Graphite conducts electricity because it has delocalized electrons that are free to move along its layers, similar to how electrons move in metals.

Quick Scoop

Core idea: structure and electrons

  • Each carbon atom in graphite forms three strong covalent bonds in a hexagonal sheet, leaving one outer-shell electron not tied to a single bond.
  • That leftover electron becomes delocalized , meaning it is shared across the whole layer and can move when a voltage is applied, carrying charge as an electric current.

Why graphite but not diamond?

  • In graphite, atoms are arranged in flat layers with delocalized electrons in π\pi π-systems above and below the plane, so electrons can flow easily within each layer.
  • Diamond also consists of carbon, but each atom bonds to four neighbors in a rigid 3D network, using all outer electrons in bonds, so there are no free electrons and it does not conduct electricity.

Directional (anisotropic) conduction

  • Graphite conducts very well along the planes of its layers, because the delocalized electrons move easily across these flat sheets.
  • Between the layers, only weak forces hold them together, so electrons do not move as freely in this direction and conductivity is much lower.

Real-world uses

  • Because of this combination of good in-plane electrical conductivity and stability, graphite is widely used for electrodes in batteries, electric arc furnaces, and other electrical contacts.
  • Its high thermal stability and ability to conduct both heat and electricity also make it useful in high-temperature industrial processes and energy technologies.

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