbon dioxide out of their environment and use it to build biomass, as in carbon respiration or the Calvin cycle, a process of carbon fixation. Some of this biomass is eaten by animals, whereas 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; dead plant or animal matter may become petroleum or coal, which can burn with the release of carbon, should bacteria not consume it.^[40]

Compounds

Organic compounds

Main article: Organic compound

Structural formula of methane, the simplest possible organic compound.

Correlation between the carbon cycle and formation of organic compounds. In plants, carbon dioxide formed by carbon fixation can join with water in photosynthesis (green) to form organic compounds, which can be utilized and further converted by both plants and animals.

Carbon has the ability to form very long chains of interconnecting C-C bonds. This property is called catenation. Carbon-carbon bonds are strong, and stable. This property allows carbon to form an almost infinite number of compounds; in fact, there are more known carbon-containing compounds than all the compounds of the other chemical elements combined except those of hydrogen (because almost all organic compounds contain hydrogen too).

Carbon is the basis for all plastic materials that are used in common household items.

The simplest form of an organic molecule is the hydrocarbon—a large family of organic molecules that are composed of hydrogen atoms bonded to a chain of carbon atoms. Chain length, side chains and functional groups all affect the properties of organic molecules. By IUPAC's definition, all the other organic compounds are functionalized compounds of hydrocarbons.^{[citation needed]}

Carbon occurs in all known organic life and is the basis of organic chemistry. When united with hydrogen, it forms various flammable compounds called hydrocarbons which are important to industry as refrigerants, lubricants, solvents, as chemical feedstock for the manufacture of plastics and petrochemicals and as fossil fuels.

When combined with oxygen and hydrogen, carbon can form many groups of important biological compounds including sugars, lignans, chitins, alcohols, fats, and aromatic esters, carotenoids and terpenes. With nitrogen it forms alkaloids, and with the addition of sulfur also it forms antibiotics, amino acids, and rubber products. With the addition of phosphorus to these other elements, it forms DNA and RNA, the chemical-code carriers of life, and adenosine triphosphate (ATP), the most important energy-transfer molecule in all living cells.

Inorganic compounds

Main article: Compounds of carbon

Commonly carbon-containing compounds which are associated with minerals or which do not contain hydrogen or fluorine, are treated separately from classical organic compounds; however the definition is not rigid (see reference articles above). Among these are the simple oxides of carbon. The most prominent oxide is carbon dioxide (CO₂). This was once the principal constituent of the paleoatmosphere, but is a minor component of the Earth's atmosphere today.^[41] Dissolved in water, it forms carbonic acid (H₂CO₃), but as most compounds with multiple single-bonded oxygens on a single carbon it is unstable.^[42] Through this intermediate, though, resonance-stabilized carbonate ions are produced. Some important minerals are carbonates, notably calcite. Carbon disulfide (CS₂) is similar.

The other common oxide is carbon monoxide (CO). It is formed by incomplete combustion, and is a colorless, odorless gas. The molecules each contain a triple bond and are fairly polar, resulting in a tendency to bind permanently to hemoglobin molecules, displacing oxygen, which has a lower binding affinity.^[43][44] Cyanide (CN^–), has a similar structure, but behaves much like a halide ion (pseudohalogen). For example it can form the nitride cyanogen molecule ((CN)₂), similar to diatomic halides. Other uncommon oxides are carbon suboxide (C₃O₂),^[45] the unstable dicarbon monoxide (C₂O),^[46][47], carbon trioxide (CO₃), ^[48][49] cyclopentanepentone (C₅O₅) ^[50], cyclohexanehexone (C₆O₆) ^[50], and mellitic anhydride (C₁₂O₉).

With reactive metals, such as tungsten, carbon forms either carbides (C^4–), or acetylides (C₂^2–) to form alloys with high melting points. These anions are also associated with methane and acetylene, both very weak acids. With an electronegativity of 2.5,^[51] carbon prefers to form covalent bonds. A few carbides are covalent lattices, like carborundum (SiC), which resembles diamond.

Organometallic compounds

Main article: organometallic chemistry

Organometallic compounds by definition contain at least one carbon-metal bond. A wide range of such compounds exist; major classes include simple alkyl-metal compounds (e.g. tetraethyl lead), η²-alkene compounds (e.g. Zeise's salt, and η³-allyl compounds (e.g. allylpalladium chloride dimer; metallocenes containing cyclopentadienyl ligands (e.g. ferrocene); and transition metal carbene complexes. Many metal carbonyls exist (e.g. tetracarbonylnickel); some workers consider the carbon monoxide ligand to be purely inorganic, and not organometallic.

