The structure of graphite

In going from its ground state to the graphite structure, a carbon atom's electronic configuration is believed to change as follows:

Three of the two 2$s$ and two 2$p$ electrons in carbon's ground state redistribute into three hybrid 2($sp^2$) orbitals which are a mathematical mixing of the $s$ orbitals with two of the three $p$ orbitals. The angular probabilities for these 2($sp^2$) orbitals can be represented by three coplanar lobes at 120$^{\circ}$ to each other [Fig.2.1]. The fourth electron of the original two 2$s$ and two 2$p$ electrons fills that $p$ orbital which does not participate in the 2($sp^2$) hybrid, the lobe for this $p$ orbital being perpendicular to the plane defined by the three 2($sp^2$) orbitals.

In the graphite structure, overlap occur between the 2($sp^2$) orbitals of neighboring atoms in the same plane. For such neighbors a side-to-side overlap also occur between their unhybridized $p$ orbitals. A side-to-side bonding known as $\pi$-bonding results between these neighbors. The electrons participating in this $\pi$-bonding seem able to move across these $\pi$-bonds from one atom to the next. This feature explains graphite's ability to conduct electricity along the sheets of carbon atom parallel to the (0001) direction. An in-plane nearest-neighbor distance is 1.421 Å. Normal to (0001), adjacent sheets of carbon atoms are held together by the weak Van der Waals bonds and separated by a distance 3.40 Å [Fig.2.2]. This gives softness to the structure [10,11].

The crystal structure is describes by hexagonal lattice with $D^4_{6h}$ ($P6_3/mmc$) space group.

Figure 2.1: Schematic presentation of sp$^3$ (left) and sp$^2$ (right) hybridization.

Figure 2.2: Diamond lattice (top): view from the $<$210$>$ direction (left), view from the $<$100$>$ direction (right). Graphite lattice (bottom): view from the $<$112$>$ direction (left), view from the $<$001$>$ direction (right).