Question

Explain the difference in properties of diamond and graphite on the basis of their structures.

Short answer

Diamond is hard with tetrahedral bonds, a poor conductor; graphite is soft with planar bonds, a good conductor.

Step-by-step solution

Identify the Structures

Diamond and graphite are both allotropes of carbon, meaning they are made from the same element but have different structures. Diamond has a tetrahedral lattice structure where each carbon atom is covalently bonded to four other carbon atoms forming a 3-dimensional network. Graphite, on the other hand, consists of layers of carbon atoms bonded in a hexagonal arrangement, with weak Van der Waals forces holding the layers together.

Analyze the Bonding in Diamond

In diamond, each carbon atom forms four strong covalent bonds with its neighboring carbon atoms. This tetrahedral bonding pattern creates a very rigid, three-dimensional structure. The strength and rigidity of this bonding contribute to diamond’s properties such as being extremely hard, having a high melting point, and being an excellent electrical insulator.

Explore the Bonding in Graphite

In graphite, each carbon atom forms three covalent bonds with other carbon atoms in a plane, creating layers that have a hexagonal pattern. These planes are held together by weak Van der Waals forces. The presence of these weaker forces between layers allows the layers to slide over each other, making graphite soft and slippery. Furthermore, this layered structure results in graphite being a good conductor of electricity due to the mobility of electrons along the layers.

Compare Electrical Conductivity

Diamond does not conduct electricity as all the electrons are involved in covalent bonding, leaving no free electrons to carry charge. Graphite conducts electricity because only three out of four outer electrons of each carbon atom are used for bonding, leaving free moving electrons (also known as delocalized electrons) that can carry an electric charge.

Compare Physical Properties

Because of its strong covalent bonding and rigid structure, diamond is the hardest natural substance known, with a high melting point. In contrast, graphite is soft and can be used as a lubricant or in pencils, due to the layers of carbon sliding over each other easily. These differences in physical properties are a direct result of their molecular structures.

Key concepts

Diamond structure

The diamond structure is a fascinating example of carbon's ability to form strong, three-dimensional networks. Each carbon atom in a diamond is surrounded by four other carbon atoms. They are connected in a way that forms a tetrahedral lattice. This structure means each carbon atom is covalently bonded to its neighbors.

This arrangement is responsible for some of diamond's most remarkable properties. Because these covalent bonds are so strong and numerous, diamonds are incredibly hard and have a high melting point. When we think about the toughness of a diamond, it's because of this robust internal framework.

Moreover, the diamond's structure influences its optical properties; it is renowned for its clarity and brilliance. However, one significant trait affected by this arrangement is diamond's electrical conductivity, or rather, lack of it. All the electrons in diamond are used in covalent bonding. As a result, there are no free electrons available to conduct electricity, making diamond an excellent insulator.

Graphite structure

Graphite has a unique layered structure, giving it a different set of properties compared to diamond. In graphite, carbon atoms are bonded together in flat, two-dimensional layers. Within each layer, the atoms form a hexagonal pattern and are bonded with strong covalent bonds.

These layers stack on top of each other with only weak Van der Waals forces in between. This weak interlayer interaction allows the layers to slide over each other easily, resulting in graphite being soft and slippery. This property makes graphite an ideal material for pencils and lubricants.

Another key aspect of this structure is graphite's electrical conductivity. Unlike diamond, not all electrons in graphite are used for bonds between atoms. Each carbon atom uses only three of its four available electrons for bonding with adjacent atoms, leaving one free electron per carbon atom. These free, delocalized electrons can move across the layers, enabling graphite to conduct electricity efficiently.

Covalent bonding in carbon

Covalent bonding is a defining feature of carbon's ability to form diverse structures like diamond and graphite. In covalent bonding, atoms share electrons to achieve a stable electronic configuration.

In diamond, each carbon atom engages in four covalent bonds, creating a strong, interconnected 3D lattice. This extensive bonding results in high hardness and insulating properties.

Conversely, in graphite, each carbon atom forms three covalent bonds within a plane, leaving one electron available to contribute to electrical conductivity. This bonding configuration leads to the formation of the layered structure that is key to graphite's properties.

Understanding covalent bonding in carbon helps explain why allotropes of carbon like diamond and graphite exhibit such contrasting qualities, ranging from hardness and insulation in diamond to softness and conductivity in graphite.