While carbon is understood to exclusively form four bonds, an interesting compound containing an octahedral hexacoordinated carbon atom has been reported. The cation of the compound is [(Ph₃PAu)₆C]²⁺. This phenomenon has been attributed to the aurophilicity of the gold ligands.^[52]

History and etymology

This section requires expansion.

The English name carbon comes from the Latin carbo for coal and charcoal,^[53] and hence comes from the French charbon, meaning charcoal. In German, Dutch and Danish, the names for carbon are Kohlenstoff, koolstof and kulstof respectively, all literally meaning coal-substance.

Carl Wilhelm Scheele

Antoine Lavoisier in his youth

Carbon was discovered in prehistory and was known in the forms of soot and charcoal to the earliest human civilizations. Diamonds were known probably as early as 2500 BCE in China, while carbon in the form of charcoal was made around Roman times by the same chemistry as it is today, by heating wood in a pyramid covered with clay to exclude air.^[54][55]

In 1722, René A. F. de Réaumur demonstrated that iron was transformed into steel through the absorption of some substance, now known to be carbon.^[56] In 1772, Antoine Lavoisier showed that diamonds are a form of carbon, when he burned samples of carbon and diamond then showed that neither produced any water and that both released the same amount of carbon dioxide per gram. Carl Wilhelm Scheele showed that graphite, which had been thought of as a form of lead, was instead a type of carbon.^[57] In 1786, the French scientists Claude Louis Berthollet, Gaspard Monge and C. A. Vandermonde then showed that this substance was carbon.^[58] In their publication they proposed the name carbone (Latin carbonum) for this element. Antoine Lavoisier listed carbon as an element in his 1789 textbook.^[59]

A new allotrope of carbon, fullerene, that was discovered in 1985^[60] includes nanostructured forms such as buckyballs and nanotubes.^[17] Their discoverers (Curl, Kroto, and Smalley) received the Nobel Prize in Chemistry in 1996.^[61] The resulting renewed interest in new forms lead to the discovery of further exotic allotropes, including glassy carbon, and the realization that "amorphous carbon" is not strictly amorphous.^[24]

Production

This section requires expansion.

Graphite

Main article: Graphite

Commercially viable natural deposits of graphite occur in many parts of the world, but the most important sources economically are in China, India, Brazil, and North Korea.^[62] Graphite deposits are of metamorphic origin, found in association with quartz, mica and feldspars in schists, gneisses and metamorphosed sandstones and limestone as lenses or veins, sometimes of a meter or more in thickness. Deposits of graphite in Borrowdale, Cumberland, England were at first of sufficient size and purity that, until the 1800s, pencils 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.

According to the USGS, world production of natural graphite in 2006 was 1.03 million tons and in 2005 was 1.04 million tons (revised), of which the following major exporters produced: China produced 720,000 tons in both 2006 and 2005, Brazil 75,600 tons in 2006 and 75,515 tons in 2005 (revised), Canada 28,000 tons in both years, and Mexico (amorphous) 12,500 tons in 2006 and 12,357 tons in 2005 (revised). In addition, there are two specialist producers: Sri Lanka produced 3,200 tons in 2006 and 3,000 tons in 2005 of lump or vein graphite, and Madagascar produced 15,000 tons in both years, a large portion of it "crucible grade" or very large flake graphite. Some other producers produce very small amounts of "crucible grade".

According to the USGS, U.S. (synthetic) graphite electrode production in 2006 was 132,000 tons valued at $495 million and in 2005 was 146,000 tons valued at $391 million, and high-modulus graphite (carbon) fiber production in 2006 was 8,160 tons valued at $172 million and in 2005 was 7,020 tons valued at $134 million.

