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Why are diamond and graphite so different even though both contain only carbon?

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A labeled diagram of the periodic table entry for carbon, showing its atomic number 6, symbol C, electron configuration, atomic weight range, solid physical state, hexagonal crystal structure, and weakly acidic oxide properties. Color coding and icons indicate it is a nonmetal, solid at room temperature, with hexagonal structure and weak acidity. — carbon Carbon and its properties.

Diamond and graphite are both composed entirely of carbon, yet they exhibit vastly different physical properties because they are allotropes—different crystalline forms of the same element characterized by distinct atomic arrangements.12

Diamond Structure

In diamond, each carbon atom is covalently bonded to four others in a rigid, three-dimensional tetrahedral lattice. This dense, interconnected structure requires the rupture of many strong bonds to break, resulting in extreme hardness, high thermal stability, and electrical insulating properties.1

Graphite Structure

Graphite consists of parallel sheets of carbon atoms arranged in fused hexagonal rings. Within these layers, the atoms are strongly bonded, but the layers themselves are held together only by weak van der Waals forces. This structure allows the layers to slide past one another, which accounts for graphite’s softness and lubricating qualities, while the arrangement also facilitates electrical conductivity.1

While graphite is the more stable form under ordinary conditions, the two materials remain distinct due to these fundamental differences in their crystal geometry.12