Q.Explain the difference in properties of diamond and graphite on the basis of their structures.

Step 1 - Diamond structure

Every carbon is $sp^3$ hybridised, covalently bonded to four other carbons in a rigid, three-dimensional tetrahedral network extending throughout the crystal - there is no plane of weakness anywhere in the structure.

Step 2 - Diamond properties from structure

Because every bond in the crystal is a strong C-C covalent bond, diamond is extremely hard (the hardest known natural material) and has a very high melting point (breaking the lattice means breaking many strong bonds). All four valence electrons of each carbon are used in localised sigma bonds, so there are no free/mobile electrons - diamond does not conduct electricity, and it is transparent.

Step 3 - Graphite structure

Each carbon is $sp^2$ hybridised, bonded to only three other carbons within a flat hexagonal sheet; the fourth valence electron of each carbon occupies a delocalised pi system that spans the whole sheet. These flat sheets are stacked and held together only by weak van der Waals forces, with a relatively large interlayer spacing.

Step 4 - Graphite properties from structure …

Compare the two structures (diamond: 3-D sp3 network, all strong bonds; graphite: 2-D sp2 sheets held by weak van der Waals forces between layers) and derive each physical property (hardness, …

• Saying diamond conducts electricity because it is a form of carbon (like graphite) - diamond has no delocalised/free electrons and does NOT conduct.
• Attributing graphite's softness to weak C-C bonds within a layer - the in-layer bonds are strong; it is only the INTER-layer forces that are weak. …