Graphite is a material found in many everyday objects, from pencil lead to industrial lubricants. This article clarifies what graphite is made of, exploring its atomic architecture and unique characteristics.
The Carbon Foundation
Graphite is an allotrope of carbon, meaning it is one of several distinct structural forms that carbon atoms can take. Other familiar allotropes include diamond and graphene. Carbon is the sole elemental component of graphite, showcasing carbon’s remarkable ability to bond with itself in various configurations.
Layered Atomic Structure
The atomic arrangement within graphite is highly ordered. Carbon atoms form flat, two-dimensional sheets, often referred to as graphene layers. Within each of these layers, carbon atoms are arranged in hexagonal rings, similar to a honeycomb lattice.
Each carbon atom within a layer forms strong covalent bonds with three other carbon atoms, creating a robust and continuous network. These strong intralayer bonds result in a carbon-carbon bond length of approximately 0.142 nanometers within each graphene sheet. The individual graphene layers are then stacked on top of each other. However, the forces holding these layers together are significantly weaker than the covalent bonds within the layers.
Weak intermolecular forces, specifically Van der Waals forces, exist between these layers. These forces allow the layers to be separated relatively easily, with an interlayer distance of about 0.335 nanometers.
Distinctive Physical Properties
Graphite’s unique layered atomic structure directly accounts for its characteristic physical properties. The weak Van der Waals forces between the graphene layers allow them to slide past each other with minimal resistance. This explains graphite’s notable softness, its greasy or slippery feel, and its widespread use as a dry lubricant. When a pencil marks paper, it is essentially thin layers of graphite shedding from the main structure.
Graphite is also an excellent electrical conductor. Within each layer, one valence electron from each carbon atom is not involved in covalent bonding and becomes delocalized, meaning it is free to move across the entire sheet. These mobile electrons allow graphite to conduct electricity efficiently along the planes of the layers.
Graphite also possesses a high melting point, reaching approximately 3,650°C, due to the extensive network of strong covalent bonds that must be broken within its layers for it to change state. Unlike diamond, another carbon allotrope where all valence electrons are rigidly bonded, graphite’s structure allows for both conductivity and a layered, opaque appearance.
Formation and Natural Occurrence
Natural graphite forms primarily through metamorphic processes deep within the Earth’s crust. This occurs when carbon-rich materials, such as organic sediments, are subjected to intense heat and pressure over long geological periods. The high temperatures and pressures cause the carbon atoms to rearrange into the stable, layered structure of graphite.
Large deposits of natural graphite are found globally, with significant sources including China, Mexico, Canada, Brazil, and Madagascar. Beyond natural occurrence, synthetic graphite is also produced on a large scale. This involves heating carbonaceous raw materials, such as petroleum coke or coal tar pitch, to extremely high temperatures, often exceeding 2,100 °C, in a process called graphitization. This industrial process forces the carbon atoms to organize into the graphite structure, resulting in a material often of higher purity than its natural counterpart.