Which statement best describes the difference in bonding between carbon atoms in diamond and graphite?

• A. In diamond, carbon atoms form double bonds, while in graphite, they form single bonds.
• B. In diamond, each carbon atom forms four single covalent bonds in a tetrahedral arrangement, while in graphite, each carbon atom forms three covalent bonds in a planar hexagonal arrangement with delocalized electrons.
• C. In both diamond and graphite, each carbon atom forms four single covalent bonds in a tetrahedral arrangement.
• D. In diamond, carbon atoms are connected by ionic bonds, while in graphite, they are connected by metallic bonds.

Answer: (B)

Verified step by step guidance

1. Identify the type of bonding in diamond: Each carbon atom in diamond forms four single covalent bonds arranged tetrahedrally, creating a strong three-dimensional network.
2. Identify the type of bonding in graphite: Each carbon atom in graphite forms three covalent bonds in a planar hexagonal arrangement, with the fourth electron delocalized over the plane, allowing electrical conductivity.
3. Understand the difference in bonding geometry: Diamond's tetrahedral bonding leads to a rigid, strong structure, while graphite's planar bonding with delocalized electrons results in layers that can slide over each other.
4. Recognize that diamond does not have double bonds between carbon atoms, nor does it have ionic bonds; similarly, graphite does not have metallic bonds but rather covalent bonds with delocalized electrons.
5. Conclude that the key difference is the number and arrangement of covalent bonds per carbon atom and the presence of delocalized electrons in graphite, which explains their distinct physical properties.