Question

Are the fullerenes network covalent solids? What makes them different from diamond and graphite? It has been shown that carbon can form chains in which every other carbon atom is bonded to the next carbon atom by a triple bond. Is this allotrope of carbon a network covalent solid? Explain.

Short answer

No, Fullerenes and the hypothetical carbon allotrope are not network covalent solids. They do not form a giant network or lattice structures like diamond or graphite. Instead, they have distinct separate structures.

Step-by-step solution

Understanding Fullerenes

Fullerenes are a form of carbon molecule where carbon atoms are connected to form a hollow sphere or tube. They do not form an extended giant structure or network, and thus, they do not meet the criteria of a network covalent solid.

Comparison with Diamond and Graphite

Diamond and graphite are both network covalent solids. In a diamond, each carbon atom is bonded to four others, forming a giant network. In graphite, each carbon atom is bonded to three others forming layered structures. Fullerenes differ from diamond and graphite as they do not form such extended giant structures or networks.

Analyzing the Hypothetical Carbon Allotrope

In the given hypothetical carbon allotrope where every other carbon atom is bonded to the next carbon atom with a triple bond, the carbon atoms are not arranged in an extended network. Therefore, this does not meet the criteria of a network covalent solid.

Key concepts

Fullerenes

Fullerenes are exceptionally fascinating carbon structures that captivate many scientists and educators alike. Unlike other carbon allotropes, fullerenes have a distinct architecture composed of carbon atoms linked to form hollow spheres, ellipsoids, or tubes. These molecules are also known as 'Buckyballs' or 'Carbon nanotubes' when shaped like a tube.

What sets fullerenes apart from network solids is their discrete molecular form, which means they're composed of individual molecules rather than continuous networks. Therefore, fullerenes do not constitute network covalent solids despite their strong covalent bonds. As molecular structures, they have distinct properties, like the ability to enclose other atoms or molecules, leading to applications in drug delivery, and potential use in electronic circuits.

Carbon Allotropes

Carbon is a versatile element and forms various allotropes with vastly different properties. An allotrope is a form of an element with a different physical structure, creating unique properties. The most well-known carbon allotropes include diamond, graphite, and fullerenes. Each allotrope exhibits distinct bonding and structures: diamonds have a rigid 3D lattice, graphite has a layered 2D structure, and fullerenes have a closed 3D spherical or cylindrical shape.

The existence of these different allotropes exemplifies carbon's ability to form sp, sp², and sp³ hybridized bonds, leading to structures ranging from the hardest materials to excellent electrical conductors. Understanding the various forms of carbon lays the foundation for exploring the diversity of bonding in solids.

Bonding in Solids

Solid materials often have their atoms, ions, or molecules bonded in specific organized patterns. The way these particles are bonded significantly influences the properties of the material. There are primarily four types of bonding in solids: ionic, covalent, metallic, and molecular.

Network covalent solids, such as diamond, are structures where atoms are bonded covalently in a continuous network extending throughout the material. This type of bonding confers high melting points, hardness, and low electrical conductivity, with diamond being a prime example. In contrast, molecular solids like fullerenes consist of discrete molecular units held together by weaker intermolecular forces, leading to lower melting points and different physical properties.

Comparison Between Diamond, Graphite, and Fullerenes

When comparing the three allotropes of carbon: diamond, graphite, and fullerenes, the key differences lie in their bonding and the resulting properties. Diamond's tetrahedral sp³ bonding creates a rigid 3D structure, making it the hardest known natural material. Graphite, with sp² bonding, forms layers that can slide over each other due to weak Van der Waals forces, making it an excellent lubricant and conductor of electricity.

Fullerenes, on the other hand, have a molecular structure with sp² bonding, giving them unique properties such as the ability to trap other molecules inside their hollow structure. They are neither as hard as diamond nor as conductive as graphite but offer potential in areas such as nanotechnology and material science. These differences illustrate how variations in bonding and structure can lead to a diverse array of materials with specialized functions.