Allotropes of Carbon: Diamond and Graphite

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Both Diamond and Graphite are different forms of carbon, known as allotropic forms. Despite sharing the same element, they display contrasting properties. Diamond is exceptional in conducting heat but lacks conductivity in electricity, whereas graphite struggles with thermal conduction but excels in electrical conduction. This variation in properties is a result of the distinct arrangement of carbon atoms within each form.

Graphite and diamond have distinct differences in their physical and chemical properties. Notably, their physical appearance sets them apart. Graphite is opaque and metallic in nature, whereas diamonds are transparent and radiant. Additionally, a significant disparity arises in terms of hardness. Minerals are evaluated on the Moh’s Hardness Scale that ranges from 1 (being the softest) to 10 (being the hardest). On this scale, graphite falls under the classification of very soft with a hardness rating of 1 to 2.

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Both diamond and graphite are forms of carbon, but they have distinct structures and properties. Diamond is the hardest natural substance known, with a hardness rating of 10. Due to its extreme hardness, diamond is commonly used as an abrasive material. Conversely, graphite serves as a lubricant.

In terms of structure, diamond has a face-centered cubic lattice structure consisting of 8 atoms per unit cell. Its cell volume measures 45.385 x 10-24cm^3 and its X-ray density is 3.5155 g/cm^3.

On the other hand, graphite features a C-centered hexagonal lattice structure comprising 4 atoms per unit cell. It has a cell volume of 35.189 x 10-24cm^3 and an X-ray density of 2.2670 g/cm^3.

Despite their structural differences, both diamond and graphite are important forms of carbon with diverse uses and unique properties.

The electronic configuration of the carbon atom in diamond is believed to change from its ground state as follows: {draw:frame}. Quantum-mechanical calculations suggest that a stronger covalent bond is formed through increased overlap between orbitals. The diamond structure consists of a three-dimensional network of strong covalent bonds, which accounts for its hardness. The diamond structure is cubic, with an edge length of ao = 3.567A, and can be visualized as two interpenetrating FCC structures displaced by (1/4,1/4,1/4) ao. Graphite.

The electronic configuration of a carbon atom changes from its ground state to the graphite structure. In the graphite structure, the 2(sp2) orbitals of neighboring atoms in the same plane overlap. There is also a side-to-side overlap between the unhybridized p orbitals of these neighboring atoms. This side-to-side bonding, known as ? -bonding, occurs between these neighbors. The electrons involved in this ? -bonding are capable of moving across these ? -bonds from one atom to another.

This feature describes how graphite is able to conduct electricity along the sheets of carbon atoms that are parallel to the (0001) direction. The distance between nearest neighbors within a single sheet is 1.421 A. Perpendicular to the (0001) direction, the sheets of carbon atoms are held together by weak Van der Waals bonds and are separated by a distance of 3.40 A. This results in a soft structure. When it comes to thermal properties, diamond crystals consist of a three-dimensional network of carbon atoms that are strongly bonded by carbon-carbon covalent bonds.

The diamond crystal is characterized by its highly symmetric cubic structure and precise alignment of carbon atoms. This makes diamond an exemplary crystal, as the atomic vibrations within crystal lattices facilitate thermal conduction in solids. Unlike non-ideal crystals, the lattices in an ideal crystal do not interact with each other, allowing for better conduction. As a result, ideal crystals demonstrate good thermal conductivity, which quantifies their ability to conduct heat.

Diamond is an excellent thermal conductor because it is an ideal crystal. However, graphite has highly anisotropic thermal properties due to the quick propagation of phonons along tightly-bound planes but slower travel between planes without bonding. The thermal conductivity of a material provides a rough assessment of the rigidity of components and the presence of imperfections within a crystal structure.

In a substance, when components move, they scatter heat carrying electrons and phonons, which decreases thermal conductivity. Therefore, while graphite is an excellent conductor in the planar direction due to its perfect structure, it becomes a poor conductor in the perpendicular direction because of structural imperfections, inadequate bonding, and slippery layers.

Electrical Properties: Diamond

In diamonds (sp3 hybridized), every carbon atom is tightly connected to four neighboring carbon atoms at the apices of a tetrahedron (a three-sided pyramid). The four valence electrons of each carbon atom contribute to the creation of extremely strong covalent bonds. These bonds possess equal strength in all directions, resulting in the high hardness of diamonds. Due to the absence of unoccupied electrons within the structure, diamonds display excellent insulation properties.

Generally, diamonds found in nature contain impurities like nitrogen and boron, giving them semi-conductor properties. On the other hand, graphite can conduct electricity as a result of electron delocalization within carbon layers, known as aromaticity. The valence electrons in graphite are mobile, enabling electricity conduction. However, this conduction is limited to the plane of the graphite layers.

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Allotropes of Carbon: Diamond and Graphite. (2018, Feb 20). Retrieved from

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