Silicon, Silicon Dioxide, and Carbon Allotropes.
| Feature | Diamond | Graphite | Graphene | C₆₀ Fullerene |
|---|---|---|---|---|
| Structure | 3D Network | Layered | Single Sheet | Sphere |
| Bonds per C | 4 | 3 | 3 | 3 |
| Conductivity | No | Yes | Yes (Very High) | Semi |
| Hardness | Very Hard | Soft (Layers slide) | Very Strong | Moderate |
Tetrahedral lattice structure identical to Diamond. Each Si bonded to 4 others. Semiconductor — used in microchips.
Giant covalent network (Sand/Quartz). Each Si bonded to 4 O, each O bonded to 2 Si. Strong, high MP, non-conductive.
Each carbon in Graphite uses only 3 of its 4 valence electrons for bonding. The 4th electron is delocalized across the entire layer — creating mobile charge carriers (like a metal).
Both diamond and SiO₂ are giant covalent structures with high melting points. But diamond is used as a cutting tool and SiO₂ (sand) is not. Why?
Diamond has each carbon bonded to 4 others in a uniform 3D tetrahedral lattice — making it the hardest known natural substance. SiO₂ also has a 3D network, but the Si–O bonds are longer and the lattice arrangement is less compact, so it is hard but not as hard as diamond. Crucially, diamond's uniform C–C bonding gives it no weak points.