Question

Graphite is a good conductor of electricity but diamond is nonconductor of electricity because (1) Diamond is hard and graphite is soft. (2) Diamond and graphite have different atomic configurations. (3) Graphite is composed of positively charged carbon ions. (4) Graphite has a hexagonal layer structure with mobile \(\pi\) -electrons while diamond has a continuous tetrahedral covalent structure with no free electrons.

Short answer

Option (4): Graphite has mobile \( \pi \) -electrons for conductivity, while diamond has no free electrons due to its tetrahedral covalent structure.

Step-by-step solution

- Understand the question

The question asks for an explanation of why graphite conducts electricity while diamond does not, based on their properties and structures.

- Analyze the options

Read through each option and analyze it to see if it correctly explains the reason for the electrical conductivity differences between graphite and diamond.

- Evaluate Option (1)

Option (1): 'Diamond is hard and graphite is soft.' This statement does not provide an explanation for electrical conductivity differences. Hardness and softness are physical properties, not directly related to electrical conductivity.

- Evaluate Option (2)

Option (2): 'Diamond and graphite have different atomic configurations.' While this is true, it is not specific enough to explain why graphite conducts electricity and diamond does not. It lacks details about how the configurations affect electrical conductivity.

- Evaluate Option (3)

Option (3): 'Graphite is composed of positively charged carbon ions.' This statement is incorrect. Graphite consists of neutral carbon atoms arranged in layers, not positively charged ions.

- Evaluate Option (4)

Option (4): 'Graphite has a hexagonal layer structure with mobile \( \pi \) -electrons while diamond has a continuous tetrahedral covalent structure with no free electrons.' This option accurately describes why graphite conducts electricity and diamond does not. The \( \pi \) -electrons in graphite are mobile, allowing for electrical conductivity, whereas diamond's electrons are bound in covalent bonds, preventing conductivity.

- Conclusion

Based on the evaluations, option (4) provides the correct reasoning for the differences in electrical conductivity between graphite and diamond.

Key concepts

graphite structure

Graphite is one of the many allotropes of carbon. The structure of graphite is unique and plays a crucial role in its electrical conductivity properties. Graphite consists of layers of hexagonally arranged carbon atoms. These layers are stacked on top of each other in a manner where one layer slides over the other easily.
The carbon atoms within each layer are bonded strongly through covalent bonds, forming what looks like a honeycomb pattern. However, the forces holding these layers together, known as Van der Waals forces, are relatively weak. This special layered structure is the key to graphite being soft and slippery.

diamond structure

Unlike graphite, diamond has a completely different arrangement of carbon atoms. In diamond, each carbon atom is tetrahedrally bonded to four other carbon atoms. This structure forms a very rigid 3D network.
These covalent bonds are extremely strong and extend throughout the entire crystal, making diamond one of the hardest known materials. This rigid and closely packed arrangement means diamond has no free electrons to conduct electricity.

pi-electrons mobility

One of the defining features of graphite is the presence of mobile \( \pi \)-electrons. In each layer of graphite, every carbon atom forms three covalent bonds with its neighboring carbon atoms. This leaves one electron per carbon atom that is not used in bonding and is free to move within the layer.
These delocalized \pi \-electrons can move across the layers of graphite, facilitating the conduction of electricity. This electron mobility is why graphite can conduct electricity so efficiently.

tetrahedral covalent bonds

In diamond, each carbon atom forms four covalent bonds in a tetrahedral arrangement. These bonds are not just strong but also spatially oriented to each form angles of 109.5 degrees. This results in an extremely hard and very stable structure.
However, this network of tetrahedral bonds means that all the electrons in diamond are locked in place and not free to move. Without any free electrons, diamond cannot conduct electricity, making it an electrical insulator.

carbon atomic configurations

The atomic configuration of carbon determines its properties. Carbon atoms can be arranged in different ways to form various allotropes like graphite and diamond.
In graphite, the carbon atoms form layers with delocalized \( \pi \)-electrons, allowing for electrical conductivity. In diamond, the tetrahedral bonding arrangement locks electrons in place, preventing any electrical conductivity.
Understanding these configurations helps explain why different forms of the same element can have vastly different physical properties.