Understanding Properties of Graphite

Graphite is an allotrope of carbon, meaning it is a different structural form of the same element. It is well-known for its unique properties, such as being soft, a good conductor of electricity, and having a high melting point. These properties are a direct result of its specific atomic structure and bonding.

Graphite Structure and Bonding

The structure of graphite consists of layers of carbon atoms arranged in a hexagonal lattice. Within each layer, every carbon atom is strongly bonded to three other carbon atoms. These strong bonds within the layer are covalent.

• Each layer is flat and composed of repeating hexagonal rings of carbon atoms.
• The distance between carbon atoms within a layer is relatively short, indicating strong covalent bonds.
• The layers are stacked on top of each other, but the forces between layers are much weaker van der Waals forces. This explains why graphite is soft and the layers can slide over each other (making it useful as a lubricant).

Hybridization of Carbon in Graphite

To understand the bonding within the graphite layers, we look at the hybridization of the carbon atoms. Since each carbon atom in a layer is bonded to three other carbon atoms, it forms three sigma ($\sigma$) bonds. Carbon has an electron configuration of $1s^2 2s^2 2p^2$. In bonding, it typically undergoes hybridization involving its $2s$ and $2p$ orbitals.

In graphite, each carbon atom undergoes $sp^2$ hybridization. This means that one $2s$ orbital and two $2p$ orbitals combine to form three $sp^2$ hybrid orbitals. These three $sp^2$ orbitals lie in a plane at $120^\circ$ to each other and form the sigma bonds with the three neighbouring carbon atoms within the layer.

After $sp^2$ hybridization, one $2p$ orbital remains unhybridized on each carbon atom. This unhybridized $2p$ orbital is perpendicular to the plane of the $sp^2$ hybrid orbitals and the graphite layer.

Electron Delocalization in Graphite

The unhybridized $2p$ orbitals on adjacent carbon atoms in the graphite layer can overlap laterally. This lateral overlap results in the formation of a delocalized pi ($\pi$) electron system extending over the entire layer. These delocalized electrons are free to move within the layer, which is why graphite is a good conductor of electricity along the layers.

Analyzing the Given Properties

Let's examine each statement based on our understanding of graphite's structure and bonding:

1. Hybridisation of each carbon atom is sp ³

   As discussed, carbon atoms in graphite are bonded to three neighbours and undergo $sp^2$ hybridization. $sp^3$ hybridization occurs when a carbon atom forms four single bonds, like in diamond. Therefore, this statement is NOT true for graphite.

2. Hybridisation of each carbon atom is sp ²

   This statement is true for graphite, as each carbon forms three sigma bonds within the layer using $sp^2$ hybrid orbitals.

3. Electrons are delocalized over the whole sheet of atoms

   This statement is true. The unhybridized $2p$ orbitals on each carbon atom overlap to form a delocalized $\pi$ system across the entire layer, allowing for electrical conductivity.

4. Each layer is composed of hexagonal rings.

   This statement is true. The fundamental structural unit within each graphite layer is a hexagonal ring of carbon atoms linked by covalent bonds.

The question asks which property is NOT true for graphite. Based on our analysis, the statement that the hybridisation of each carbon atom is $sp^3$ is incorrect for graphite.

Identifying the Property NOT True for Graphite

Comparing the statements with the known properties of graphite, we find that the statement "Hybridisation of each carbon atom is $sp^3$" does not describe graphite. Graphite carbon atoms are $sp^2$ hybridized. Therefore, this property is not true for graphite.

Thus, the property that is NOT true for graphite is that the hybridisation of each carbon atom is $sp^3$.

Revision Table: Key Graphite Properties

Additional Information: Comparing Graphite and Diamond

Graphite and diamond are both allotropes of carbon, but their dramatically different properties arise from their different structures and bonding.

• Diamond: Each carbon atom is $sp^3$ hybridized and bonded to four other carbon atoms in a tetrahedral arrangement. This forms a rigid, 3D network structure. All valence electrons are localized in strong $\sigma$ bonds. Diamond is very hard, is an excellent electrical insulator, and has a high melting point.
• Graphite: As discussed, carbon atoms are $sp^2$ hybridized, forming layered hexagonal structures with delocalized $\pi$ electrons. Graphite is soft, is a good electrical conductor, and also has a high melting point (due to strong covalent bonds within layers).

Understanding the hybridization of carbon atoms (sp² vs sp³) is key to explaining the contrasting properties of these two important carbon allotropes.