Diamond

Main article: Diamond

Diamond output in 2005

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 X-ray fluorescence, after which the final sorting steps are done by hand. Before the use of X-rays 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.^[63]

Historically diamonds were known to be found only in alluvial deposits in southern India.^[64] India led the world in diamond production from the time of their discovery in approximately the 9th century BCE^[65] 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.^[66]

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.^[67] Interestingly 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.^[67]

In the United States, diamonds have been found in Arkansas, Colorado, and Montana.^[68][69] In 2004, a startling discovery of a microscopic diamond in the United States^[70] led to the January 2008 bulk-sampling of kimberlite pipes in a remote part of Montana.^[71]

Today, most commercially viable diamond deposits are in Russia, Botswana, Australia and the Democratic Republic of Congo.^[72] In 2005, Russia produced almost one-fifth of the global diamond output, reports the British Geological Survey. Australia boasts the richest diamantiferous pipe with production reaching peak levels of 42 metric tons (41 LT; 46 ST) per year in the 1990s.^[68]

There are also commercial deposits being actively mined in the Northwest Territories of Canada, Siberia (mostly in Yakutia territory, for example Mir pipe and Udachnaya pipe), Brazil, and in Northern and Western Australia. Diamond prospectors continue to search the globe for diamond-bearing kimberlite and lamproite pipes.

Applications

Pencil lead for mechanical pencils are made of graphite.

Sticks of vine and compressed charcoal.

A cloth of woven carbon filaments

Silicon carbide single crystal

The C₆₀ fullerene in crystalline form

Tungsten carbide milling bits

Carbon is essential to all known living systems, and without it life as we know it could not exist (see alternative biochemistry). The major economic use of carbon other than food and wood is in the form of hydrocarbons, most notably the fossil fuel methane gas and crude oil (petroleum). Crude oil is used by the petrochemical industry to produce, amongst others, gasoline and kerosene, through a distillation process, in refineries. Cellulose is a natural, carbon-containing polymer produced by plants in the form of cotton, linen, and hemp. Cellulose is mainly used for maintaining structure in plants. Commercially valuable carbon polymers of animal origin include wool, cashmere and silk. Plastics 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 alloys with iron, of which the most common is carbon steel. Graphite is combined with clays to form the 'lead' used in pencils used for writing and drawing. It is also used as a lubricant and a pigment, as a molding material in glass manufacture, in electrodes for dry batteries and in electroplating and electroforming, in brushes for electric motors and as a neutron moderator in nuclear reactors.

Charcoal is used as a drawing material in artwork, for grilling, and in many other uses including iron smelting. Wood, coal and oil are used as fuel for production of energy and space heating. Gem quality diamond is used in jewelry, and Industrial diamonds are used in drilling, cutting and polishing tools for machining metals and stone. Plastics are made from fossil hydrocarbons, and carbon fiber, made by pyrolysis of synthetic polyester fibers is used to reinforce plastics to form advanced, lightweight composite materials. Carbon fiber is made by pyrolysis of extruded and stretched filaments of polyacrylonitrile (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 specific tensile strength than steel.^[73]

Carbon black is used as the black pigment in printing ink, artist's oil paint and water colours, carbon paper, automotive finishes, India ink and laser printer toner. Carbon black is also used as a filler in rubber products such as tyres and in plastic compounds. Activated charcoal is used as an absorbent and adsorbent in filter material in applications as diverse as gas masks, water purification and kitchen extractor hoods and in medicine to absorb toxins, poisons, or gases from the digestive system. Carbon is used in chemical reduction at high temperatures. Coke is used to reduce iron ore into iron. Case hardening of steel is achieved by heating finished steel components in carbon powder. Carbides of silicon, tungsten, boron and titanium, are among the hardest known materials, and are used as abrasives in cutting and grinding tools. Carbon compounds make up most of the materials used in clothing, such as natural and synthetic textiles and leather, and almost all of the interior surfaces in the built environment other than glass, stone and metal.

Diamonds

The diamond industry can be broadly separated into two basically distinct categories: one dealing with gem-grade diamonds and another for industrial-grade diamonds. While a large trade in both types of diamonds exists, the two markets act in dramatically different ways.

A large trade in gem-grade diamonds exists. Unlike precious metals such as gold or platinum, gem diamonds do not trade as a commodity: there is a substantial mark-up in the sale of diamonds, and there is not a very active market for resale of diamonds.

