II.

Properties

Of all the elements, carbon is the only one suitable for building the variety of molecules necessary to sustain life. Carbon atoms can attach to each other to form chains, rings, or a crystal mesh. The chains may be thousands of carbon atoms long and either linear or branched, and the rings usually contain from three to six carbon atoms. Most organic compounds contain many carbon-hydrogen bonds. Some of the other elements that bond to carbon include oxygen, nitrogen, fluorine, chlorine, bromine, iodine, sulfur, and phosphorus.

A.

Isotopes

Every carbon atom contains six positively charged particles called protons in its nucleus and six or more neutral particles called neutrons. The carbon atom’s nucleus is surrounded by six negatively charged electrons. The number of neutrons in a carbon atom’s nucleus determines which isotope it is. Isotopes are atoms of the same element that have different numbers of neutrons in the nucleus. Three different isotopes of carbon exist in nature. The important isotopes of carbon are carbon-12, carbon-13, and carbon-14. Scientists identify them by their mass number, which is the sum of the number of protons and neutrons in an atom. Carbon-12 contains six protons and six neutrons, carbon-13 contains six protons and seven neutrons, and carbon-14 contains six protons and eight neutrons.

In nature, carbon-12 accounts for about 98.89 percent of all carbon. Carbon-13 has a natural abundance of 1.11 percent, and the amount of carbon-14 is negligible. The atomic mass of carbon is 12.011 atomic mass units (AMU), which is the average mass of the isotopes of carbon based on their abundance.

Scientists have found some important uses for the less abundant isotopes of carbon. The nucleus of carbon-13 is magnetic. This property enables scientists to detect nuclei of carbon-13 atoms using a technique called nuclear magnetic resonance (NMR). By detecting the location of carbon-13 atoms in carbon-based molecules, scientists can learn about the structure of these molecules. Carbon-14 is radioactive, that is, its nucleus is unstable and can spontaneously change into the nucleus of another element (see Radioactivity). In a given sample, half of the carbon-14 nuclei will disintegrate in about 5,730 years. Living organisms constantly replenish carbon in their systems, so that the amount of carbon-14 remains constant as long as an organism is alive. Knowing the original amount of carbon-14 in organisms, scientists can measure the amount of carbon-14 that has disintegrated in a fossilized organism and determine the amount of time that has passed since it died. This technique for determining the age of fossils is called carbon dating.

B.

Bonding

As with all atoms, the electrons in a carbon atom reside in layers, or shells, around the nucleus. Carbon atoms have two electrons in their inner shell, and this shell can only contain two electrons, so it is full. Carbon atoms have four outer, or valence, electrons in their next shell. This outer electron shell can hold eight electrons, and atoms in general are much more stable when their outer shell is full. To obtain a full outer shell, carbon atoms form four covalent bonds with other atoms. A covalent bond is a bond formed when two atoms share a pair of electrons. When two atoms share one pair of electrons, the covalent bond is called a sigma bond and it holds the electrons tightly between the two atoms. One pair of shared electrons is also called a single bond. When two atoms share two pairs of electrons (creating a double bond), the first shared pair forms a sigma bond, while the second pair forms a pi bond. The pi bond does not hold electrons as tightly as the sigma bond holds the first pair. When two atoms share three pairs of electrons (creating a triple bond), two of the bonds are pi bonds. Electrons in pi bonds are much more reactive than are electrons in sigma bonds. That is, pi electrons more easily split away from the bond and create bonds with other atoms, adding those atoms to the molecule.

Carbon atoms can bond together in chains, rings, and meshlike networks. If a carbon atom bonds with four identical atoms, those atoms will be equally distant from each other—at the tips of an imaginary tetrahedron, or a pyramid with a triangular base. Any two of the bonds form an angle of 109.5° when carbon is in a tetrahedral form.

C.

