Question

Both diamond and graphite are network solids consist solely of carbon atoms: In diamond, every carbon atom is bonded to four other carbon atoms. In graphite, every carbon atom is bonded to three other carbon atoms, creating sheets of atoms that lie on top of one another. (a) Which type of intermolecular forces exist between the sheets of carbon atoms in graphite? (b) Diamond is an extremely hard substance, whereas graphite is a soft, slippery substance often used as a lubricant. Use intramolecular and intermolecular forces to explain this difference in physical properties.

Short answer

In summary, the intermolecular forces between graphite layers are weak van der Waals forces, while diamond lacks such forces entirely. The difference in physical properties between diamond and graphite arises from the disparity in their intramolecular forces (strong covalent bonds in a tetrahedral arrangement for diamond vs. covalent bonds in a hexagonal arrangement for graphite) and intermolecular forces (absent in diamond but present as van der Waals forces in graphite). These characteristics make diamond an extremely hard substance and graphite a soft, slippery substance.

Step-by-step solution

Understand the Structures of Diamond and Graphite

In diamond, each carbon atom is bonded to four other carbon atoms in a tetrahedral arrangement, forming a strong three-dimensional network. In graphite, each carbon atom is bonded to three other carbon atoms, forming a two-dimensional network of hexagonal sheets. These sheets are stacked on top of each other, separated by a larger distance than the bond length between carbon atoms within the same sheet.

(a) Identify Intermolecular Forces in Graphite

The main intermolecular force between the sheets of carbon atoms in graphite is van der Waals forces (London dispersion forces). These forces arise from temporary charge imbalances due to the movement of electrons, resulting in small instantaneous dipoles. These dipoles induce dipoles in the neighboring sheets, creating attractive forces between them.

(b) Understanding Intra- and Intermolecular Forces in Diamond and Graphite

In both diamond and graphite, carbon atoms are covalently bonded to other carbon atoms within their respective structures. These strong covalent bonds are known as intramolecular forces. - In diamond, there are strong covalent bonds in all directions, forming a robust three-dimensional network. This highly symmetrical and interconnected structure makes diamond an extremely hard substance. - In graphite, the intramolecular forces (covalent bonds) are limited to the two-dimensional hexagonal sheets. The sheets of graphite are only held together by weak van der Waals forces (intermolecular forces). As a result, graphite layers can slide across each other more easily, making graphite a soft and slippery substance often used as a lubricant.

Conclusion

In summary, - The intermolecular forces present between the sheets of carbon atoms in graphite are van der Waals forces. - The difference in physical properties between diamond and graphite arises from the disparity in their intramolecular forces (strong covalent bonds in a tetrahedral arrangement for diamond vs. covalent bonds in a hexagonal arrangement for graphite) and intermolecular forces (absent in diamond but present as van der Waals forces in graphite). These characteristics make diamond an extremely hard substance and graphite a soft, slippery substance.

Key concepts

Diamond and Graphite Structures

Understanding the structure of diamond and graphite is fundamental to grasping why these substances exhibit remarkably different physical properties despite both being forms of carbon.

In diamond, carbon atoms create a formidable lattice by covalently bonding with four other carbon atoms in a tetrahedral configuration. This leads to a durable three-dimensional network that extends throughout the entire crystal, making diamond one of the hardest known materials.

Contrastingly, graphite is structured in layers. Here, each carbon atom forms a plane by bonding to three other carbons in a hexagonal pattern. These planes, or sheets, are weakly interconnected by van der Waals forces, which permits them to slide over each other with ease. This unique structure makes graphite soft and slippery, which is why it finds use as a lubricant.

Van der Waals Forces

Van der Waals forces play a crucial role in the physical properties of substances such as graphite. They are a type of weak intermolecular force, including London dispersion forces, that occur between molecules or, in the case of graphite, between layers of atoms.

These forces emerge from temporary fluctuations in electron distribution within atoms or molecules, which produce momentary dipoles. A dipole in one molecule can induce a dipole in a neighboring molecule, resulting in an attraction between them. These attractions, although weak compared to covalent or ionic bonds, are significant in substances with a large number of atoms or in situations where molecules or atom layers come into close proximity, as in the layers of graphite.

Despite their relatively low strength, van der Waals forces are pivotal in determining the physical behaviors of many materials, including biological molecules and nanomaterials.

Covalent Bonds in Chemistry

At the heart of both diamond and graphite structures lie covalent bonds, the strong intramolecular forces that hold atoms together within a molecule or a crystal lattice.

A covalent bond is formed when two atoms share a pair of electrons, allowing them to attain a more stable electronic configuration. In diamond, the sharing is between one carbon atom and four others, which results in a rigid, three-dimensional framework.

In graphite, the covalent bonding occurs between a carbon atom and three nearby atoms, creating flat sheets. The type and arrangement of covalent bonds directly influence the hardness, melting point, and other physical properties of a substance. Covalent bonds are directional, which means they can support complex geometries, such as the tetrahedral shape in diamond or the planar hexagonal rings in graphite.

Physical Properties of Substances

Physical properties of substances such as hardness, melting point, and malleability are greatly affected by the type of chemical bonding they possess and the forces at play between these bonds.

The robust covalent bonds in diamond result in its high melting point and exceptional hardness. As these bonds extend throughout the entire crystal in a 3D network, it requires substantial energy to disrupt the structure, explaining the significant hardness and high melting point.

In graphite, while the covalent bonds within the sheets are strong, the interlayer van der Waals forces are much weaker. This makes the sheets easy to separate and slide past each other, leading to graphite's unique lubricating properties. The physical properties of these two forms of carbon showcase the significant impact of both intra- and intermolecular interactions.