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 2s 22p 2; likewise, all the members of Group 4—sometimes known as the "carbon family"—have configurations of ns 2np 2, 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.
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.
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
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.
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.
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.
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 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.
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."
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.
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 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.
Different versions of the same element, distinguished by molecular structure.
Having no definite structure.
Naturally occurring compounds of carbon, hydrogen, and oxygen. These are primarily produced by green plants through the process of photosynthesis.
The process whereby nutrients from plants are broken down in an animal's body to create carbon dioxide.
A type of chemical bonding in which two atoms share valence electrons.
A term describing a type of solid in which the constituent parts have a simple and definite geometric arrangement that is repeated in all directions.
A form of bonding in which two atoms share two pairs of valence electrons. Carbon is also capable of single bonds and triple bonds.
The relative ability of an atom to attract valence electrons.
An atom or group of atoms that has lost or gained one or more electrons, and thus has a net electric charge.
A form of chemical bonding that results from attractions between ions with opposite electric charges.
Substances which have the same chemical formula, but which are different chemically due to differences in the arrangement of atoms.
Atoms that have an equal number of protons, and hence are of the same element, but differ in their number of neutrons. This results in a difference ofmass. Isotopes may be either stable or unstable. The latter type, known as radioisotopes, are radioactive.
A term describing the distribution of valence electrons that takes place in chemical bonding for most elements, which usually end up with eight valence electrons.
The study of carbon, its compounds, and their properties. (Many carbon-containing oxides and carbonates are not considered organic, however.)
The biological conversion of light energy (that is, electromagnetic energy) to chemical energy inp lants.
A term describing a phenomenon whereby certain isotopes known as radioisotopes are subject to a form of decay brought about by the emission of high-energy particles. "Decay" does not mean that the isotope "rots"; rather, it decays to form another isotope until eventually (though this may take a long time) it becomes stable.
A form of bonding in which two atoms share one pair of valence electrons. Carbon is also capable of double bonds and triple bonds.
Capable of bonding to four other elements.
A form of bonding in which two atoms share three pairs of valence electrons. Carbon is also capable of single bonds and double bonds.
Electrons that occupy the highest principal energy level in an atom. These are the electrons involved in chemical bonding.
"Carbon." Science of Everyday Things. . Encyclopedia.com. (July 25, 2017). http://www.encyclopedia.com/science/news-wires-white-papers-and-books/carbon
"Carbon." Science of Everyday Things. . Retrieved July 25, 2017 from Encyclopedia.com: http://www.encyclopedia.com/science/news-wires-white-papers-and-books/carbon
Note: This article, originally published in 1998, was updated in 2006 for the eBook edition.
Carbon is an extraordinary element. It occurs in more different forms than any other element in the periodic table. The periodic table is a chart that shows how chemical elements are related to each other. More than ten million compounds of carbon are known. No other element, except for hydrogen, occurs in even a fraction of that number of compounds.
As an element, carbon occurs in a striking variety of forms. Coal, soot, and diamonds are all nearly pure forms of carbon. Carbon also occurs in a form, discovered only recently, known as fullerenes or buckyballs. Buckyball carbon holds the promise for opening a whole new field of chemistry (see accompanying sidebar).
Carbon occurs extensively in all living organisms as proteins, fats, carbohydrates (sugars and starches), and nucleic acids.
Group 14 (IVA)
Carbon is such an important element that an entirely separate field of chemistry is devoted to this element and its compounds. Organic chemistry is the study of carbon compounds.
Discovery and naming
Humans have been aware of carbon since the earliest of times. When cave people made a fire, they saw smoke form. The black color of smoke is caused by unburned specks of carbon. The smoke may have collected on the ceiling of their caves as soot.
Later, when lamps were invented, people used oil as a fuel. When oil burns, carbon is released in the reaction, forming a sooty covering on the inside of the lamp. That form of carbon became known as lampblack. Lampblack was also often mixed with olive oil or balsam gum to make ink. And ancient Egyptians sometimes used lampblack as eyeliner.
