Question

Network Covalent Atomic Solids Explain why the graphite structure of carbon allows graphite to be used as a lubricant, but the diamond structure of carbon does not.

Short answer

Graphite is a good lubricant due to its layers of carbon atoms held by weak van der Waals forces, allowing the layers to slide past one another. Diamond lacks such layers and is extremely rigid, meaning it cannot act as a lubricant.

Step-by-step solution

Understanding Graphite Structure

Graphite consists of layers of carbon atoms. Each carbon atom is bonded to three other carbon atoms in the same layer, forming a hexagonal array. These layers are held together by weaker van der Waals forces, allowing them to slide past one another with relative ease.

Understanding Diamond Structure

In diamond, each carbon atom is bonded tetrahedrally to four other carbon atoms. This creates a very strong and rigid three-dimensional network, resulting in an extremely hard structure.

Contrasting the Two Structures

The difference in bonding between graphite and diamond leads to distinct physical properties. Graphite's layers can slip because of the weak forces between them, making it a good lubricant. Diamond's rigid structure does not have planes that can slide easily, hence it is not a good lubricant.

Key concepts

Graphite Structure

Graphite is a fascinating allotrope of carbon that has unique properties due to its special structure. Each carbon atom in graphite is linked to three other carbon atoms in a plane, forming a hexagonal pattern that resembles a honeycomb structure. These planes of carbon, often referred to as graphene layers, stack on top of one other, creating the bulk material we know as graphite.

The bonds within each layer are incredibly strong — they are covalent bonds, robust and directional in nature. However, what's particularly interesting about graphite is the way these planes interact with each other. Unlike the strong covalent bonds within the layers, the forces between the layers are much weaker, known as van der Waals forces. These are intermolecular attractions that do not involve the sharing or transferring of electrons like in covalent or ionic bonds.

Because of these weaker forces, the layers can easily slide over each other. This property is what makes graphite an excellent lubricant, as the individual layers can move with little resistance, allowing materials to glide smoothly across one another. Not only does it serve as a lubricant, but this layered structure also makes graphite a good electrical conductor along the planes, although it conducts poorly perpendicular to the layers.

Diamond Structure

In stark contrast to the planar layers of graphite, diamond boasts an incredibly strong three-dimensional lattice. This structural difference is the secret behind diamond's ranking as the hardest natural substance known. Every carbon atom in diamond is tetrahedrally bonded to four other carbon atoms. This means they are linked in a pyramidal shape, with strong covalent bonds extending in all spatial directions, creating a highly interconnected network.

The tetrahedral bonding in diamond accounts for several of its remarkable characteristics. Diamonds are not only hard but also have a high refractive index, giving them their characteristic sparkle. The strength of the bonds throughout the structure makes it nearly impossible for the diamond to cleave in any particular direction, which means it lacks the planar slip systems present in graphite. As a result, diamonds are not suitable as a lubricant; instead, their hardness and thermal conductivity find uses in cutting tools and heat sinks respectively.

Understanding the tetrahedral bonding pattern is crucial for grasping why diamonds are such formidable materials. It's this spatially extensive covalent bonding that also makes diamonds have a low electrical conductivity, as there are no free moving electrons within the structure.

Van der Waals Forces

Van der Waals forces are the gentle whisper in the world of intermolecular attractions, acting between molecules without the need for chemical bonding. These forces arise due to temporary fluctuations in electron distribution within atoms and molecules, creating instant and ephemeral dipoles that attract one another. They are named after the Dutch physicist Johannes Diderik van der Waals, who was the first to describe these interactions.

There are three types of van der Waals forces: dipole-dipole interactions, London dispersion forces, and hydrogen bonds. While van der Waals forces are involved in numerous processes, they are particularly crucial in explaining the properties of molecules and solids where no stronger bonds are at play. For example, in the case of graphite, we've seen how these forces allow layers to slip over one another.

It's important to note that although these forces are generally weaker than both covalent and ionic bonds, their impact is significant in biological systems, contributing to the 3D structure of proteins and the adhesion of cells. Van der Waals forces play a silent yet indispensable role in many natural and synthetic materials, dictating physical properties like boiling points, melting points, and viscosity.