Graphite is one of the most well-known examples of a conductive non-metal. Its ability to conduct electricity makes it useful in electrodes, batteries, and industrial processes. For IB Chemistry students, understanding why graphite conducts is essential in bonding, hybridization, and structure–property relationships. This article breaks down the explanation clearly and links it to concepts you need for your exams.
The Structure of Graphite
Graphite is made of carbon atoms arranged in layers.
Each carbon atom uses three of its electrons to form three sigma (σ) bonds with neighboring carbon atoms.
This creates:
- Flat hexagonal sheets
- Strong covalent bonds within the layers
- Weak forces between layers
These layers can slide over each other, giving graphite its lubricating and soft properties.
The Key Reason: Delocalized Electrons
Each carbon atom in graphite has four valence electrons, but uses only three for bonding.
The remaining one electron per carbon atom becomes delocalized.
These delocalized electrons:
- Are free to move throughout the entire layer
- Act as mobile charge carriers
- Allow electrical conductivity similar to metals
This is the fundamental reason graphite conducts electricity: it contains free-moving electrons.
Hybridization in Graphite (IB Concept)
Graphite is sp² hybridized.
This means:
- Three sp² orbitals form σ-bonds
- One unhybridized p-orbital remains