One of the most common forms of carbon is charcoal. Charcoal is made by heating wood in the absence of air so it does not catch fire. Instead, it gives off water vapor, leaving pure carbon. This method for producing charcoal was known as early as the Roman civilization (509 b.c.-a.d. 476).
French physicist René Antoine Ferchault Reaumur (1683-1757) believed carbon might be an element. He studied the differences between wrought iron, cast iron, and steel. The main difference among these materials, he said, was the presence of a "black combustible material" that he knew was present in charcoal.
Carbon was officially classified as an element near the end of the eighteenth century. In 1787, four French chemists wrote a book outlining a method for naming chemical substances. The name they used, carbone, is based on the earlier Latin term for charcoal, charbon.
Coal, soot (nearly pure carbon), and diamonds are all nearly pure forms of carbon.
Carbon exists in a number of allotropic forms. Allotropes are forms of an element with different physical and chemical properties. Two allotropes of carbon have crystalline structures: diamond and graphite. In a crystalline material, atoms are arranged in a neat orderly pattern. Graphite is found in pencil "lead" and ball-bearing lubricants. Among the non-crystalline allotropes of carbon are coal, lampblack, charcoal, carbon black, and coke. Carbon black is similar to soot. Coke is nearly pure carbon formed when coal is heated in the absence of air. Carbon allotropes that lack crystalline structure are amorphous, or without crystalline shape.
The allotropes of carbon have very different chemical and physical properties. For example, diamond is the hardest natural substance known. It has a rating of 10 on the Mohs scale. The Mohs scale is a way of expressing the hardness of a material. It runs from 0 (for talc) to 10 (for diamond). The melting point of diamond is about 3,700°C (6,700°F) and its boiling point is about 4,200°C (7,600°F). Its density is 3.50 grams per cubic centimeter.
On the other hand, graphite is a very soft material. It is often used as the "lead" in lead pencils. It has a hardness of 2.0 to 2.5 on the Mohs scale. Graphite does not melt when heated, but sublimes at about 3,650°C (6.600°F). Sublimination is the process by which a solid changes directly to a gas when heated, without first changing to a liquid. Its density is about 1.5 to 1.8 grams per cubic centimeter. The numerical value for these properties varies depending on where the graphite originates.
The amorphous forms of carbon, like other non-crystalline materials, do not have clear-cut melting and boiling points. Their densities vary depending on where they originate.
Carbon does not dissolve in or react with water, acids, or most other materials. It does, however, react with oxygen. It burns in air to produce carbon dioxide (CO2) and carbon monoxide (CO). The combustion (burning) of coal gave rise to the Industrial Revolution (1700-1900).
Another highly important and very unusual property of carbon is its ability to form long chains. It is not unusual for two atoms of an element to combine with each other. Oxygen (O2), nitrogen (N2), hydrogen (H2), chlorine (Cl2), and bromine (Br2) are a few of the elements that can do this. Some elements can make even longer strings of atoms. Rings of six and eight sulfur atoms (S6 and S8), for example, are not unusual.
Carbon has the ability to make virtually endless strings of atoms. If one could look at a molecule of almost any plastic, for example, a long chain of carbon atoms attached to each other (and to other atoms as well) would be evident. Carbon chains can be even more complicated. Some chains have side chains hanging off them.
There is almost no limit to the size and shape of molecules that can be made with carbon atoms. (See accompanying diagrams.)
Buckyballs are a recently discovered form of pure carbon. These spheres are made up of exactly 60 linked carbon atoms.
Occurrence in nature
Carbon is the sixth most common element in the universe and the fourth most common element in the solar system. It is the second most common element in the human body after oxygen. About 18 percent of a person's body weight is due to carbon.
The black color of smoke is caused by unburned specks of carbon.
Carbon is the 17th most common element in the Earth's crust. Its abundance has been estimated to be between 180 and 270 parts per million. It rarely occurs as a diamond or graphite.
Both allotropes are formed in the earth over millions of years, when dead plant materials are squeezed together at very high temperatures. Diamonds are usually found hundreds or thousands of feet beneath the earth's surface. Africa has many diamond mines.
