Short Answer

Diamond and graphite are both allotropes of carbon, but they have different structures, resulting in contrasting properties.

• Diamond has a 3D tetrahedral network of sp³ bonded carbon atoms, making it extremely hard, transparent, and an electrical insulator.
• Graphite has planar layers of sp² bonded carbon atoms, making it soft, slippery, a good conductor of electricity, and opaque.
• The differences arise because of bonding type and atomic arrangement, even though both are made of carbon.

Detailed Explanation :

Structure and Bonding

1. Diamond:
   • Each carbon atom forms four strong covalent bonds with neighboring atoms.
   • Creates a rigid 3D tetrahedral lattice, extending in all directions.
   • No free electrons → cannot conduct electricity.
2. Graphite:
   • Each carbon atom is bonded to three other carbons in planar hexagonal sheets.
   • Layers are held together by weak van der Waals forces, allowing layers to slide over each other.
   • One electron per carbon is free → conducts electricity along the planes.

Physical Properties

Property	Diamond	Graphite
Hardness	Hardest natural substance	Soft and slippery
Density	Higher (~3.5 g/cm³)	Lower (~2.2 g/cm³)
Appearance	Transparent, brilliant	Opaque, black/grey
Electrical Conductivity	Insulator	Conductor along layers
Thermal Conductivity	Very high	Moderate
Cleavage	No cleavage, brittle	Perfect cleavage along layers

• Diamond’s hardness makes it useful in cutting and polishing tools.
• Graphite’s softness and lubricating property make it ideal for pencils and lubricants.
• Thermal conductivity is higher in diamond due to strong covalent bonding in all directions.

Chemical Properties

1. Diamond:
   • Chemically inert at room temperature.
   • Reacts only at high temperatures or in the presence of strong oxidizers.
2. Graphite:
   • Chemically stable at room temperature.
   • Burns in air at high temperatures to produce CO₂.
   • Can act as a reducing agent in some chemical reactions.

Applications

1. Diamond:
   • Cutting, drilling, polishing tools
   • Jewelry due to brilliance and transparency
   • High-performance heat sinks in electronics
2. Graphite:
   • Lubricants and pencils
   • Electrodes in batteries and electrolysis
   • Moderator in nuclear reactors

Reasons for Property Differences

• Bonding Type: sp³ in diamond → 3D network → hard, strong bonds
• sp² in graphite → planar sheets → layers can slide, free electrons for conduction
• Electron Mobility: No free electrons in diamond → insulator; delocalized electrons in graphite → conductor
• Layering: Graphite’s layers allow flexibility and softness, unlike diamond’s rigid lattice.

Conclusion

Diamond and graphite, although composed entirely of carbon, show vastly different properties due to differences in bonding and structure. Diamond is hard, transparent, and an electrical insulator, whereas graphite is soft, slippery, opaque, and a good conductor. Understanding these differences is essential in industrial applications, materials science, and chemistry.