In graphite, each carbon atom is bonded to three neighboring carbon atoms, forming a hexagonal lattice structure. Each carbon atom in graphite has a valency of 4, meaning it can form up to four covalent bonds. However, in the case of graphite, only three out of the four valence electrons of each carbon atom participate in bonding with other carbon atoms, leaving one electron uninvolved in bonding. This uninvolved electron in each carbon atom is known as a pi (π) electron.
The three covalent bonds in graphite are formed by each carbon atom's three valence electrons, which are used to share electrons with the neighboring carbon atoms. These bonds create a two-dimensional hexagonal network of carbon atoms, forming the hexagonal parallel planes characteristic of graphite.
The uninvolved pi (π) electron, which is not involved in bonding, is located above and below the plane of the hexagonal lattice. These uninvolved electrons give rise to unique properties of graphite, such as its ability to conduct electricity along the parallel planes and its characteristic slippery and lubricating nature.
So, in graphite, while the carbon atom's valency of 4 is not completely satisfied in terms of forming covalent bonds with four other atoms, it is still stable due to the strong covalent bonding formed by three of its valence electrons in the hexagonal lattice structure.