Carbon also occurs in a number of minerals. Among the most common of these minerals are the carbonates of calcium (CaCO3) and magnesium (MgCO3). Carbon also occurs in the form of carbon dioxide (CO2) in the atmosphere. Carbon dioxide makes up only a small part of the atmosphere (about 300 parts per million), but it is a crucial gas. Plants use carbon dioxide in the atmosphere in the process of photosynthesis. Photosynthesis is the process by which plants convert carbon dioxide and water to carbohydrates (starches and sugars). This process is the source of life on Earth.
Carbon also occurs in coal, oil, and natural gas. These materials are often known as fossil fuels. They get that name because of the way they were formed. They are the remains of plants and animals that lived millions of years ago. When they died, they fell into water or were trapped in mud. Over millions of years, they slowly decayed. The products of that decay process were coal, oil, and natural gas.
Some forms of coal are nearly pure carbon. Oil and natural gas are made primarily of hydrocarbons, which are compounds made of carbon and hydrogen.
Three isotopes of carbon occur in nature, carbon-12, carbon-13, and carbon-14. One of these isotopes, carbon-14, is radioactive. Isotopes are two or more forms of an element. Isotopes differ from each other according to their mass number. The number written to the right of the element's name is the mass number. The mass number represents the number of protons plus neutrons in the nucleus of an atom of the element. The number of protons determines the element, but the number of neutrons in the atom of any one element can vary. Each variation is an isotope.
Five artificial radioactive isotopes of carbon are known also. A radioactive isotope is one that breaks apart and gives off some form of radiation. Artificial radioactive isotopes can be made by firing very small particles (such as protons) at atoms. These particles stick in the atoms and make them radioactive.
Carbon-14 has some limited applications in industry. For example, it can be used to measure the thickness of objects, such as sheets of steel. The steel must always be the same thickness.
Carbon is the sixth most common element in the universe and the fourth most common element in the solar system.
In this process, a small sample of carbon-14 is placed above the conveyor belt carrying the steel sheet. A detection device is placed below the sheet. The detection device counts the amount of radiation passing through the sheet. If the sheet gets thicker, Less radiation gets through. If the sheet gets thinner, more radiation gets through. The detector records how much radiation passes through the sheet. If the amount becomes too high or too low, the conveyor belt is turned off. The machine making the sheet is adjusted to produce steel of the correct thickness.
The most important use of carbon-14 is in finding the age of old objects (see accompanying sidebar for more information).
Diamond, graphite, and other forms of carbon are taken directly from mines in the earth. Diamond and graphite can also be made in laboratories. Synthetic diamonds, for example, are made by placing pure carbon under very high pressures (about 800,000 pounds per square inch) and temperatures (about 2,700°C). The carbon is heated and squeezed in the same way organic material is heated and squeezed in the earth. Today, about a third of all diamonds used are synthetically produced.
W hen an organism is alive, it takes in carbon dioxide from the air around it. Most of that carbon dioxide is made of carbon-12, but a tiny portion consists of carbon-14. So the living organism always contains a very small amount of radioactive carbon, carbon-14. A detector next to the living organism would record radiation given off by the carbon-14 in the organism.
When the organism dies, it no longer takes in carbon dioxide. No new carbon-14 is added, and the old carbon-14 slowly decays into nitrogen. The amount of carbon-14 slowly decreases as time goes on. Over time, less and less radiation from carbon-14 is produced. The amount of carbon-14 radiation detected for an organism is a measure, therefore, of how long the organism has been dead. This method of determining the age of an organism is called carbon-14 dating.
The decay of carbon-14 allows archaeologists (people who study old civilizations) to find the age of once-living materials. Measuring the amount of radiation remaining indicates the approximate age.
There are many uses for carbon's two key allotropes, diamond and graphite. Diamonds are one of the most beautiful and expensive gemstones in the world. But they also have many industrial uses. Because they are so hard they are used to polish, grind, and cut glass, metals, and other materials. The bit on an oil-drilling machine may be made of diamonds. The tool used to make thin tungsten wires is also made of diamonds.
