THE 1996 Nobel Prize in Chemistry was recently awarded to three scientists for their discovery of a new form of carbon. The scientists are Mr Richard E. Smalley and Mr Robert F. Curl Jnr of Rice University, Houston, in the US, and Mr Harold W. Kroto, University of Sussex, Brighton, Britain.
This new form of carbon displays a fantastic array of chemical properties and has initiated enormous research in physics and chemistry.
Carbon is a unique element in chemistry because it forms a vast number of compounds, larger than elements combined. The most common compounds are those carbon and hydrogen.
The first carbon compounds were identified in living matter and therefore the study of carbon compounds was called organic chemistry.
There are 92 naturally occurring elements, ranging from the lightest, hydrogen, to the heaviest, uranium. Carbon is the sixth lightest element and is denoted by the letter symbol C. Up until recently it was believed that elemental carbon can occur in nature only in three forms - diamond, graphite and amorphous carbon.
The carbon atoms are arranged differently in these forms of solid, and, technically speaking, we say that each has a different crystalline structure. When an element appears in more than one form, the various forms are called allotropes.
Allotropes can seem entirely different in appearance and properties. This is well illustrated by the allotropes of carbon. Diamond is the hardest substance known. The hardness of substances is commonly measured on the Mohs Scale, which assigns minerals - numbers from one, for talc, to 10, for diamond. A mineral of a particular number is capable of scratching all minerals with lower numbers.
The relative hardnesses of diamond and graphite are explained by their crystalline structures. In diamond, carbon atoms are arranged with absolute three dimensional symmetry. Each carbon atom is bonded to four others at equal distances, each of the four located at the apexes of a tetrahedron of which the carbon atom under consideration forms the centre. The tetrahedron belongs to the group of five Platonic solids, which also includes the cube, octahedron, dodecahedron, and icosahedron.
This compact arrangement means that diamond is substantially denser than graphite. It also means that diamond will not pull apart in any direction except under massive force. Other atoms are also capable of taking up the diamond configuration, but the carbon atoms hold together the tightest because they are the smallest.
Thus, diamond is the hardest known substance. Diamonds are naturally produced under tremendous pressure and temperature in molten rock, far below the earth's surface. Microscopic diamonds are also abundant in outer space.
In silicon carbide (carborundum), silicon atoms replace half the carbon atoms. But the silicon atoms are larger than carbon atoms and they do not bond as closely to their neighbours. Therefore, silicon carbide is not as hard as diamond, although it is very hard, at number nine on the Mohs Scale, and is a much used abrasive.
In graphite, the carbon atoms are arranged in layers. In each layer, the atoms are arranged in hexagons, and the hexagons fit together exactly like hexagonal tiles on the bathroom floor. Each carbon atom is strongly bonded to three others in equal fashion.
Each layer of carbon atoms is a comparatively large distance from the layers above and below, so that [the bonds between layers are weak, and one layer can be easily made to slide over the other. Therefore, graphite is not particularly hard and can actually be used as a lubricant.
Amorphous carbon is characterised by a very low degree of crystallinity. Amorphous carbon can be made by heating pure sugar at 900 C in the absence of air. Amorphous carbon is found in varying degrees of purity in charcoal, coal, coke, carbon black, and lamp black.
The arrangement of carbon atoms in graphite is more stable than the arrangement in diamond, under conditions on the surface of the earth. Consequently, there is a tendency for diamond to turn spontaneously into graphite. If you have a large collection of diamonds, please feel free to panic for a moment.
However, let me now reassure you that you are in no danger of waking up one morning to find your splendid diamond collection has turned into pencil leads overnight. The carbon atoms in your diamonds are held together so tightly that it will take millions of years before they change into graphite.
ON the other hand, it is possible to convert graphite into diamond. Diamond is 55 per cent denser than graphite and the conversion is accomplished by putting the graphite under pressure and thereby forcing the atoms into the tighter packed arrangement characteristic of diamond. Let me quickly add, as you contemplate the lead in your pencil with quickened interest, that high temperatures (5,000 C) and fantastically high pressures (200,000 atmospheres) are essential for the conversion of graphite to synthetic diamonds. These are usually too small and impure to be used as gems, but they are produced commercially as abrasives and cutting tools.
In 1985, scientists produced a stable form of carbon consisting of 60 carbon atoms, C60, in a roughly spherical shape. Because of its structure, the molecule was named Buckminster Eullerene ("Buckyball") in honour of R. Buckminster Fuller, the inventor of the geodesic dome which conforms to the same underlying structural formula.
The fullerenes can be considered, after graphite and diamond, to be a third well defined allotrope of carbon. They promise to have synthetic, pharmaceutical, and industrial applications.
In the fullerene molecule, an even number of carbon atoms is arranged on the surface of a closed hollow cage. Each atom is linked to three near neighbours and the bonds outline a polyhedral network, consisting of 12 pentaons and 20 linked exagons. The 60 carbon atoms lie on the surface of a sphere and are distributed with the symmetry of an icosahedron. The 12 pentagons are isolated and distributed symmetrically among the 20 linked hexagons - the symmetry is that of a modern soccer ball.
The properties of fullerenes that have been determined to date suggest there is likely to be a wide range of areas in which fullerenes or their derivatives will have uses, for example, as charge carriers in batteries. Thin films of C60 are transparent at low light intensity and opaque at high intensities. This means that a thin layer of C60 molecules may be used to protect the eye from high transient light pulses.
The whole fullerene area was born from pure fundamental science. Fullerenes were discovered as a consequence of astrophysically motivated chemical physics experiments. Pullerene was discovered as a direct result of physicochemical investigations that simulated processes occurring in stars and in space.
Consequently, the likelihood that fullerenes and their analogues are present in space is an exciting conjecture. But to me, the most striking feature of the fullerene story is how, once again, a whole new area of applied practical research and development has grown entirely out of a discovery made in basic research.