Graphite is a naturally occurring form of carbon, widely recognized for its unique properties. It finds applications in everyday items like pencil lead and industrial uses such as lubricants and electrodes in batteries. Unlike many other non-metals, graphite is an effective conductor of electricity, a characteristic that makes it valuable. This ability to conduct electricity stems directly from its specific atomic structure and the behavior of its electrons.
The Atomic Arrangement of Graphite
Graphite is composed of carbon atoms arranged in distinct layers. Within each of these layers, carbon atoms are strongly bonded together in hexagonal rings, forming flat sheets. These strong connections within the layers are a result of a type of bonding known as sp2 hybridization, where each carbon atom forms three covalent bonds with its neighboring carbon atoms. While bonds within these layers are very strong, the forces holding the layers together are much weaker. These weak Van der Waals forces allow the layers to easily slide past one another, contributing to graphite’s soft and slippery texture.
Delocalized Electrons and Electrical Flow
The electrical conductivity of graphite arises from the unique arrangement of its electrons. In graphite, each carbon atom uses three of its four outer electrons to form covalent bonds with three adjacent carbon atoms within its layer. One valence electron per carbon atom is not involved in these direct bonds. These unbonded electrons occupy unhybridized p-orbitals, which extend above and below the hexagonal planes of the carbon layers.
These p-orbitals from neighboring carbon atoms overlap, creating a continuous “sea” of electrons that are not confined to any single atom or bond. This collection of electrons is referred to as delocalized electrons, free to move throughout the entire layer. When an electrical voltage is applied, these mobile delocalized electrons flow easily across the layers, allowing graphite to conduct electricity.
Why Graphite Conducts, Unlike Diamond
To understand graphite’s conductivity, it helps to compare it with diamond, another form of carbon that behaves very differently. Diamond has a three-dimensional structure where each carbon atom is bonded to four other carbon atoms in a tetrahedral arrangement. All four of carbon’s valence electrons are used in forming strong covalent bonds, a type of bonding called sp3 hybridization. Because all electrons in diamond are tightly held within these localized covalent bonds, there are no free or delocalized electrons available to carry an electrical charge. This absence of mobile electrons means that diamond is an electrical insulator.