Synthetic diamonds are more commonly used in industry than in jewelry. Industrial diamonds do not have to be free of flaws, as do jewelry diamonds.
Graphite works well as pencil lead because it rubs off easily. It is also used as a lubricant. Graphite is added to the space between machine parts that rub against each other. The graphite allows the parts to slide over each other smoothly.
Graphite is also used as a refractory. Refractory material can withstand very high temperatures by reflecting heat away from itself. Refractory materials are used to line ovens used to maintain high temperatures.
Graphite is used in nuclear power plants. A nuclear power plant converts nuclear energy to electrical power. Graphite acts as a moderator by slowing down the neutrons used in the nuclear reaction.
Graphite is used to make black paint, in explosives and matches, and in certain kinds of cathode ray tubes, like the ones used in television sets.
Amorphous forms of carbon have many uses. These include the black color in inks, pigments (paints), rubber tires, stove polish, typewriter ribbons, and phonograph records.
One form of carbon is known as activated charcoal. The term activated means that the charcoal has been ground into a very fine powder. In this form, charcoal can remove impurities from liquids that pass through. For example, activated charcoal removes color and odor from oils and water solutions.
The decay of carbon-14 allows archaeologists (people who study old civilizations) to find the age of once-living materials.
Carbon dioxide (CO2) is used to make carbonated beverages (it's the fizz in soda pop and beer), in fire extinguishers, and as a propellant in aerosol products. A propellant is a gas that pushes liquids out of a spray can, such as those used for deodorant or hair spray. Carbon dioxide can also be frozen to a solid called dry ice. It is widely used as a way of keeping objects cold.
Carbon monoxide (CO) is another compound formed between carbon and oxygen. Carbon monoxide is a very toxic gas produced when something burns in a limited amount of air. Carbon monoxide is always formed when gasoline burns in the engine of an automobile and is a common part of air pollution. Old heating units can produce carbon monoxide. This colorless and odorless gas can cause headaches, illness, coma, or even death.
Carbon monoxide has a few important industrial uses. It is often used to obtain a pure metal from the ore of that metal:
It would take a very large book to describe all the uses of organic compounds, which are divided into a number of families. An organic family is a group of organic compounds with similar structures and properties. The largest organic family is the hydrocarbons, compounds that contain only carbon and hydrogen. Methane, or natural gas (CH4), ethane (C2H6), propane (C3H8), ethylene (C2H4), and benzene (C6H6) are all hydrocarbons.
Hydrocarbons are used as fuels. Gas stoves burn natural gas, which is mostly methane. Propane gas is a popular camping fuel, used in small stoves and lanterns. Another important use of hydrocarbons is in the production of more complicated organic compounds.
Buckyballs and nanotubes
In the 1980s, chemists discovered a new allotrope of carbon. The carbon atoms in this allotrope are arranged in a sphere-like form of 60 atoms. The form resembles a building invented by American architect Buckminster Fuller (1895-1983). The building is known as a geodesic dome.
Each of the points on the dome is occupied by one carbon atom. The discoverers named the new form of carbon buckminsterfullerene in honor of Fuller. That name is too long to use in everyday conversation so it is usually shortened to fullerene or buckyball.
The discovery of the fullerene molecule was very exciting to chemists. They had never seen a molecule like it. They have been studying ways of working with this molecule. One interesting technique has been to cut open just one small part of the molecule. Then they cut open a small part on a second molecule. Finally, they join the two buckyballs together. They get a double-buckyball.
Repeating this process over and over could result in triple-buckyballs, quadruple-buckyballs, and so on. As this process is repeated, the buckyball becomes a long narrow tube called a nanotube. Nanotubes are long, thin, and extremely tiny tubes somewhat like a drinking straw or a long piece of spaghetti.
Scientists are beginning to find ways of using nanotubes. One idea is to run a thin chain of metal atoms through the center of a nanotube. This allows it to act like a tiny electrical wire. Nanotubes may completely change many devices that will be made in the future.
