As a materials science expert with extensive knowledge in the field of crystallography and electronic properties of materials, I'm well-equipped to discuss the differences in electrical conductivity between graphite and diamond, both of which are allotropes of carbon.

Graphite and diamond are distinct in their atomic structure, which is the fundamental reason for their differing electrical properties.
Graphite has a layered structure where each carbon atom is bonded to three other carbon atoms in a planar hexagonal lattice. This arrangement leaves one electron per carbon atom free to move within the layer, which contributes to graphite's ability to conduct electricity. These free electrons are delocalized over the entire layer, allowing them to move freely and thus enabling electrical conductivity.

In contrast, diamond has a tetrahedral structure where each carbon atom is covalently bonded to four other carbon atoms. This results in a rigid three-dimensional network where all four valence electrons of each carbon atom are engaged in strong covalent bonds. Because of this, there are no free or delocalized electrons available to move through the structure, which is why diamond does not conduct electricity.

The difference in bonding can be further explained by considering the sp2 and sp3 hybridization of carbon atoms in graphite and diamond, respectively. In graphite, the carbon atoms are sp2 hybridized, forming sigma bonds with three neighboring carbon atoms and a pi bond with the fourth electron in the p-orbital, which is free to move. This pi electron system is responsible for the electrical conductivity in graphite. On the other hand, in diamond, the carbon atoms are sp3 hybridized, with all four electrons participating in sigma bonds, leaving no free electrons for conduction.

Another aspect to consider is the physical properties of these materials. Graphite is soft and flaky, allowing the layers to slide over one another, which can disrupt the flow of electrons. However, within a single layer, the movement of electrons is relatively unimpeded. Diamond, being the hardest known natural material, has a rigid structure that does not allow for any such movement, further inhibiting the flow of electrons.

In summary, the electrical conductivity of graphite is due to the presence of delocalized electrons within its planar layers, while diamond's lack of conductivity is attributed to the absence of free electrons due to its fully saturated covalent bonding in a three-dimensional network.

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Graphite can conduct electricity because of the delocalised (free) electrons in its structure. These arise because each carbon atom is only bonded to 3 other carbon atoms. ... However, in diamond, all 4 outer electrons on each carbon atom are used in covalent bonding, so there are no delocalised electrons.read more >>