Explain why diamond is harder than graphite. Why is graphite...

Key Concepts

Electrical Conductivity

Graphite conducts electricity due to the presence of delocalized ? electrons within its layers, which can move freely and carry charge. In diamond, all four electrons of carbon participate in forming strong, localized covalent bonds through sp³ hybridization, resulting in no free electrons for conduction, thus making diamond an electrical insulator.

Hybridization

The differences in hybridization account for the contrasting properties. In diamond, carbon atoms are sp³ hybridized, leading to a robust, three-dimensional framework. In graphite, carbon atoms are sp² hybridized, forming planar structures with one electron delocalized in a ?-bond system across the layers, which influences both its mechanical and electrical properties.

Crystal Structure

Diamond and graphite differ in their crystal structure. Diamond has a three-dimensional network of carbon atoms linked by strong covalent bonds in a tetrahedral arrangement, which gives it exceptional hardness. In contrast, graphite is composed of layers of carbon atoms arranged in a planar hexagonal lattice; the strong bonds hold atoms within each layer, but weak van der Waals forces connect the layers, making graphite much softer.

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