Allotropic forms of Carbon

Allotropy

The existence of an element in more than one form is called “allotropy” & the different crystalline forms of that element are called its allotropic forms or allotropes or allotropic modifications.

The allotropes of a particular element possess similar chemical properties but different physical properties. The difference in properties is due to the arrangement of atoms in the lattice.

Allotropic Forms of Carbon

The allotropic forms of carbon can be classified into two groups

1. Amorphous

2. Crystalline

1. Amorphous forms:

There are various amorphous allotropes of carbon namely coal, coke, charcoal, lamp black etc.

2. Crystalline forms:

Following are the two crystalline forms of carbon.

a. Diamond

b. Graphite.

a. Diamond

Diamond is a transparent, colourless, crystalline solid. It is the hardest substance known. Diamonds easily scratch other diamonds, but this damages both diamonds. It has very high refractive index (2.45) due to which it looks highly brilliant. Its melting point is 3500^oC. It is bad conductor of electricity. It is inert at ordinary temperature & burns at 900^oC to form CO₂. Pure diamond is transparent to X-Rays.

Structure of Diamond:

In diamond each carbon atom is sp³ hybridized. Thus, all the four valence electrons of carbon are involved in sigma bonding. Each carbon atom is covalently linked with four other carbon atoms to give basic tetrahedral unit.These tetrahedral units unite with one another and produce the cubic unit cell of diamond.

Uses of Diamonds:

1. The hardness of diamonds contributes to its suitability as a gemstone. Because it maintains its polish extremely well.

2. The black diamonds (Bort and carbonado) are used for drilling and boring rocks.

3. As the hardest known naturally-occurring material, diamond can be used to polish, cut, or wear away any material, including other diamonds.

b. Graphite

Graphite is a dark grey crystalline solid which possesses dull metallic luster. It is soft and greasy to feel. Its density is about 2.2 gms/cm³ and lighter than diamond. It marks the paper. It burns to give CO₂ when strongly heated in air. It is not attacked by dilute acids or alkalis.

Structure & properties of Graphite:

In graphite carbon atom is SP² hybridized. Each carbon atom uses three hybridized orbitals to form three covalent bonds and finally form a layer structure. The remaining unhybridized p-orbital of each carbon atom merge to form a giant pi orbital above and below the plane. The layers of graphite are held together by weak vander Waal forces of attraction. The distance between the layers is 3.34^oA and C-C. On applying pressure the layers slip over one another and hence graphite is slippery

Since each carbon atoms uses three valence electron for bonding, the fourth electron of each carbon is free. These free electrons are responsible

for making graphite a good conductor.

Uses of Graphite:

1. It is used in the manufacture of electrodes.

2. It is used in making pencils.

3. It is used as lubricant in industries.

4. It is used in zinc-carbon batteries.

5. Fine flakes of graphite are used in brake shoes for heavier vehicles