Carbon

Carbon

Carbon is not commonly found in a free state. The free element is mainly found as graphite and diamond. Carbon occurs in a number of other forms, e.g. wood charcoal, animal charcoal, coke, soot (lampblack).

Carbon compounds are found in many naturally occurring substances such as coal, petroleum, wood, coal gas, natural gas, carbonates, shells, organic matter of all kinds, all living things, and occurs in the air to a small, but very important extent (0.03–0.04% by volume) as carbon dioxide. The carbohydrates, proteins and lipids in all living things contain carbon.

Pure carbon is found in the form of diamond in Tanzania (Mwadui), Sierra Leone, India, South Africa, Russia and South America; and impure carbon, as graphite, is found in Sri Lanka.

Allotropic forms of carbon

The allotropic forms of carbon are:

1. graphite.
2. diamond; and
3. amorphous carbon.

These allotropes have got different molecular structures. The structural differences are mainly due to the way their atoms are packed.

Graphite

Graphite has a layer structure. Figure 1.18 illustrates one layer of the structure of graphite. Each layer consists of carbon atoms covalently bonded together into hexagonal rings. These rings form flat parallel layers, one over the other. The force that holds the carbon atoms together are very strong. Adjacent layers are held together by weak van der Waals' forces as shown in figure 1.19. The layers readily slide over one another, accounting for the soft and greasy texture of graphite.

Carbon atoms in one graphite layer

Graphite structure

Carbon atom has got four electrons in its outer shell. Each carbon atom forms three covalent bonds to other carbon atoms. Thus, three of its four outermost electrons are paired up to form covalent bonds. The fourth electron is not attached to any particular atom (delocalized) and is free to move anywhere along the layers. Graphite conducts electricity in the plane of the layers but not at right angles to them.

Because conduction of electricity involves movement of unshared electrons from one atom to another atom, graphite is a good conductor of electricity since the hexagonal layers permit this movement. It is also a good conductor of heat for a similar reason.

Physical properties of graphite

1. It is a black, soft and slippery substance. It feels soapy and greasy. It has a metallic lustre and is opaque to light.
2. It has low relative density (2.3) as compared to diamond (3.5).
3. Graphite is a good conductor of heat and electricity due to the delocalized electrons.
4. It has a very high melting point (3730°C) and boiling point (4830°C). The melting and boiling points are high because of strong covalent bonds between the carbon atoms which require more energy (heat) to break in order to melt graphite.

Uses of graphite

1. It is used as an electrode in electrolysis and as a positive terminal in dry cells. The use of graphite as electrode in electrolysis has an advantage because it does not react readily with most substances (it is an inert electrode).
2. It is used as a lubricant, particularly when high temperatures are involved, where the usual lubricating oils easily decompose due to extreme heat. It is a lubricant for dynamos, electric motors and fast-moving parts of machinery.
3. The major use of graphite is in making lead pencils of different hardness, by mixing it with different proportions of clay. The weakly held layers of carbon atoms in graphite easily slide over each other and are left behind on paper as black marks.
4. Being resistant to chemicals and having a high melting point and also because it is a good conductor of heat, graphite is used to make crucibles.
5. Graphite has the ability to absorb fast-moving neutrons, thus, it is used in nuclear reactors to control the speed of the nuclear fission reaction.

Diamond

This basic unit is repeated in three dimensions as shown in figure 1.21 to form a giant tetrahedral structure of millions of carbon atoms, all forming four covalent bonds to each other. The melting point of diamond is high. This is because of the strong covalent bonds between carbon atoms, which require a large amount of heat energy to break up.

