Hybridisation of Graphite

Carbon has different allotropes, such as graphite and diamond. Graphite is soft, opaque, and can conduct electricity; on the other hand, diamond is hard, transparent, and cannot conduct electricity. Both are pure carbon, but still, they have different properties. It is because of the arrangement of carbon atoms in the molecule. The process of hybridisation of carbon forms graphite. Hybridisation is when two or more atomic orbitals are combined to form a new hybridised orbital. Graphite has layers of atoms of carbon that are put together in six-membered hexagonal rings. Carbon atoms undergo sp² hybridisation to form graphite.

The Hybridisation of Graphite

Carbon atoms undergo sp² hybridisation to form graphite. In sp² hybridisation, one s orbital and two p orbitals combine to form three sp² orbitals. Carbon has a general electronic configuration of (1s2, 2s2, 2p2) thereby spreading the four valence electrons in the s and p orbitals. 

•  These newly formed orbitals connect each atom of carbon to the other three atoms by a covalent bond.
• The structure made by these atoms is a trigonal planar geometry and the C-C-C bond formed has a 120° angle.
•  The hybridised carbon atoms thus formed are put together in several layers, and in those layers, they are arranged in a hexagonal arrangement.

Graphite has different properties as the layers have a weak force between them. This weak force could be the reason that graphite is highly anisotropic, with mechanical, electrical, and thermal properties that are very dissimilar within in-plane and out-of-plane directions.

Molecular Geometry of Graphite

As we know, the hybridisation of graphite is an sp² hybridisation. The graphite structure is a trigonal planar, and the atoms have a bond angle of 120° between them. 

• Graphite has layers of carbon arranged in hexagonal rings along with six members. These hexagonal rings are connected at their edges. 
• An endless sequence of fused benzene rings could be used to make layers of linked rings.
•  Carbon atoms arranged in these rings are in an sp² hybridised state. 
• Each carbon atom is connected with other carbon atoms in this type of molecular orbital model.

Graphene layers are singular sheets of carbon atoms, and the ring arrays are organised in large sheets.

•  It could also mean that the carbon sheets are very powerfully bonded. However, these sheets interact with each other weakly. 
• The sheets are not connected through covalent bonds; their atoms interact via London dispersion forces. 
• The graphene layers are set on top of one another parallel to the “C” crystallographic axis of the hexagonal 4-axis system where graphite gets crystallised.

Carbon atoms are capable of forming spherical molecules that are known as buckminsterfullerenes. However, it also raises the question of why there are no covalent bonds between graphite sheets? 

• Diamond has a three-dimensional bond, but graphite and graphene have two-dimensional bonds. 
• As we know, to form four bonds attached to each carbon atom, it is required to hybridise four atomic orbitals. 
• We may think that only three bonds are needed in graphite as each carbon is attached with three others, which is not entirely true.

Hybridisation of Diamond

Diamond is formed as a result of sp3 hybridisation. Here, one s orbital, and three p orbitals are mixed to form four sp3 orbitals.

•  Sp3 hybridisation is tetrahedral hybridisation.
•  These orbitals form a 109.28-degree angle.
•  Carbon atoms are connected through a powerful covalent bond.
•  Each carbon atom is bonded with the other four adjacent atoms in a  diamond. 
• Each of these atoms is connected together with the other four of their adjacent atoms.
• Hence, a tough three-dimensional structure is formed, which gives the diamond its hardness.

Conclusion

Graphite is one of the different allotropes of carbon. Allotropes are various forms of the same element, but they have other properties. Graphite is soft, opaque, black, or grey and can conduct electricity. However, diamond is also an allotrope of carbon, but its properties completely differ from graphite. The hybridisation of graphite is an sp² hybridisation with a bond angle of 120°, while in diamond, it is sp3 hybridisation with a 109.28-degree bond angle. Also, carbon atoms in diamonds have a three-dimensional network, while in graphite, it is two-dimensional. These bonds give a diamond its hardness and make graphite a soft material.