Properties of Graphite

Graphite is a crystalline form of carbon, typically occurring in metamorphic rocks as flakes or crystalline layers. It is formed by the metamorphosis of carbonaceous sediments. In nature, it is also found in igneous rocks, meteorites, and the mineral plumbago. Artificially, it can be prepared by heating powdered coke mixed with some sand and ferric oxide.

Graphite is a covalent solid as its constituent carbon atoms are held strongly by covalent bonds. It is the most stable form of carbon, and when subjected to high pressure and temperature, it can turn into a diamond, but that takes millions of years. The properties of graphite differ from other covalent solids because of its structure.

Overview

Although graphite is a form of pure carbon, a non-metal, it displays the properties of both a metal and a nonmetal. The properties of graphite can be attributed to its crystalline structure. 

Graphite has a two-dimensional planar structure where each carbon is sp2 hybridised. In a single layer, each individual carbon atom is bonded to three of its neighbouring carbon atoms through covalent bonds forming hexagonal planer rings. The last free valence electron of each atom is free to move between different layers. In the rings, the Carbon–Carbon covalent bond length is 141.5 pm which indicates a strong bonding. Thus, graphite has two-dimensional sheet-like polymeric rings. Each sheet or layer may be regarded as a fused system of benzene rings. Any two successive sheets are about 340 pm apart. This large distance between two successive layers does not permit the formation of covalent bonds. Successive layers are able to slide one over the other due to weak Van der Waals forces holding them.

Physical Properties of Graphite

Graphite is a greyish black, opaque substance with a metallic lustre. It marks a black stain on the paper. Although graphite is flexible, it is not elastic. Graphite’s structure, as discussed above, can be used to explain the characteristic properties of graphite as follows:

1. Slippery nature: The successive layers in graphite are bonded by weak Van der Waals forces, which allows the sliding of one layer over another. This, in turn, also allows the cleavage of the structure along the lines of planes. Consequently, black lead is soft and slippery.
2. Lustrous: Graphite contains delocalized electrons that can move freely through different layers of the graphite lattice, causing reflection of any light incident on them. As this reflection is specular and not diffused, the surface appears shiny or lustrous.
3. High melting point: The carbon atoms in each layer are bonded by strong covalent bonds, due to which the melting point of graphite is high, about 3500°C. 
4. Conductivity: Each carbon in graphite is sp2 hybridised. Thus, one valence electron of each carbon atom is free to move from one point to another. The unhybridized orbitals containing one electron each overlap laterally to form bonds in the same layer. These free electrons are delocalized and move under the influence of heat and electric fields. Thus, it makes a good thermal and electric conductor.
5. Low density: Its density is 2.26 g/cm3, which is greater than that of diamond due to the large distance between successive layers. 
6. Low coefficient of thermal expansion: Graphite has an extraordinarily low coefficient of thermal expansion. This implies that it can be heated and be exposed to extremely high temperatures without expanding all that much. It retains its form and strength at high temperatures.
7. Graphite displays low adsorption of X-rays and neutrons. 
8. The chemical bonds in the black lead are stronger than those that make up diamonds.

Chemical Properties of Graphite

The following are the chemical properties of graphite:

1. At standard temperature and pressure, graphite is generally an inert material. 
2. It does not rapidly react with water.
3. Graphite is also highly refractory – resistant to chemical attack, heat, and pressure. 
4. When graphite is exposed to very high temperatures and strong oxidising atmospheres, it undergoes oxidation and decomposes. It burns in air at 700°C to form carbon dioxide.

C + ½ O2 → CO

C + O2 → CO2

1. In bulk form, it is non-flammable but combustible.
2. It acts as a reducing agent. It has violent reactions with fluorine, chlorine dioxide, and potassium peroxide, which are very strong oxidising agents.

Uses of Graphite

The properties of graphite make it a unique material that has extensive use in the following fields :-

• Owing to graphite’s soft and slippery nature, it is used as a grease, a type of lubricant.
• Due to its high electrical conductivity, graphite is reflected in the making of electrodes for electric furnaces.
• The high melting point of graphite makes it suitable for making crucibles for melting metals. It is also used in fireproofing systems in the form of graphite plates.
• Graphite is highlighted as an excellent material for heat sinks, carbon fibre, and heat exchangers due to its thermal conductivity and low coefficient of thermal expansion. 
• Graphite leaves a black mark on paper because of which it is extensively used for making pencil leads.

• Graphite’s mechanical properties at high temperatures and chemically inert nature make it very valuable in nuclear applications. It is used in nuclear reactors as moderators to slow down the speed of neutrons.

Conclusion

The uniqueness of graphite can be attributed to its crystalline structure. Graphite, or black lead, is a greyish black opaque covalent crystalline solid made purely of carbon. Although it is non-metal, graphite has a metallic lustre and is slippery. It has a low density, high melting point, and it is a coefficient of thermal expansion. Unlike other non-metals, it has high thermal and electrical conductivity. Chemically, graphite is inert, resistant to heat, pressure, and does not react with water or air at standard temperature. However, at very high temperatures and strong oxidising atmospheres, graphite undergoes oxidation to form carbon dioxide. In bulk form, it is non-flammable but combustible. It acts as a reducing agent, reacting violently with strong oxidising agents.

Due to its distinct properties, graphite is used as a lubricant in the making of pencil leads, electrodes for electrical furnaces, high thermal applications, in the steel and glass industry, and in nuclear reactors.