Question

Which of the following statements is not correct? (1) Diamond has a 3D polymeric structure while graphite has a layered lattice structure. (2) Diamond is an insulator while graphite is a good conductor of electricity. (3) Diamond is a good conductor of heat than graphite. (4) Diamond and graphite have same densities.

Short answer

Statement 4 is incorrect.

Step-by-step solution

- Understand the Structures

First, compare the structures of diamond and graphite. Diamond has a 3D polymeric structure where each carbon atom is tetrahedrally bonded to four other carbon atoms. Graphite has a layered structure with each layer consisting of carbon atoms in a hexagonal arrangement, and layers are held together by weak van der Waals forces.

- Electrical Conductivity

Next, analyze the electrical conductivity of both substances. Diamond has no free electrons, making it an electrical insulator. On the other hand, graphite has free electrons between its layers, making it a good conductor of electricity.

- Heat Conductivity

Consider the heat conductivity properties. Diamond is known to be a very good conductor of heat due to its strong covalent bonding and rigid lattice structure. Graphite, although it conducts heat, does so less efficiently than diamond.

- Compare Densities

Finally, compare the densities of diamond and graphite. Diamond has a higher density (approximately 3.51 g/cm³) compared to graphite (approximately 2.26 g/cm³) because of its more compact tetrahedral bonding network.

- Identify the Incorrect Statement

Review each statement in light of the above properties. The only incorrect statement is that diamond and graphite have the same densities, as they do not; diamond is denser than graphite.

Key concepts

Diamond Structure

Diamond has a unique 3D structure. Every carbon atom in diamond is covalently bonded to four other carbon atoms in a tetrahedral shape. This creates a very strong and rigid lattice structure. Because of these strong covalent bonds, diamond is incredibly hard and has a high melting point.
Furthermore, each carbon atom in a diamond is symmetrically surrounded by others, contributing to its remarkable physical properties.

Graphite Structure

Graphite has a very different structure from diamond. In graphite, carbon atoms are arranged in a layered configuration. Each layer is made up of carbon atoms bonded together in hexagonal patterns.
These layers are held together by weak van der Waals forces, which allow them to slide over each other easily.
This gives graphite its slippery feel and is why it is used as a lubricant and in pencils. Additionally, the layers can easily separate, unlike the stiff, interlocked structure of diamond.

Electrical Conductivity

The electrical conductivity of diamond and graphite varies significantly. Diamond does not conduct electricity because all of its electrons are involved in strong covalent bonds. There are no free electrons to carry electric current.
On the other hand, graphite is a good conductor of electricity. This is because, within its layers, each carbon atom forms only three bonds. The fourth electron is free to move, allowing graphite to conduct electricity. This property is utilized in applications such as batteries and electrodes.

Heat Conductivity

Diamond excels at conducting heat. Its rigid lattice structure allows for efficient transfer of thermal energy. Its strong covalent bonds and symmetrical lattice provide a pathway for heat to travel quickly through the material.
In contrast, graphite, while capable of conducting heat, is less effective in comparison. Its layered structure does not promote as efficient thermal conductivity as diamond. Still, graphite's ability to conduct heat is beneficial in various industrial applications.

Material Density

The density of a material relates to how closely packed its atoms are. Diamond has a density of about 3.51 g/cm³. This high density results from its tightly packed 3D tetrahedral arrangement of carbon atoms.
Graphite, however, has a lower density of around 2.26 g/cm³. This difference arises because graphite's layered structure is less compact. The layers are held together by much weaker forces, creating more space within the material and thus a lower density.
It is important to understand that this density difference affects many of the physical properties and uses of these carbon allotropes.