Other organic families contain carbon, hydrogen, and oxygen. Methyl alcohol (wood alcohol) and ethyl alcohol (grain alcohol) are the most common members of the alcohol family.
Methyl alcohol is used to make other organic compounds and as a solvent (a substance that dissolves other substances). Ethyl alcohol is used for many of the same purposes. It is also the alcohol found in beer, wine, and hard liquor, such as whiskey and vodka.
All alcohols are poisonous but some alcohols are more poisonous than others. If not drunk in moderation, alcoholic beverages can damage the body and brain. And, if drunk in large quantities, they can cause death. Methyl alcohol is more toxic than ethyl alcohol. People who have drunk methyl alcohol by mistake have died.
The list of everyday products made from organic compounds is very long. It includes drugs, artificial fibers, dyes, artificial colors and flavors, food additives, cosmetics, plastics of all kinds, detergents, synthetic rubber, adhesives, antifreeze, pesticides and herbicides, synthetic fuels, and refrigerants.
Carbon is essential to life. Nearly every molecule in a living organism contains carbon. The study of carbon compounds that occur in living organisms is called biochemistry (bio- = life + -chemistry ).
Carbon can also have harmful effects on organisms. For example, coal miners sometimes develop a disease known as black lung. The name comes from the appearance of the miner's lungs. Instead of being pink and healthy, the miner's lungs are black. The black color is caused by coal dust inhaled by the miner. The longer a miner works digging coal, the more the coal dust is inhaled. That worker's lungs become more and more black.
Color is not the problem with black lung disease however. The coal dust in the lungs blocks the tiny holes through which oxygen gets into the lungs. As more coal dust accumulates, more holes are plugged up, making it harder for the miner to breathe. Many miners eventually die from black lung disease because they lose the ability to breathe.
Carbon is essential to life. Nearly every molecule in a living organism contains carbon.
Carbon monoxide poisoning is another serious health problem. Carbon monoxide is formed whenever coal, oil, or natural gas bums. For example, the burning of gasoline in cars and trucks produces carbon monoxide. Today, almost every person in the United States inhales some carbon monoxide every day.
Small amounts of carbon monoxide are not very dangerous. But larger amounts cause a variety of health problems. At low levels, carbon monoxide causes headaches, dizziness, nausea, and loss of balance. At higher levels, a person can lose consciousness. At even higher levels, carbon monoxide can cause death.
"Carbon (revised)." Chemical Elements: From Carbon to Krypton. . Encyclopedia.com. (July 25, 2017). http://www.encyclopedia.com/science/news-wires-white-papers-and-books/carbon-revised
"Carbon (revised)." Chemical Elements: From Carbon to Krypton. . Retrieved July 25, 2017 from Encyclopedia.com: http://www.encyclopedia.com/science/news-wires-white-papers-and-books/carbon-revised
Carbon is the non-metallic chemical element of atomic number 6 in Group 14 of the periodic table , symbol C, atomic weight 12.01, specific gravity as graphite 2.25, as diamond 3.51. Its stable isotopes are 12C (98.90%) and 13C (1.10%). The weight of the 12C atom is the international standard on which atomic weights are based. It is defined as weighing exactly 12.00000 atomic mass units.
Carbon has been known since prehistoric times. It gets its name from carbo, the Latin word for charcoal, which is almost pure carbon. In various forms, carbon is found not only on Earth, but in the atmospheres of other planets, in the Sun and stars, in comets and in some meteorites.
On Earth, carbon can be considered the most important of all the chemical elements , because it is the essential element in practically all of the chemical compounds in living things. Carbon compounds are what make the processes of life work. Beyond Earth, carbon-atom nuclei are an essential part of the nuclear fusion reactions that produce the energy of the Sun and of many other stars. Without carbon, the Sun would be cold and dark.
