Explain the difference between Diamond and graphite ?

The key differences between diamond and graphite can be explained based on their atomic structure, electrical conductivity, and bonding properties. Below is a detailed comparison to help differentiate between diamond and graphite:

1. Hybridization of Carbon

In diamond, carbon atoms undergo sp3 hybridization, forming a strong, three-dimensional covalent network. This arrangement results in each carbon atom being bonded to four other carbon atoms, creating a rigid, hard structure.

In contrast, graphite consists of carbon atoms that undergo sp2 hybridization, forming a two-dimensional structure of layers. Each carbon atom is bonded to three other carbon atoms within a plane, leaving one electron free for conduction between layers.

2. Electrical Conductivity

Diamond is a bad conductor of electricity because all of its valence electrons are tightly bonded in the crystal structure and are not free to move. This makes diamond an insulator.

On the other hand, graphite is an excellent conductor of electricity. The free electrons in the sp2 hybridized layers of graphite can move along the planes, allowing graphite to conduct electricity efficiently.

3. Structural Arrangement

Diamond has a giant three-dimensional polymeric structure, where each carbon atom is bonded in a tetrahedral arrangement, creating a rigid and compact lattice. This structure gives diamond its hardness and high melting point.

Graphite, in contrast, has a two-dimensional layer structure, where each layer consists of carbon atoms bonded in a hexagonal arrangement. The layers are weakly held together by van der Waals forces, allowing the layers to slide over each other, making graphite slippery and soft.

4. Bond Length and Bond Angles

The C–C bond length in diamond is approximately 1.54 Å, and the bond angle between the carbon atoms is 109.28°, typical for sp3 hybridization in a tetrahedral geometry.

In graphite, the C–C bond length is slightly shorter, around 1.42 Å, with bond angles of 120° due to the sp2 hybridization and the planar arrangement of atoms.

Conclusion

To differentiate between diamond and graphite, it is essential to look at the bonding, structure, and electrical conductivity. Diamond, with its sp3 hybridization and three-dimensional structure, is a hard insulator, while graphite, with sp2 hybridization and a layered structure, is soft and an excellent conductor of electricity.

Diamond Structure :

Graphite Structure :