Question

Graphite conducts electricity because of (1) weak Van der Waal's forces between layers (2) covalent bonding between carbon atoms of layers (3) delocalized clectrons in cach laycr (4) sp \(^{2}\) hybridisation of carbon atom in cach layer

Short answer

Graphite conducts electricity because of delocalized electrons in each layer.

Step-by-step solution

- Analyze the structure of graphite

Graphite has a layered structure where carbon atoms are arranged in hexagonal layers. Each carbon atom is bonded to three others in the same layer via covalent bonds, forming a network of hexagons.

- Understand the type of bonding between layers

Between these layers, weak Van der Waal’s forces hold the layers together. These forces allow the layers to slide past each other easily.

- Identify the type of electron arrangement

Within each layer, each carbon atom has one free electron that is not used in bonding. These free electrons become delocalized over the entire layer, allowing them to move freely.

- Determine the cause of electrical conductivity

The delocalized electrons in each layer of carbon atoms can move freely, carrying an electric charge and allowing graphite to conduct electricity.

- Choose the correct option

Based on the analysis, the correct reason for graphite's conductivity is the presence of delocalized electrons in each layer.

Key concepts

delocalized electrons

Graphite’s ability to conduct electricity stems from the presence of delocalized electrons. In each graphite layer, carbon atoms form a hexagonal arrangement. Each carbon atom uses three of its four available electrons to form covalent bonds with three other carbon atoms.

This leaves one electron per carbon atom unbonded. These unbonded electrons are free to move within the layer. Such electrons are termed as 'delocalized electrons'.📈

Delocalized electrons can move throughout the graphite layer, making them essential for electrical conductivity. When an electric potential is applied, these electrons carry the charge through the material.

• Delocalized electrons spread out over the entire layer
• Enable the flow of electric current
• Key to graphite's conductive properties

carbon atom bonding

The bonding between carbon atoms in graphite is crucial for its properties. Carbon atoms in a graphite layer are bonded to three other carbon atoms through covalent bonds. This creates a planar, hexagonal structure.

Each carbon atom undergoes a process known as 'sp² hybridization'. This means one s orbital and two p orbitals mix to form three sp² hybrid orbitals. These orbitals overlap with those of adjacent carbon atoms to form strong covalent bonds.

Here are key points about carbon atom bonding in graphite:

• Hexagonal arrangement due to covalent bonding
• One free electron per carbon atom
• sp² hybridization leads to strong in-plane bonds

This strong covalent bonding within the layers gives graphite its high melting point and mechanical strength within the layers.

Van der Waals forces

In graphite, different layers of carbon atoms are held together by weak Van der Waals forces. These forces are much weaker than covalent bonds.

Van der Waals forces arise from temporary attractions between molecules. In graphite, this force occurs between the layers, allowing them to slide over each other easily.

This is why graphite feels slippery and is used in applications like lubricants and pencil lead.

• Weak forces between the layers allow sliding
• Essential for graphite’s lubricating properties
• Weaker than covalent bonds but crucial for graphite's structure

The presence of Van der Waals forces ensures that while the layers hold together, they are not rigid, giving graphite its unique properties.