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 20,000 kg annually), unsuitable for use as gemstones and known as bort, are destined for industrial use.^[74] In addition to mined diamonds, synthetic diamonds found industrial applications almost immediately after their invention in the 1950s; another 3 billion carats (600 metric tons) of synthetic diamond is produced annually for industrial use.^[75] The dominant industrial use of diamond is in cutting, drilling, grinding, and polishing. Most uses of diamonds in these technologies do not require large diamonds; in fact, most diamonds that are gem-quality except for their small size, can find an industrial use. Diamonds are embedded in drill tips or saw blades, or ground into a powder for use in grinding and polishing applications.^[76] Specialized applications include use in laboratories as containment for high pressure experiments (see diamond anvil cell), high-performance bearings, and limited use in specialized windows.^[77][78] With the continuing advances being made in the production of synthetic diamonds, future applications are beginning to become feasible. Garnering much excitement is the possible use of diamond as a semiconductor suitable to build microchips from, or the use of diamond as a heat sink in electronics.^[79]

Precautions

Worker at carbon black plant in Sunray, Texas (photo by John Vachon, 1942)

Pure carbon has extremely low toxicity and can be handled and even ingested 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, for example. Consequently if it gets into body tissues it is likely to remain there indefinitely. Carbon black was probably one of the first pigments to be used for tattooing, and Ötzi the Iceman was found to have carbon tattoos that survived during his life and for 5200 years after his death.^[80] However, inhalation of coal dust or soot (carbon black) in large quantities can be dangerous, irritating lung tissues and causing the congestive lung disease coalworker's pneumoconiosis. Similarly, diamond dust used as an abrasive can do harm if ingested or inhaled. Microparticles of carbon are produced in diesel engine exhaust fumes, and may accumulate in the lungs.^[81] In these examples, the harmful effects may result from contamination of the carbon particles, with organic chemicals or heavy metals for example, rather than from the carbon itself.

Carbon may also burn vigorously and brightly in the presence of air at high temperatures, as in the Windscale fire, which was caused by sudden release of stored Wigner energy in the graphite core. Large accumulations of coal, which have remained inert for hundreds of millions of years in the absence of oxygen, may spontaneously combust when exposed to air, for example in coal mine waste tips.

The great variety of carbon compounds include such lethal poisons as tetrodotoxin, the lectin ricin from seeds of the castor oil plant Ricinus communis, cyanide (CN^-) and carbon monoxide; and such essentials to life as glucose and protein.

See also

• Carbon chauvinism
• Carbon footprint
• Low-carbon economy
• Organic chemistry
• Timeline of carbon nanotubes

References

1. ^ Chemical Rubber Company Handbook of Chemistry and Physics, 59th Edition, CRC Press, Inc, 1979
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External links

Wikimedia Commons has media related to: Carbon

Look up carbon in Wiktionary, the free dictionary.

• Carbon - Periodic Table of Videos
• Carbon on Britannica
• WebElements.com – Carbon
• Chemicool.com – Carbon
• It's Elemental – Carbon
• Extensive Carbon page at asu.edu
• Electrochemical uses of carbon
• Computational Chemistry Wiki
• Carbon - Super Stuff. Animation with sound and interactive 3D-models.
• BBC Radio 4 series "In Our Time", on Carbon, the basis of life, 15 June 2006
• Introduction to Carbon Properties geared for High School students.
• Comprehensive Data on Carbon

v • d • e   Inorganic carbon compounds CCl4 · CF · CF4 · CI4 · CO · CO2 · CO3 · CS · CS2 · CSe2 · C3O2

v • d • e Allotropes of carbon sp3 forms Diamond (cubic) ♦ Lonsdaleite (hexagonal diamond) sp2 forms Graphite ♦ Fullerenes (buckyballs (C20+)) ♦ Nanotubes ♦ Glassy carbon sp forms Linear acetylenic carbon mixed sp3/sp2 forms Amorphous carbon ♦ nanobuds ♦ Carbon nanofoam other forms C1 - C2 - C3 - C8 related Chaoite - Activated carbon - Carbon black - Charcoal - Carbon fiber - Fullerite - Aggregated diamond nanorods

v • d • e Periodic table
H																															He
Li	Be																									B	C	N	O	F	Ne
Na	Mg																									Al	Si	P	S	Cl	Ar
K	Ca															Sc	Ti	V	Cr	Mn	Fe	Co	Ni	Cu	Zn	Ga	Ge	As	Se	Br	Kr
Rb	Sr															Y	Zr	Nb	Mo	Tc	Ru	Rh	Pd	Ag	Cd	In	Sn	Sb	Te	I	Xe
Cs	Ba	La	Ce	Pr	Nd	Pm	Sm	Eu	Gd	Tb	Dy	Ho	Er	Tm	Yb	Lu	Hf	Ta	W	Re	Os	Ir	Pt	Au	Hg	Tl	Pb	Bi	Po	At	Rn
Fr	Ra	Ac	Th	Pa	U	Np	Pu	Am	Cm	Bk	Cf	Es	Fm	Md	No	Lr	Rf	Db	Sg	Bh	Hs	Mt	Ds	Rg	Uub	Uut	Uuq	Uup	Uuh	Uus	Uuo
Alkali metals Alkaline earth metals Lanthanoids Actinoids Transition metals Other metals Metalloids Other nonmetals Halogens Noble gases