In the form of chemical compounds, carbon is distributed throughout the world as carbon dioxide gas, CO2, in the atmosphere and dissolved in all the rivers , lakes and oceans . In the form of carbonates, mostly calcium carbonate (CaCO3), it occurs as huge rocky masses of limestone , marble and chalk. In the form of hydrocarbons , it occurs as great deposits of natural gas , petroleum and coal . Coal is important not only as a fuel, but because it is the source of the carbon that is dissolved in molten iron to make steel.
All plants and animals on Earth contain a substantial proportion of carbon. After hydrogen and oxygen , carbon is the most abundant element in the human body, making up 10.7% of all the body's atoms. Carbon is found as the free (uncombined) element in three different allotropic forms-different geometrical arrangements of the atoms in the solid. The two crystalline forms (forms containing very definite atomic arrangements) are graphite and diamond. Graphite is one of the softest known materials, while diamond is one of the hardest.
There is also a shapeless, or amorphous , form of carbon in which the atoms have no particular geometric arrangement. Carbon black, a form of amorphous carbon obtained from smoky flames, is used to make rubber tires and inks black. Charcoal—wood or other plant material that has been heated in the absence of enough air to actually burn—is mostly amorphous carbon, but it retains some of the microscopic structure of the plant cells in the wood from which it was made. Activated charcoal is charcoal that has been steam-purified of all the gummy wood-decomposition products, leaving porous grains of pure carbon that have an enormous microscopic surface area . It is estimated that one cubic inch of activated charcoal contains 200,000 ft2 (18,580 m2) of microscopic surface. This huge surface has a stickiness, called adsorption, for molecules of gases and solids; activated charcoal is therefore used to remove impurities from water and air, such as in home water purifiers and in gas masks.
Graphite is a soft, shiny, dark gray or black, greasy-feeling mineral that is found in large masses throughout the world, including the United States, Brazil, England, western Europe , Siberia, and Sri Lanka. It is a good conductor of electricity and resists temperatures up to about 6,332°F (3,500°C), which makes it useful as brushes (conductors that slide along rotating parts) in electric motors and generators, and as electrodes in high-temperature electrolysiscells. Because of its slipperiness, it is used as a lubricant. For example, powdered graphite is used to lubricate locks, where oil might be too viscous. The "lead" in pencils is actually a mixture of graphite, clay, and wax. It is called "lead" because the metallic element lead (Pb) leaves gray marks on paper and was used for writing in ancient times. When graphite-based pencils came into use, they were called "lead pencils."
The reason for graphite's slipperiness is its unusual crystalline structure. It consists of a stack of one-atom-thick sheets of carbon atoms, bonded tightly together into a hexagonal pattern in each sheet, but with only very weak attractions—much weaker than actual chemical bonds—holding the sheets together. The sheets of carbon atoms can therefore slide easily over one another; graphite is slippery in the same way as layers of wet leaves on a sidewalk.
Diamond, the other crystalline form of pure carbon, is the world's hardest natural material, and is used in industry as an abrasive and in drill tips for drilling through rock in oil fields and human teeth in dentists' offices. On a hardness scale of one to ten, which mineralogists refer to as the Mohs scale of hardness, diamond is awarded a perfect ten. But that's not why diamonds are so expensive. They are the most expensive of all gems, and are kept that way by supply and demand. The supply is largely controlled by the De Beers Consolidated Mines, Inc. in South Africa , where most of the world's diamonds are mined, and the demand is kept high by the importance that is widely attributed to diamonds.
A diamond can be considered to be a single huge molecule consisting of nothing but carbon atoms that are strongly bonded to each other by covalent bonds, just as in other molecules. A one-carat diamond "molecule" contains 1022carbon atoms.
The beauty of gem-quality diamonds comes from their crystal clarity, their high refractivity (ability to bend light rays) and their high dispersion—their ability to spread light of different colors apart, which makes the diamond's rainbow "fire." Skillful chipping of the gems into facets (flat faces) at carefully calculated angles makes the most of their sparkle. Even though diamonds are hard, meaning that they can't be scratched by other materials, they are brittle—they can be cracked.
