As we have already mentioned, different allotropes (different forms of the same element) can have quite different properties. The carbon allotrope graphite is soft, grey/black, opaque, conducts electricity, and slippery – it makes a good lubricant.64 Diamond is hard, transparent, and does not conduct electricity. How can this be possible if both are pure carbon? The answer lies in how the carbon atoms are organized with respect to one another. Whereas the carbon atoms in diamond form a three-dimensional network, in graphite, the atoms are organized in two-dimensional sheets that stack one on top of the other. Within each two-dimensional sheet the carbon atoms are linked by covalent bonds in an extended array of six-membered rings. This means that the carbon sheets are very strongly bonded, but the interactions between sheets are much weaker. Although there are no covalent bonds between the sheets, the atoms of the sheets do interact through London dispersion forces, very much like the interactions that hold helium atoms together. Because the sheets interact over very much larger surface areas, however, these interactions are much stronger than those in helium. Yet another allotrope of carbon, graphene, consists of a single sheet of carbon atoms.65 These sheets can be rolled into tubes to form nanotubes that are the subject of intense research interest because of their inherently high tensile strength. Carbon atoms can also form spherical molecules, known as buckminsterfullerenes or buckyballs.66
The obvious question is, why don’t covalent bonds form between graphite sheets? Why are the patterns of covalent bonding so different: three-dimensional (tetrahedral) in diamond, with each carbon bonded to four others, and two-dimensional (planar) in graphite and graphene, with each carbon atom bonded to only three others? One way to describe the molecular structure is to use the hybrid orbital bonding model. As we discussed previously, to form the four bonds attached to each carbon atom in diamond, we needed to hybridize four atomic orbitals to form four bonding orbitals. We might think we only need three bonds in graphite/graphene because each carbon is only connected to three others. This is not exactly true. In graphite and grapheme we use a model in which only three atomic orbitals are hybridized—an s and two 2p orbitals in order to form three sp2 bonding orbitals. These orbitals attach each carbon atom to three other atoms. Just like in diamond the three bonds associated with each carbon atom in graphite/graphene move as far apart as possible to minimize electron pair repulsion; they lie at the points of a triangle (rather than a tetrahedron). This geometry is called trigonal planar and the C–C–C bond angle is 120° (see Figure).
All well and good, but this does not really explain why the carbons in graphite/graphene are attached to three other carbon atoms, whereas in diamond each carbon is attached to four others. Perhaps surprisingly there is no good answer for why carbon takes up different forms—except that it can. But in fact carbon does form four bonds in graphite (carbon almost always forms four bonds—a central principle of organic chemistry). The trick is that the four bonds are not always equivalent; in graphite the fourth bond is not formed by the sp2 bonding orbitals but rather involves an unhybridized 2p atomic orbital. These p orbitals stick out at right angles to the sheet and can overlap with p orbitals from adjacent carbons in the same sheet (see Figure). Remember that p orbitals have two regions of electron density. To explain the fact that graphite conducts electricity, we use an idea from molecular orbital (MO) theory, namely that bonding and antibonding MOs are formed from the adjacent p orbitals that extend over the sheet surface. The energy different between these orbitals is not large and electrons can move from one to the other, allowing the movement of electrons throughout the whole sheet of graphite, which gives it many of the properties that we associate with metals. Note that we use both the hybridization model, which explains the planar framework of C–C bonds in graphite, and molecular orbital theory, which explains graphite’s electrical conductivity. So before we delve further into the properties associated with graphite, let us take a look at bonding in metals.
Questions to Answer
1. Diamond and graphite appear to be quite different substances, yet both contain only carbon atoms. Why are the observable properties of diamond and graphite so different when they are made of the same substance?
2. The electron configuration of C is 1s2 2s2 2p2. Using the idea that each atom provides one electron to a bond, if carbon used atomic orbitals to bond, how many bonds would it form? Would they all be the same? What would be the bond angles if this were to happen? (Draw a picture of what this might look like.)
3. The electron configuration of C is 1s2 2s2 2p2 this means that carbon has 6 electrons. Why doesn't it form 6 bonds?
4. We have seen that carbon can form materials in which it bonds to 4 other atoms (sp3 hybridization) or three other atoms (sp2 hybridization). What would be the hybridization for a carbon that was only bonded to two atoms? How would the other (unhybridized) p orbitals influence the behavior of such material (assuming that it could form)?
Questions to Ponder
1. Could carbon form a three-dimensional structure by linking to two other carbon atoms?
2. Do you think diamonds are transparent to all forms of light, such as X-rays?
3. What does the color of graphite imply about the energies of the photons it absorbs?
64 In fact the sheets in graphite do not slip relative to each other very readily. On Earth graphite is a lubricant, but in space in the absence of small molecules like O2, N2 and H2O, graphite does not lubricate. It is thought that the sheets slip relative to each other as if they were rolling on ball bearings (the small molecules). As you might imagine, this discovery caused some consternation in high-flying airplanes where the engines began to fail because of lack of lubrication.
65 The Nobel Prize in Physics was awarded in 2010 for the discovery of graphene. http://nobelprize.org/nobel_prizes/p...aureates/2010/
66 In 1996 Smalley, Kroto, and Curl were awarded the Nobel Prize in Chemistry for the discovery of fullerenes. http://nobelprize.org/nobel_prizes/c...aureates/1996/