Fig. 2.3. Schematic presentation of sp3, sp2, and sp1 bonding hybridizations [2.9], The unshaded lobes denote strong bonds and the shaded lobes denote weak bonds.
forms. In addition, many solid-state variants of graphite and diamond have been developed, and these are also briefly reviewed in this chapter.
Carbon has an atomic number of 6 and is in a \s22s22p2 electronic ground state configuration. In the graphite structure, strong in-plane bonds are formed between a carbon atom and its three nearest-neighbors from 2s, 2px, and 2py orbitals; this bonding arrangement is denoted by sp2 (see Fig. 2.3). The remaining electron with a pz orbital provides only weak in-terplanar bonding [2.10], but is responsible for the semimetallic electronic behavior in graphite [2.11]. In contrast, the carbon atoms in the diamond structure are tetrahedrally bonded to their four nearest-neighbors using linear combinations of 2s, 2px, 2py, and 2pz orbitals in an sp3 configuration (see Fig. 2.3). The difference in the structural arrangement of these allotropie forms of carbon gives rise to the wide differences in their physical properties.
The ideal crystal structure of graphite (see Fig. 2.2) consists of layers in which the carbon atoms are arranged in an open honeycomb network, such that the A and B atoms on consecutive layers are on top of one another, but the A' atoms in one plane are over the unoccupied centers of the adjacent layers, and similarly for the B' atoms [2.4], This gives rise to an ABAB planar stacking arrangement shown in Fig. 2.2, with an in-plane nearest-neighbor distance ac_c of 1.421 A, an in-plane lattice constant a0 of 2.462 A, a c-axis lattice constant c0 of 6.708 A, and an interplanar distance of c0/2 = 3.3539 A (see Table 2.1). This structure is consistent with the Dlh(P63/mmc) space group and has four atoms per unit cell (see Fig. 2.2).
Disorder tends to have little effect on the in-plane lattice constant, largely because the in-plane C-C bond is very strong and the nearest-neighbor
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