Physical properties of diamond

1. It is the hardest natural substance known. This due to the strong covalent bonds between the carbon atoms in diamond. Again, the compact tetrahedral arrangement of carbon atoms contributes to its hardness.
2. It has the highest melting point (3550°C) and boiling point (4289°C).
3. It has a high relative density (3.5) compared to graphite (2.3).
4. It is a poor conductor of heat and electricity. This is because there are no free electrons to conduct electricity. All electrons are firmly held in covalent bonds.
5. It is colourless, transparent and has a dazzling (amazing) brilliant lustre and radiance.
6. It has a high refractive index of 2.5. The high refractive index results in high dispersion of light, making it suitable for use in jewelers.

Uses of diamond

1. It is used in making jewellery.
2. Due to its extreme hardness, it is used to make glass cutters, drilling devices, rock borers, and as an abrasive for smoothing very hard materials.

Comparison of the properties of diamond and graphite

Diamond	Graphite
Colourless, transparent and glittering	Black, opaque with metallic lustre
Hardest natural substance known, used to cut glass and in drills	Soft to touch, greasy or soapy, can be used as a lubricant and in making lead pencils
High relative density (3.5)	Low relative density (2.3)
Non-conductor	Good conductor of heat and electricity
Burns in air least readily (at about 900°C)	Burns in air readily (at 700°C)
Have strong C–C covalent bonds arranged octahedrally to form a giant molecular crystal	Have strong C–C bonds within the hexagonal rings in the sheets but weak Van der Waals' forces in between layers
Its cleavage is difficult, and it occurs along the boundaries of the octahedral crystal unit	Cleavage easy and occurs along the sheets or the layers
Prepared from graphite at very high pressure and temperature	Prepared from coke and silica mixture at high temperature
Not attacked by potassium chlorate and nitric acid together	Attacked by these reagents

Amorphous carbon

Amorphous carbon is carbon that does not have any clear shape, form or crystalline structure. Amorphous carbon is made of tiny bits of graphite with varying amounts of other elements considered as impurities. It is formed when a materials containing carbon is burned in a limited supply of oxygen, resulting in incomplete combustion.

Amorphous carbon exists in many forms.

The following are the major forms in which it occurs:

1. Charcoal
2. Lampblack (soot) or carbon black
3. Coke

Charcoal

Charcoal is made by heating organic material (animal or plant parts) to a high temperature in the absence of air or in the presence of limited amounts of oxygen or other reagents, catalysts, or solvents. This process is called destructive distillation. There are three categories of charcoal, namely, wood charcoal, animal charcoal and sugar charcoal.

Wood charcoal

Wood charcoal is made by heating wood or other vegetable matter (for example coconut shells) in the almost complete absence of air. Wood charcoal is light, porous and is a remarkably good absorbent for liquid or gases (1 cm³ of wood charcoal will absorb 100 cm³ of ammonia gas at 0°C).

Uses of wood charcoal

1. It is used in gas masks to absorb poisonous gases in air in industrial process to recover volatile materials from waste gases.
2. Wood charcoal can be used in metal refining instead of coke.
3. Wood charcoal is a good source of energy. Thus, it is used as fuel for cooking and heating in homes.

Chemical properties of carbon

1. Carbon burns in excess oxygen to form carbon dioxide:

   C(s) + O₂(g) → CO₂(g)

   In insufficient oxygen, carbon monoxide is formed:

   2C(s) + O₂(g) → 2CO(g)

2. Carbon has got reducing properties and thus acts as a reducing agent. Carbon reduces oxides of metals below it in the electrochemical and activity series to their respective metals. This occurs on strong heating, and this reaction is used industrially for extraction of metals from their ores:

   PbO(s) + C(s) → Pb(s) + CO(g)

   Fe₂O₃(s) + 3C(s) → 2Fe(s) + 3CO(g)

   ZnO(s) + C(s) → Zn(s) + CO(g)

3. Sulphur vapours react with red hot carbon to give carbon disulphide:

   C(s) + 2S(g) → CS₂(l)

4. Carbon dioxide is reduced by red hot carbon to carbon monoxide:

   C(s) + CO₂(g) → 2CO(g)

   This reaction is used in the industrial manufacture of producer gas.