Carbon is unique among the elements because its atoms can form an endless variety of molecules with an endless variety of sizes, shapes and chemical properties. No other element can do that to anywhere near the degree that carbon can. In the evolution of life on Earth, nature has always been able to "find" just the right carbon compound out of the millions available, to serve just about any required function in the complicated chemistry of living things.
Carbon-containing compounds are called organic compounds, and the study of their properties and reactions is called organic chemistry. The name organic was originally given to those substances that are found in living organisms—plants and animals. Almost all of the chemical substances in living things are carbon compounds (water and minerals are the obvious exceptions), and the name organic was eventually applied to the chemistry of all carbon compounds, regardless of where they come from.
Having the atomic number six, every carbon atom has a total of six protons. Therefore, all carbon atoms with a neutral charge also have a total of six electrons. Two are in a completed inner orbit, while the other four are valence electronsouter electrons that are available for forming bonds with other atoms. An ion is an atom with either a negative or positive charge has either less or more electrons than the number of protons (respectively), and is referred to as either an anion (negatively charged) or a cation (positively charged).
It is impossible to summarize the properties of carbon's millions of compounds. Organic compounds can be classified into families that have similar properties, because they have certain groupings of atoms in common.
See also Carbon dating; Chemical bonds and physical properties; Chemical elements; Gemstones; Geochemistry; Historical geology
"Carbon." World of Earth Science. . Encyclopedia.com. (July 25, 2017). http://www.encyclopedia.com/science/encyclopedias-almanacs-transcripts-and-maps/carbon
"Carbon." World of Earth Science. . Retrieved July 25, 2017 from Encyclopedia.com: http://www.encyclopedia.com/science/encyclopedias-almanacs-transcripts-and-maps/carbon
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°C; graphite sublimes about 3,375°C; b.p. 4,827°C; 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).
There are more carbon compounds than there are compounds of all other elements combined. The study of carbon compounds, both natural and synthetic, is called organic chemistry. Plastics, foods, textiles, and many other common substances contain carbon. Hydrocarbon fuels (e.g., natural gas), marsh gas, and the gases resulting from the combustion of fuels (e.g., carbon monoxide and carbon dioxide) are compounds of carbon. With oxygen and a metallic element, carbon forms many important carbonates, such as calcium carbonate (limestone) and sodium carbonate (soda). Certain active metals react with it to make industrially important carbides, such as silicon carbide (an abrasive known as carborundum), calcium carbide, used for producing acetylene gas, and tungsten carbide, an extremely hard substance used for rock drills and metalworking tools.
Natural Occurrence and Uses
Carbon has been known to humans in its various forms since ancient times. Although carbon makes up only .032% of the earth's crust, it is very widely distributed and forms a vast number of compounds. Carbon exists in the stars; a series of thermonuclear reactions called the carbon cycle (see nucleosynthesis) is a source of energy for some stars. Carbon in the form of diamonds has been found in meteorites. Coke is used as a fuel in the production of iron. Carbon electrodes are widely used in electrical apparatus. The "lead" of the ordinary pencil is graphite mixed with clay. The successful linking in the 1940s of carbon with silicon has led to the development of a vast number of new substances known collectively as the silicones.
All living organisms contain carbon; the human body is about 18% carbon by weight. In green plants carbon dioxide and water are combined to form simple sugars (carbohydrates); light from the sun provides the energy for this process (photosynthesis). The energy from the sun is stored in the chemical bonds of the sugar molecule. Anabolism, the synthesis of complex compounds (such as fats, proteins, and nucleic acids) from simpler substances, involves the utilization of energy stored by photosynthesis. Catabolism is the release of stored energy by the oxidative destruction of organic compounds; water and carbon dioxide are two byproducts of catabolism. This continuing synthesis and degradation involving carbon dioxide is known as the biological carbon cycle.
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).
