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Allotropes of Carbon: Diamond and Graphite

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    Transport Phenomenon (Electrical and Thermal) in two allotropic forms of Carbon (Diamond and Graphite) Graphite and Diamond both are formed from carbon (two allotropic forms of carbon). Though they have similar constituent element, they differ a lot in their properties. Diamond is a good thermal conductor but a bad electrical conductor, while graphite is a bad thermal conductor but a good electrical conductor. This is one example of their property difference. The difference in their properties arises because of different arrangement of carbon atoms present in them.

    Basic physical and chemical properties of graphite and diamond to highlight their differences are as follows: Differences between Graphite and Diamond Physical Appearance: Graphite is opaque and metallic- to earthy-looking while diamonds are transparent and brilliant. Another important physical difference is their hardness. The hardness of minerals is compared using the Moh’s Hardness Scale, a relative scale numbered 1 (softest) to 10 (hardest). Graphite is very soft and has a hardness of 1 to 2 on this scale.

    Diamonds are the hardest known natural substance and have a hardness of 10. Diamond is used as an abrasive because of its great hardness, whereas graphite is used as a lubricant. Structural Differences: Diamond Space Group Fd3m face-centered cubic Atoms/unit cell 8 Cell volume 45. 385 x 10-24cm3 X-ray density 3. 5155 g/cm3 {draw:frame} Graphite Space Group C6/mmc; C-centered hexagonal Lattice parameters a = 2. 4612 A c = 6. 7079 A Atoms/unit cell 4 Cell volume 35. 189 x 10-24cm3 X-ray density 2. 2670 g/cm3 Electronic Configuration:

    Diamond The carbon atom’s electronic configuration is believed to change from its ground state in diamond as follow: {draw:frame} Quantum-mechanical calculations indicate that greater overlap between orbitals results in a stronger covalent bond. The diamond structure represents a three-dimensional network of strong covalent bonds, which explains why diamond is so hard. The diamond structure is cubic with a cube edge length of ao = 3. 567A and can be viewed as two interpenetrating FCC structures displaced by (1/4,1/4,1/4) ao. Graphite

    In going from its ground state to the graphite structure, a carbon atom’s electronic configuration is believed to change as follows: {draw:frame} In the graphite structure, overlap occur between the 2(sp2) orbitals of neighbouring atoms in the same plane. For such neighbours a side-to-side overlap also occurs between their unhybridized p orbitals. A side-to-side bonding known as ? -bonding results between these neighbors. The electrons participating in this ? -bonding seem able to move across these ? -bonds from one atom to the next.

    This feature explains graphite’s ability to conduct electricity along the sheets of carbon atom parallel to the (0001) direction. An in-plane nearest-neighbour distance is 1. 421 A. Normal to (0001), adjacent sheets of carbon atoms are held together by the weak Van der Waals bonds and separated by a distance 3. 40 A. This gives softness to the structure . Thermal Properties: Diamond: Diamond crystal is a three-dimensional network of carbon atoms. All carbon atoms in the network are strongly bonded by carbon-carbon covalent bonds.

    The diamond crystal has a highly symmetric cubic structure. The carbon atoms in diamond are precisely aligned. Thus diamond is an ideal crystal. Atoms in the crystal lattices in solids vibrate. These vibrations, called the atomic vibrations facilitate thermal conduction (transport of heat) in solids. In an ideal crystal, the lattices are so precisely aligned that they do not interact with each other. Therefore an ideal crystal conducts better than a non-ideal crystal resulting in ideal crystals having good thermal conductivity, which is a measure of heat conduction.

    Diamond being an ideal crystal is thus a good thermal conductor. Graphite The thermal properties of graphite are highly anisotropic, since phonons propagate very quickly along the tightly-bound planes, but are slower to travel from one plane to another, as there is no bonding between the layers. The thermal conductivity of a substance gives a rough indication for how rigidly components within a crystalline solid are held together and how much imperfections are incorporated within a crystal.

    Moving components within a substance tend to scatter the heat carrying electrons and phonons, in effect reduce the thermal conductivity of that material. As a result graphite is a good conductor in planar direction as the structure is perfect, however in perpendicular direction, structural imperfections, lack of bonding and slippery layers tend to scatter heat carrying electrons and phonons, resulting in graphite being a bad conductor of heat in perpendicular direction.

    Electrical Properties: Diamond: In diamonds (sp3 hybridized), each carbon atom is strongly bonded to four adjacent carbon atoms located at the apices of a tetrahedron (a three-sided pyramid). The four valence electrons of each carbon atom participate in the formation of very strong covalent bonds. These bonds have the same strength in all directions. This gives diamonds their great hardness. Since there are no free electrons to wander through the structure, diamonds are excellent insulators.

    However, generally diamonds in nature are found with certain impurity levels, mainly doped with nitrogen and boron. This renders diamond semi-conductor like properties. Graphite: Hence, Graphite can conduct electricity due to the vast electron delocalization within the carbon layers (a phenomenon called aromaticity). These valence electrons are free to move, so are able to conduct electricity. However, the electricity is only conducted within the plane of the layers.

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