Carbon

Carbon is unique in that no other element forms the huge variety of compounds that carbon can form with other elements. All forms of life on Earth are based on carbon compounds, as are most plastics, drugs, and dyes.

The preeminence of carbon in the ranks of more than 100 elements has long been recognized by chemists, who have divided the subject into two broad classes: organic and inorganic chemistry. Organic chemistry concerns the chemical reactions of compounds that contain carbon and usually hydrogen, while inorganic chemistry concerns compounds of the other elements. Although this might seem an unbalanced division, there are many more compounds of carbon than all the conceivable combinations of all the other elements—and by a huge factor.

Properties and occurrence

The atomic mass of carbon is approximately 12 times the atomic mass of hydrogen, the lightest of all elements. Each carbon atom has six electrons: two in an inner shell and four in an incomplete outer shell. A carbon atom tends to find four more electrons to complete its outer shell of eight electrons; in most cases, this bonding is achieved by sharing electrons in covalent bonds with other atoms. Because of this tendency, chemists say that carbon has a valency of four, meaning that it could bind with four hydrogen atoms, for example, each hydrogen atom donating an electron to its bond with the carbon atom.

Carbon is present as a mere 0.3 percent of Earth’s crust. Most of this carbon is bound up in carbonates such as limestone and chalk or in fossil fuels such as coal, gas, and oil. The fossil fuels formed through the action of bacteria, heat, and pressure on the remains of plants and animals over many millions of years. Coal derives from vegetation, while petroleum and natural gas derive from the decomposed remains of marine organisms.

Slightly impure carbon is formed when organic matter is heated in the absence of air. Heating wood in the absence of air drives off gases and volatile liquids to leave charcoal, for example, while treating coal in a similar way produces coke.

When ignited in air, charcoal glows at a high temperature forming little smoke, making charcoal ideal for cooking food in barbecues. Until the industrial revolution, charcoal was the main source of heat and carbon for producing iron from iron-oxide ores. Since then, the use of coke for the same purpose has permitted iron to be made on a much larger scale because coke is hard enough to support the weight of the charge in a blast furnace, whereas charcoal would not be.

Allotropes of carbon

A small proportion of the carbon in Earth’s crust is present as the element—an ability shared by few other elements, of which gold is the besnown example. Pure carbon occurs in two physical forms, called allotropes, and each corresponds to a different arrangement of carbon atoms.

Carbon is not alone in being allotropic: phosphorus exists as two allotropes—white and red phosphorus—that also have different atomic arrangements. What distinguishes the allotropes of carbon—diamond and graphite—is the vast difference between their properties. Diamond, for example, is a colorless transparent solid that is the hardest of all known substances, while graphite, on the other hand, is a soft darray material used in the lead for pencils and as an industrial lubricant. Diamond is a poor conductor of heat and electricity, while graphite is a good conductor of both, though not as good as a metal. The contrasting properties of graphite and diamond are due to their different structures.

The fact that diamond and graphite are two forms of the same element can be confirmed by burning them: the product is carbon dioxide in either case. The French chemist Antoine Lavoisier first performed this experiment in 1773.

The difference between the properties of the two allotropes of carbon stems from their different arrangements of carbon atoms. Each carbon atom has four electrons that are available for forming bonds. In diamond, each of these electrons forms a bond with a neighboring carbon atom, and the neighboring carbon atoms lie at the points of a tetrahedron with the first carbon atom at its center. Each of the carbon atoms in the tetrahedron connects with four other carbon atoms—including the first—and so on, so every carbon atom in a diamond is held firmly in place by four bonds. In a way, a diamond is one giant carbon molecule: there are no freonding electrons, and since electrons are responsible for conduction, diamond is neither a conductor of heat nor of electricity. The strength of bonding in a diamond makes it the hardest substance known—a great advantage in cutting applications.

The carbon atoms in graphite, however, are arranged in flat planes of hexagonal rings stacked one on another. Each carbon atom bonds to just three others within the plane. The remaining bonding electron wanders relatively loosely between the planes, accounting for graphite’s conductivity. The planes are held together by weaker attractions, a fact that is reflected in the larger separation between planes than between adjacent carbon atoms within a plane. The planes can slip over each other easily, making graphite soft and a useful lubricant. The free electrons in graphite are capable of absorbing and reflecting all frequencies of visible light, thus making graphite both dark and shiny. In contrast, bonding electrons of transparent colorless diamonds do not absorb light at visible frequencies.

Graphite can be turned into diamonds when temperatures of at least 4530°F (2500°C) are combined with a pressure of 700 tons per sq. in. (99 tonnes/cm²). Scientists at the U.S. General Electric Company first achieved this conversion in 1955. Synthetic diamonds made in this way are used to make cutting tools.