Carbon is a non-metallic element with the chemical symbol C and atomic number 6. It belongs to group 14 of the periodic table. Carbon is a unique element because it forms a vast number of compounds due to its ability to form strong covalent bonds with itself and other elements such as hydrogen, oxygen, and nitrogen.
Carbon exists in both free and combined states. In the free state, it occurs as diamond, graphite, and coal. In the combined state, it is found in compounds such as carbon dioxide (CO₂), carbonates (CaCO₃), and hydrocarbons.
Allotropy is the existence of an element in two or more different physical forms in the same physical state. The different forms are called allotropes. These allotropes differ in physical properties but have similar chemical properties.
Carbon exhibits allotropy and exists mainly in two forms:
The crystalline forms of carbon have atoms arranged in a definite geometric pattern. The main crystalline forms are:
Structure: Each carbon atom in diamond is tetrahedrally bonded to four others by strong covalent bonds, forming a three-dimensional network. This structure makes diamond the hardest known natural substance.
Properties of Diamond:
$$ \text{C (s)} + \text{O}_2 (g) \rightarrow \text{CO}_2 (g) $$
Uses of Diamond:
Conversion of Graphite to Artificial Diamond
Artificial diamonds can be made from graphite by subjecting it to very high pressure and temperature in the presence of a metal catalyst such as nickel (Ni) or iron (Fe).
Process: Graphite is heated at about 3000°C and subjected to pressures of about 60,000 atmospheres. Under these conditions, the graphite structure rearranges into the tetrahedral structure of diamond.
Equation (representation):
$$ \text{Graphite} \xrightarrow[\text{Ni/Fe}]{\text{High temp/pressure}} \text{Diamond} $$
Use: These synthetic diamonds are used industrially for cutting, grinding, and polishing purposes.
Structure: Each carbon atom in graphite is bonded to three others in flat hexagonal layers. The fourth electron is free to move between layers, allowing electrical conductivity.
Properties of Graphite:
Uses of Graphite:
Structure: Fullerene is a molecule made entirely of carbon, forming hollow spheres, tubes, or ellipsoids. The most common fullerene is C₆₀ (Buckminsterfullerene), which looks like a soccer ball with 60 carbon atoms arranged in hexagons and pentagons.
Properties of Fullerene:
Uses of Fullerene:
| Property | Diamond | Graphite |
|---|---|---|
| Structure | Each carbon atom is tetrahedrally bonded to four other carbon atoms forming a three-dimensional network. | Each carbon atom is covalently bonded to three others in flat hexagonal layers; the layers are held by weak van der Waals forces. |
| Type of Bonding | Strong covalent bonds throughout the structure. | Strong covalent bonds within layers but weak forces between layers. |
| Hardness | Extremely hard — the hardest natural substance known. | Soft and slippery; the layers can slide over one another easily. |
| Electrical Conductivity | Does not conduct electricity because it has no free electrons. | Good conductor of electricity due to delocalized electrons between layers. |
| Density | High density (about 3.5 g/cm³). | Lower density (about 2.3 g/cm³). |
| Appearance | Transparent and brilliant. | Opaque and black in colour. |
| Uses | Used in cutting and drilling tools, jewelry, and abrasives. | Used as a lubricant, electrode material, and in pencil leads. |
| Conversion | Diamond can be formed artificially by subjecting graphite to high temperature and pressure in the presence of a catalyst such as nickel or cobalt. | Graphite can be converted into diamond under high pressure and temperature conditions. |