Allotropes

Carbon has multiple allotropes. Allotropes are different physical forms of the same element, such as a hard, highly structured crystal and a soft, less-structured substance. Allotropes differ in the way the atoms bond with each other and arrange themselves into a structure. Because of their different structures, allotropes have different physical and chemical properties. The three common allotropes of carbon are diamond, graphite, and amorphous carbon (examples of amorphous carbon include charcoal, soot, and the coal-derived fuel called coke). The density of diamond is about 3.5 grams per cubic centimeter (g/cm³), graphite ranges from 1.9 to 2.3 g/cm³, and amorphous carbon ranges from 1.8 to 2.1 g/cm³. Diamond is one of the hardest known materials, while graphite is one of the softest. These differences arise from the differences in bonding between the carbon atoms.

In diamond, each carbon atom bonds tetrahedrally to four other carbon atoms to form a three-dimensional lattice. The shared electron pairs are held tightly in sigma bonds between adjacent atoms. Pure diamond is an electrical insulator—it does not conduct electric current. It is colorless and, because of its hardness, is used in industrial cutting tools. Cut diamonds sparkle brilliantly, which makes them treasured gemstones in jewelry.

Graphite is black and slippery and conducts electricity. In graphite, the atoms form planar, or flat, layers. Each layer is made up of rings containing six carbon atoms. The rings are linked to each other in a structure that resembles the hexagonal mesh of chicken wire. Each atom has three sigma bonds (with 120° between any two of the bonds) and belongs to three neighboring rings. The fourth electron of each atom becomes part of an extensive pi bond system. Graphite conducts electricity, because the electrons in the pi bond system can move around throughout the graphite. Bonds between atoms within a layer of graphite are strong, but the forces between the layers are weak. Because the layers can slip past each other, graphite is soft and can be used as a lubricant. Rubbing off layers of carbon in graphite is easy; you do it every time you write with a “lead” pencil. The “lead” is not actually lead at all but graphite mixed with clay. Diamond makers can transform graphite into diamond by applying extremely high pressure (more than 100,000 times the atmospheric pressure at sea level) and temperature (about 3000°C or 5000°F). High temperatures break the strong bonds in graphite so that the atoms can rearrange themselves into a diamond lattice. About 90 percent of the diamonds used in tools in the United States are made this way.

Amorphous carbon is actually made up of tiny crystal-like bits of graphite with varying amounts of other elements, which are considered impurities. For example, the coal industry divides coal up into various grades depending on the amount of carbon in the coal and the amount of impurities. The highest grade, anthracite, contains about 90 percent carbon. Lower grades include bituminous coal, which is 76 percent to 90 percent carbon, subbituminous coal, with 60 percent to 80 percent, and lignite, with 55 percent to 73 percent.

In 1985 chemists created a new allotrope of carbon by heating graphite to extremely high temperatures. They named the allotrope buckminsterfullerene, after American architect R. Buckminster Fuller. Fuller designed geodesic domes, rigid structures with a three-dimensional geometry that resemble this form of carbon. Unlike diamond and graphite, which can have an unending crystal structure, the original fullerene forms molecules of 60 carbon atoms (with a molecular formula of C₆₀). The molecules are shaped like tiny soccer balls (called buckyballs), with an atom at each point where the lines on a soccer ball would normally meet. The 60 carbon atoms bond in 20 six-membered rings and 12 five-membered rings. Each carbon atom is at a corner where two six-membered rings and one five-membered ring come together. Scientists have since discovered other fullerenes, including very narrow, long tubes and the C₇₀ fullerene, an elongated structure shaped more like a football but rounded on the ends. After scientists discovered fullerenes in the lab, geologists discovered fullerenes in nature—in ancient rocks in New Zealand and in the meteorite-created Ries Crater in Germany.

Scientists, excited by the properties of these recently discovered materials, are exploring ways to use them. When cooled, some fullerene-based compounds that include other noncarbon atoms are superconductors, that is, they can conduct electricity with no resistance. Some pure carbon fullerene tubes are stronger than metals and conduct electricity. Someday we may use them as electrical wires or as fibers to reinforce plastic, making materials that are even stronger than those reinforced with current carbon fibers (see Composite Material). Other compounds based on C₆₀ appear to inhibit the activity of the virus that causes acquired immunodeficiency syndrome (AIDS).