Question

In graphite, \(\mathrm{C}\) atom is in.........state. (a) \(s p^{3}\) (b) \(s p\) (c) \(s p^{2}\) (d) None of these.

Short answer

In graphite, a \text{C} atom is in the \text{sp}^{2} state, so the answer is (c) sp².

Step-by-step solution

Understand the Structure of Graphite

Graphite is composed of layers of carbon atoms. Each carbon atom is bonded to three other carbon atoms in the same layer, forming a hexagonal lattice structure. The fourth valence electron is not localized but is part of a delocalized pi system that allows conduction of electricity.

Determine the Hybridization

Since each carbon atom is bonded to three other carbon atoms and has one non-bonded valence electron participating in a pi bond, this implies that three orbitals (one s orbital and two p orbitals) are hybridized to form three sigma bonds. Therefore, the carbon atom's hybridization in graphite is sp².

Key concepts

sp² Hybridization

Understanding the concept of sp² hybridization is crucial to explaining the structure of graphite, a form of carbon. In chemistry, hybridization is the process where atomic orbitals mix to form new hybrid orbitals, which can influence molecular geometry and bonding properties.

When we discuss sp² hybridization in carbon atoms, we are referring to the mixing of one s orbital and two p orbitals. This process generates three sp² hybrid orbitals that lie on the same plane with a 120-degree angle between them. Each of these has an unpaired electron capable of forming a strong covalent bond, known as a sigma bond.

Carbon atoms with sp² hybridization readily form planar structures. In graphite, every carbon atom uses these three sp² orbitals to form sigma bonds with three neighboring carbon atoms, resulting in a flat, hexagonal lattice structure. This hexagonal lattice is what makes graphite's layers so strong, yet the layers can slide over each other easily, leading to graphite's lubricating qualities.

Carbon Atomic Structure

The carbon atomic structure in graphite reveals much about its physical properties and uses. Carbon, with its atomic number 6, has a configuration of 2s² 2p² in its ground state, which means it has four valent electrons. In graphite, each carbon atom undergoes hybridization, adjusting its electron configuration to maximize bonding efficiency.

The unique aspect of the carbon atomic structure in graphite is the presence of both sp² hybridized sigma bonds and delocalized pi bonds. Each carbon atom contributes one electron to a delocalized system of pi electrons that resides above and below the planar carbon layers. These pi electrons are free to move across the plane, which gifts graphite its ability to conduct electricity, unlike diamond, another form of carbon, which is an excellent electrical insulator.

Sigma and Pi Bonds

The bonding in graphite can be divided into sigma (σ) and pi (π) bonds, each crucial for the material's properties. Sigma bonds are formed by the head-on overlap of orbitals, such as the sp² hybrid orbitals of carbon in graphite. They create a strong bond that defines the basic shape of the molecule.

In contrast, pi bonds are formed by the side-on overlap of p orbitals, allowing for electron density above and below the atomic plane. In graphite, each carbon atom is involved in forming three sigma bonds with neighboring carbons, creating a stable planar configuration. Additionally, the remaining p orbital of each carbon atom, perpendicular to the plane of sp² orbitals, contributes to a pi bond. The electrons in these pi bonds are delocalized and are less tightly bound than those in sigma bonds, offering graphite its ability to conduct electricity and its lesser overall bond strength, which results in the ease with which the layers of graphite separate.