"carbon." The Columbia Encyclopedia, 6th ed.. . Encyclopedia.com. (July 25, 2017). http://www.encyclopedia.com/reference/encyclopedias-almanacs-transcripts-and-maps/carbon
"carbon." The Columbia Encyclopedia, 6th ed.. . Retrieved July 25, 2017 from Encyclopedia.com: http://www.encyclopedia.com/reference/encyclopedias-almanacs-transcripts-and-maps/carbon
melting point: 3,500°C
boiling point: 4,827°C
density: 2.62 g/cm3
most common ions: CO32−, HCO2−
Carbon is the sixth most abundant element in the universe and possibly the most widespread element on earth. Named from the Latin word carbo, meaning charcoal, it has been known since ancient times, although not recognized as an element until more modern times. It is found in all living things, but is also commonly found in minerals such as limestone and marble and as a small but important constituent of the atmosphere, as carbon dioxide.
In its elemental form, carbon can be found as diamond, the hardest naturally occurring substance; graphite, an excellent lubricant; or as a fullerene (or "buckyball"). Although five isotopes are known, only C-12 (98.9 percent natural abundance) and C-13 (1.1 percent) are stable. Nevertheless, the presence of trace amounts of C-14, a radioactive isotope (with a half-life of 5,730 years), permits "carbon dating" of historical objects.
Carbon is unique among the elements in that carbon atoms can form bonds with other carbon atoms. This property, known as "concatenation," is the reason that there are more than several million known organic (containing carbon) compounds.
Carbon is a major constituent of most of our fuels—natural gas, petroleum, coal, wood, and other biomass . When burned for their heat content, these fuels produce carbon dioxide, which escapes into the atmosphere. In the "carbon cycle," this carbon dioxide is trapped by plants and reincorporated into useful substances such as cellulose and starch via photosynthesis. Since the Industrial Revolution of the nineteenth century, which greatly increased the quantity of fuel combustion , the concentration of carbon dioxide in the atmosphere has increased. This atmospheric change has been implicated as the cause of global warming.
Carbon is a major constituent of most polymers, both naturally occurring ones such as cellulose, starch, RNA , DNA , silk, and wool, as well as synthetic ones, including nylon, Teflon, polyethylene, and polystyrene. Some of the strongest and most modern materials replacing metals are made of carbon fibers.
see also Allotropes; Global Warming; Organic Chemistry.
George H. Wahl Jr.
Jefferson Lab. "It's Elemental: The Element Carbon." Available from <http://education.jlab.org/itselemental/ele006.html>.
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"carbon." A Dictionary of Earth Sciences. . Retrieved July 25, 2017 from Encyclopedia.com: http://www.encyclopedia.com/science/dictionaries-thesauruses-pictures-and-press-releases/carbon
There are two stable isotopes of carbon (proton numbers 12 and 13) and four radioactive ones (10, 11, 14, 15). Carbon–14 is used in carbon dating.
"carbon." A Dictionary of Biology. . Encyclopedia.com. (July 25, 2017). http://www.encyclopedia.com/science/dictionaries-thesauruses-pictures-and-press-releases/carbon-2
"carbon." A Dictionary of Biology. . Retrieved July 25, 2017 from Encyclopedia.com: http://www.encyclopedia.com/science/dictionaries-thesauruses-pictures-and-press-releases/carbon-2
car·bon / ˈkärbən/ • n. the chemical element of atomic number 6, a nonmetal that has two main forms (diamond and graphite) and that also occurs in impure form in charcoal, soot, and coal. (Symbol: C) ∎ Chem. an atom of this element. ∎ a rod of carbon in an arc light. ∎ a piece of carbon paper or a carbon copy.
"carbon." The Oxford Pocket Dictionary of Current English. . Encyclopedia.com. (July 25, 2017). http://www.encyclopedia.com/humanities/dictionaries-thesauruses-pictures-and-press-releases/carbon-0
"carbon." The Oxford Pocket Dictionary of Current English. . Retrieved July 25, 2017 from Encyclopedia.com: http://www.encyclopedia.com/humanities/dictionaries-thesauruses-pictures-and-press-releases/carbon-0
So carbonate XVIII; see -ATE 1. Hence carbonaceous XVIII, carbonic XVIII.
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