Chemistry Regents Practice Test

At standard temperature and pressure (STP), diamond and graphite are both solid forms of carbon. What distinguishes their properties?

1. Their atomic weights

2. Different crystal structures and different properties

3. Their electrical conductivity

4. Their isotopic compositions

The correct answer is: Different crystal structures and different properties

Diamond and graphite are both allotropes of carbon, meaning they are different structural forms of the same element. The correct choice highlights that the distinction lies in their crystal structures and the resultant properties. In diamond, each carbon atom is covalently bonded to four other carbon atoms in a three-dimensional tetrahedral arrangement. This structure creates a very strong and rigid lattice, resulting in diamond being one of the hardest known materials. Additionally, diamond has a high melting point and is typically an electrical insulator due to the lack of free-moving charged particles. Conversely, in graphite, carbon atoms are bonded in layers of hexagonal lattices, where each carbon atom is bonded to three others. These layers can slide over each other, making graphite slippery and giving it the property of lubricity. Furthermore, the layers are held together by weaker van der Waals forces, which results in lower hardness compared to diamond. Importantly, graphite has delocalized electrons, allowing it to conduct electricity. The other options relate to characteristics that do not fundamentally differentiate diamond and graphite in significant ways. Their atomic weights are identical since they are both composed of carbon. Despite differences in electrical conductivity, the core distinction lies in their unique structural arrangements that lead to their varying properties