Final answer:
Diamond and graphite, both made of carbon, differ due to their atomic structures. Dense three-dimensional covalent bonds give diamonds their hardness, while graphite's layers bonded by weaker forces allow it to be soft and slippery.
Step-by-step explanation:
Diamond and graphite are both forms of pure carbon, yet they have distinctly different properties due to their unique atomic structures. Diamond is renowned for its extreme hardness, which is a result of a three-dimensional network of strong covalent bonds connecting each carbon atom to four others in a tetrahedral structure. This arrangement allows diamond to scratch almost any other material, including glass.
On the other hand, graphite is composed of layers of carbon atoms arranged in a hexagonal lattice. Each carbon atom is bonded to three others in the same plane, forming sheets. Between these sheets are weaker van der Waals forces, allowing the layers to slide over each other easily. This is why graphite feels slippery and is used as 'lead' in pencils; as it writes, layers of carbon are transferred to the paper.
The differences in atomic arrangement between diamond and graphite illustrate why the same element, carbon, can exhibit such contrasting characteristics. These forms of carbon, known as allotropes, showcase the vast potential of element arrangements affecting physical properties like hardness, transparency, and electrical conductivity.