What is allotrope? Write the properties and structure of diamond and graphite.
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Allotropes are the elements which exist in two or more different forms in the same physical state. Allotropes generally differ in physical properties and may also differ in chemical activity.
Diamond has a tetrahedral geometry around each carbon atom with an sp3 hybridization. ... Diamond is hard due to its strong network covalent bonds where as graphite is slippery due to the weak interactions between the 2D layers within the structure of graphite.
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Carbon is capable of forming many allotropesdue to its valency. Well-known forms of carbon include diamond and graphite. In recent decades many more allotropes are forms of carbon have been discovered and researched including ball shapes such as buckminsterfullerene and sheets such
as graphene. Larger scale structures of carbon include nanotubes, nanobuds and nanoribbons. Other unusual forms of carbon exist at very high temperatures or extreme pressures.
Graphite conducts electricity, due to delocalization of the pi bond electrons above and below the planes of the carbon atoms. These electrons are free to move, so are able to conduct electricity. However, the electricity is only conducted along the plane of the layers. In diamond, all four outer electrons of each carbon atom are 'localised' between the atoms in covalent bonding. The movement of electrons is restricted and diamond does not conduct an electric current. In graphite, each carbon atom uses only 3 of its 4 outer energy level electrons in covalently bonding to three other carbon atoms in a plane. Each carbon atom contributes one electron to a delocalised system of electrons that is also a part of the chemical bonding. The delocalised electrons are free to move throughout the plane. For this reason, graphite conducts electricity along the planes of carbon atoms, but does not conduct in a direction at right angles to the plane.
Graphite powder is used as a dry lubricant. Although it might be thought that this industrially important property is due entirely to the loose interlamellar coupling between sheets in the structure, in fact in a vacuumenvironment (such as in technologies for use in space), graphite was found to be a very poor lubricant. This fact led to the discovery that graphite's lubricity is due to adsorbed air and water between the layers, unlike other layered dry lubricants such as molybdenum disulfide. Recent studies suggest that an effect called superlubricity can also account for this effect.
When a large number of crystallographic defects bind these planes together, graphite loses its lubrication properties and becomes what is known as pyrolytic carbon, a useful material in blood-contacting implants such as prosthetic heart valves.
Graphite is the most stable allotrope of carbon. Contrary to popular belief, high-purity graphite does not readily burn, even at elevated temperatures.[3] For this reason, it is used in nuclear reactors and for high-temperature crucibles for melting metals.[4] At very high temperatures and pressures (roughly 2000 °C and 5 GPa), it can be transformed into diamond
The term allotrope refers to one or more forms of an elementary substance.
as graphene. Larger scale structures of carbon include nanotubes, nanobuds and nanoribbons. Other unusual forms of carbon exist at very high temperatures or extreme pressures.
Graphite conducts electricity, due to delocalization of the pi bond electrons above and below the planes of the carbon atoms. These electrons are free to move, so are able to conduct electricity. However, the electricity is only conducted along the plane of the layers. In diamond, all four outer electrons of each carbon atom are 'localised' between the atoms in covalent bonding. The movement of electrons is restricted and diamond does not conduct an electric current. In graphite, each carbon atom uses only 3 of its 4 outer energy level electrons in covalently bonding to three other carbon atoms in a plane. Each carbon atom contributes one electron to a delocalised system of electrons that is also a part of the chemical bonding. The delocalised electrons are free to move throughout the plane. For this reason, graphite conducts electricity along the planes of carbon atoms, but does not conduct in a direction at right angles to the plane.
Graphite powder is used as a dry lubricant. Although it might be thought that this industrially important property is due entirely to the loose interlamellar coupling between sheets in the structure, in fact in a vacuumenvironment (such as in technologies for use in space), graphite was found to be a very poor lubricant. This fact led to the discovery that graphite's lubricity is due to adsorbed air and water between the layers, unlike other layered dry lubricants such as molybdenum disulfide. Recent studies suggest that an effect called superlubricity can also account for this effect.
When a large number of crystallographic defects bind these planes together, graphite loses its lubrication properties and becomes what is known as pyrolytic carbon, a useful material in blood-contacting implants such as prosthetic heart valves.
Graphite is the most stable allotrope of carbon. Contrary to popular belief, high-purity graphite does not readily burn, even at elevated temperatures.[3] For this reason, it is used in nuclear reactors and for high-temperature crucibles for melting metals.[4] At very high temperatures and pressures (roughly 2000 °C and 5 GPa), it can be transformed into diamond
The term allotrope refers to one or more forms of an elementary substance.
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