This entry is from Wikipedia, the leading user-contributed encyclopedia. It may not have been reviewed by professional editors (see full disclaimer)

Dansk (Danish)
n. - kulstof, sod, kul

idioms:

• carbon copy    gennemslag
• carbon dating    datering med kulstof 14-metoden
• carbon dioxide    kuldioxid, carbondioxid, kulsyre
• carbon monoxide    kulilte, carbonmonoxid
• carbon paper    karbonpapir

Nederlands (Dutch)
koolstof

Français (French)
n. - carbone

idioms:

• carbon copy    (Imprim) copie carbone, (fig) réplique exacte
• carbon dating    datation au carbone 14
• carbon dioxide    dioxyde de carbone
• carbon monoxide    monoxyde de carbone
• carbon paper    papier carbone

Deutsch (German)
n. - Kohlenstoff, Durchschlag, Kohlepapier

idioms:

• carbon copy    Durchschlag
• carbon dating    Radiokarbonmethode (zur Datierung)
• carbon dioxide    Kohlendioxid, Kohlensäure
• carbon monoxide    Kohlenmonoxyd
• carbon paper    Kohlepapier

Ελληνική (Greek)
n. - (χημ.) άνθρακας, καρμπόν αντιγραφής, αντίγραφο με καρμπόν

idioms:

• carbon copy    αντίγραφο με καρμπόν, πιστό αντίγραφο
• carbon dating    χρονολόγηση με άνθρακα 14
• carbon dioxide    (χημ.) διοξείδιο του άνθρακα
• carbon monoxide    μονοξείδιο του άνθρακα
• carbon paper    καρμπόν, χαρτί αντιγραφής

Italiano (Italian)
copia, carbonio, carbone, cartacarbone

idioms:

• carbon copy    copia carbone
• carbon dating    datazione al carbonio
• carbon dioxide    anidride carbonica
• carbon monoxide    ossido di carbonio
• carbon paper    cartacarbone

Português (Portuguese)
n. - carbono (m) (Quím.)

idioms:

• carbon copy    cópia (f) em papel carbono
• carbon dating    datação (f) por carbono (Fís.)
• carbon dioxide    dióxido (m) de carbono
• carbon monoxide    monóxido (m) de carbono
• carbon paper    papel (m) carbono

Русский (Russian)
углерод, уголь, копировальная бумага

idioms:

• carbon copy    копия, точная копия
• carbon dating    установление даты по углероду
• carbon dioxide    углекислый газ
• carbon monoxide    угарный газ
• carbon paper    копировальная бумага

Español (Spanish)
n. - copia al carbón, papel carbónico

idioms:

• carbon copy    copia hecha con papel carbón
• carbon dating    datación por carbono 14
• carbon dioxide    dióxido de carbono, ácido carbónico
• carbon monoxide    monóxido de carbono, óxido de carbono
• carbon paper    papel carbónico, papel de calco

Svenska (Swedish)
n. - kol, kolspets, sot

中文（简体）(Chinese (Simplified))
碳, 复写的副本, 复写纸, 碳精棒

idioms:

• carbon copy    复写的副本, 极相似的人
• carbon dating    碳年代测定
• carbon dioxide    二氧化碳
• carbon monoxide    一氧化碳
• carbon paper    复写纸

中文（繁體）(Chinese (Traditional))
n. - 碳, 複寫的副本, 複寫紙, 碳精棒

idioms:

• carbon copy    複寫的副本, 極相似的人
• carbon dating    碳年代測定
• carbon dioxide    二氧化碳
• carbon monoxide    一氧化碳
• carbon paper    複寫紙

한국어 (Korean)
n. - 탄소의, 탄소봉, 복사지

日本語 (Japanese)
n. - 炭素, カーボン紙

idioms:

• carbon copy    写し, そっくりのもの
• carbon dating    放射性炭素年代測定法
• carbon dioxide    二酸化炭素
• carbon monoxide    一酸化炭素
• carbon paper    カーボン紙

العربيه (Arabic)
‏(الاسم) فحم, الكربون‏

עברית (Hebrew)
n. - ‮פחמן, פחם, העתק‬

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