Carbon Alloy
Carbon Alloy
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Allotropes of Carbon
Carbon has several allotropes that are its different versions and are 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 one signs a credit-card receipt where in 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 diamond that also is crystalline in structure. People are most familiar with diamond in the form of jewelry, but in fact it is widely applied for a number of other purposes. According to the Mohs scale, which measures the hardness of minerals, diamond has the hardness of 10, in other words it is 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 is, 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, an invention for which they later jointly received a Nobel Prize. 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, due to which they have been dubbed as buckminsterfullerene.
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.
About the Author
Dr. Badruddin Khan teaches Chemistry in the University of Kashmir, Srinagar, india.
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