Carbon is one of the most versatile elements in the periodic table. One reason for this versatility is its ability to form allotropes—different structural forms of the same element. In IB Chemistry, allotropy appears in Topic 4 (Bonding) and Topic 3 (Periodic Trends). Understanding carbon allotropes helps explain differences in physical properties, bonding, conductivity, and real-world applications.

This article introduces two of the most common carbon allotropes and explains why they behave so differently.

What Are Allotropes?

Allotropes are different structural forms of the same element in the same physical state.

The atoms are the same, but the bonding, arrangement, and structure differ.
This causes major changes in:

• Hardness
• Electrical conductivity
• Melting point
• Density
• Chemical reactivity

Carbon is famous for its many allotropes, but the most commonly studied in IB Chemistry are diamond and graphite.

Allotrope 1: Diamond

Structure

Diamond has a giant covalent lattice. Each carbon atom forms four single covalent bonds (sp³ hybridization) with four other carbon atoms in a tetrahedral arrangement. This creates a rigid, three-dimensional network.

Properties

1. Extremely Hard

The continuous lattice of strong covalent bonds makes diamond the hardest known natural material.

2. Very High Melting Point

Huge amounts of energy are required to break the covalent network.

3. Electrical Insulator

Diamond does not conduct electricity because all four valence electrons of each carbon atom are involved in bonding—no delocalized electrons are present.

4. Transparent and Lustrous

The symmetry and strong bonding give diamond its famous sparkle.

Uses

• Cutting tools
• Drill bits
• Jewelry
• High-pressure scientific instruments

Allotrope 2: Graphite

Structure

Graphite is also a giant covalent structure, but with a crucial difference: each carbon atom forms three covalent bonds (sp² hybridization) to create hexagonal layers. The fourth electron becomes delocalized and moves freely across the sheet.

Properties

1. Conducts Electricity

The delocalized electrons allow graphite to conduct electricity, making it unusual for a non-metal.

2. Soft and Slippery

Layers are held together by weak London dispersion forces. They slide easily, which explains why graphite is used as a lubricant.

3. High Melting Point

Even though layers slide, the covalent bonds within each sheet are strong and require a lot of energy to break.

4. Opaque and Black

The delocalized electrons absorb light, giving graphite its dark appearance.

Uses

• Pencils ("lead")
• Lubricants
• Electrodes
• Batteries
• Conductive materials

Why These Allotropes Are So Different

Diamond and graphite differ because of:

Bonding and Hybridization

• Diamond: sp³ → 4 bonds → 3D lattice
• Graphite: sp² → 3 bonds + delocalized electron → layered structure

Electron Delocalization

• Diamond: none → electrical insulator
• Graphite: delocalized electrons → good conductor

Forces Between Layers

• Diamond: no layers → rigid and hard
• Graphite: layers slide → soft and flaky

Even though both are pure carbon, different bonding patterns create drastically different properties.

Other Carbon Allotropes (IB Extension Knowledge)

Although the question only requires two, it’s helpful to know others:

• Fullerenes (C₆₀) – spherical molecules
• Graphene – single layer of graphite
• Carbon nanotubes – rolled sheets of graphene
• Amorphous carbon – soot, charcoal

These may appear in HL questions or data-based assessments.

FAQs

Why does graphite conduct electricity but diamond does not?

Graphite has delocalized electrons that move through its layers. Diamond has no free electrons because all valence electrons are tied in covalent bonds.

Are diamond and graphite both giant covalent structures?

Yes. Both are giant covalent lattices, but with different bonding patterns, leading to different physical properties.

Which allotrope is more stable?

Graphite is the thermodynamically more stable allotrope under standard conditions, even though diamond is kinetically stable and does not convert spontaneously.

Conclusion

Two common allotropes of carbon are diamond and graphite. Diamond forms a strong, three-dimensional tetrahedral network, making it extremely hard and unable to conduct electricity. Graphite forms layered sheets with delocalized electrons, giving it lubricating properties and electrical conductivity. These differences illustrate how powerful variations in bonding and structure can be, even within a single element.