With respect to graphite and diamond, which of the statement(s) given below is (are) correct? (A) Graphite is harder than diamond. (B) Graphite has higher electrical conductivity than diamond. (C) Graphite has higher thermal conductivity than diamond. (D) Graphite has higher \(\mathrm{C}-\mathrm{C}\) bond order than diamond.

Understanding the nature of graphite and diamond is essential for comprehending their distinctive characteristics. Both graphite and diamond are allotropes of carbon, meaning they share the same element but have different atom arrangements and bonding. This leads to varying physical and chemical properties.

Graphite is composed of layers of carbon atoms arranged in hexagons. Between these layers are weak van der Waals forces that allow the layers to slide over one another easily, contributing to graphite's softness and greasiness. This structural form is excellent for uses such as pencil lead or lubricants.

Diamond, on the other hand, showcases carbon atoms bonded in a robust three-dimensional tetrahedral lattice. This incredibly strong structure results in the hardest known natural material, making diamond ideal for cutting tools and gemstones. Each type serves different purposes due to their distinct properties, which is why knowing the difference between graphite and diamond is valuable for various applications.

Electrical conductivity in carbon allotropes is a fascinating subject that sheds light on their unique applications. The electrical properties of graphite and diamond are starkly different due to their contrasting structures and bonding patterns.

In graphite, one of the outer electrons of each carbon atom is not used in bonding and is free to move. This electron is called a 'delocalized electron', and it can conduct electricity along the layers of graphite. The movement of these free electrons is what allows graphite to conduct electricity. This property makes graphite useful in electrodes and batteries.

Diamond lacks free electrons entirely. In diamond, each carbon atom uses all of its outer electrons to form strong covalent bonds with four other carbon atoms. Without delocalized electrons, there is no movement of charge, rendering diamond an excellent electrical insulator. This explains why diamond does not conduct electricity and is used as an insulating material in certain electronic components.

The bond order between carbon atoms within these allotropes significantly affects their properties and is key to understanding their molecular structure. Bond order refers to the number of chemical bonds between a pair of atoms.

Graphite's structure is arranged such that each carbon atom is bonded to three others in a planar hexagonal lattice with a bond order of approximately 1.33, as there is a resonance structure with some double bonds between the carbon layers. This partial double-bond character allows for easier sliding of the layers over each other.

Conversely, the structure of diamond forms with a bond order of 1, as each carbon atom is bonded to four other carbon atoms through single bonds in a tetrahedral configuration. This single-bond structure contributes to diamond's extreme hardness because the tetrahedral lattice lacks any easily breakable bonds. As a result, diamonds are significantly harder and more rigid compared to the semi-flexible layers of graphite.