Diamond is much harder than graphite mainly because diamond has a three dimensional covalent network structure, whereas graphite has weakly bonded layers. Which explanation best describes this difference in hardness?

Introduction / Context:
This question compares two allotropes of carbon, diamond and graphite, and asks why their hardness is so different. Both substances are made only of carbon atoms, yet diamond is one of the hardest known materials, while graphite is soft and used as a lubricant and pencil lead. Understanding the structural reason for this difference is a classic topic in solid state chemistry and materials science.

Given Data / Assumptions:

• Both diamond and graphite are allotropes of carbon.
• Diamond is experimentally found to be much harder than graphite.
• We are given several possible structural explanations and must choose the best one.

Concept / Approach:
The key concept is the arrangement and bonding of carbon atoms. In diamond, each carbon atom is sp3 hybridized and bonded covalently to four other carbon atoms in a tetrahedral three dimensional network. This creates a very rigid, strong structure throughout the crystal. In graphite, each carbon atom is sp2 hybridized and bonded to three other carbons in planar hexagonal layers. These layers are held together by weak forces, so they can slide over each other, making graphite soft. The correct explanation must mention the rigid three dimensional covalent network in diamond versus layered structure in graphite.

Step-by-Step Solution:

Verification / Alternative check:
Solid state chemistry diagrams and models of diamond show a repeating tetrahedral framework, while graphite is represented as parallel sheets of hexagons. Measurements of mechanical properties show that diamond has extremely high hardness and high bulk modulus. Graphite, in contrast, is soft enough to leave marks on paper. These observations are consistent with strong three dimensional bonding in diamond and weak interlayer forces in graphite. Textbooks repeatedly emphasize this structural difference when explaining the contrasting properties of carbon allotropes.

Why Other Options Are Wrong:

• Option B: Neither diamond nor graphite contains metal atoms; both are pure carbon.
• Option C: Density does not determine hardness by itself; many dense materials are not extremely hard and some less dense materials can be quite hard.
• Option D: Both diamond and graphite involve covalent bonding; diamond does not consist of purely ionic bonds.
• Option E: The number of atoms in the crystal is not the reason for hardness; it is the bonding pattern that matters.

Common Pitfalls:
Students sometimes believe that hardness is simply a result of density or the number of atoms, which is not correct. Others may think that diamond must contain some special element, but it is purely carbon, just like graphite. The key lesson is that the same element can form very different structures, called allotropes, and these structural differences lead to dramatically different physical properties. Focusing on bonding and geometry rather than just composition helps avoid these misunderstandings.

Final Answer:
Diamond is harder than graphite because Diamond has a rigid three dimensional covalent network structure, while graphite consists of weakly bonded planar layers.