The property of graphite to have layers that easily slide over one another can be attributed to the arrangement of its carbon atoms and the behavior of electrons within its structure.
Graphite is a form of carbon where carbon atoms are arranged in a layered structure. Each carbon atom in graphite forms strong covalent bonds with three neighboring carbon atoms, resulting in a hexagonal lattice structure within each layer. These layers are stacked on top of each other.
Within each layer, carbon atoms are strongly bonded and form a planar structure, but the bonding between layers is relatively weak. The weak bonding occurs due to the presence of delocalized electrons in graphite.
In the structure of graphite, each carbon atom has three out of its four valence electrons involved in forming strong covalent bonds within the layer. The remaining fourth valence electron of each carbon atom becomes delocalized and moves freely throughout the entire structure. These delocalized electrons are often referred to as π (pi) electrons.
The presence of these delocalized π electrons gives rise to unique properties in graphite. The π electrons are not tightly bound to any specific carbon atom or layer but can move easily within and between the layers. This results in weak interlayer forces, allowing the layers to slide over one another with relatively low resistance.
The sliding of graphite layers is facilitated by the weak van der Waals forces between the layers. These forces are relatively weak compared to the strong covalent bonds within each layer. As a result, when an external force is applied to graphite, the layers can easily slide over one another due to the weak interlayer forces, resulting in the lubricating property of graphite.
In summary, the ability of graphite layers to slide over one another is a consequence of the hexagonal lattice structure, the weak interlayer forces, and the presence of delocalized π electrons that allow for relatively free movement of electrons within and between the layers.