Carbon forms two well-known macroscopic solids—diamond and graphite. Both are composed solely of carbon atoms, yet they display dramatically different physical properties. Consider the following particle-level descriptions of their structures.
• Diamond: each carbon atom is covalently bonded to four other carbon atoms in a tetrahedral arrangement, generating a continuous three-dimensional framework. The C–C bond length in this network is approximately 154 pm.
• Graphite: each carbon atom is covalently bonded to three other carbon atoms in a trigonal planar arrangement, producing extensive two-dimensional hexagonal sheets. Adjacent sheets are separated by about 340 pm and are held together only by weak intermolecular forces.
Use these descriptions to address the questions that follow.
(a) Describe the difference in the dimensionality of the covalent bonding in diamond and graphite.
(b) Diamond is extremely hard, whereas a layer of graphite can be scratched by a fingernail. Explain this contrast in hardness in terms of the atomic-level structures provided.
(c) Electrical measurements show that diamond is an electrical insulator, whereas graphite conducts electricity parallel to the sheets but not perpendicular to them. Explain these observations using the bonding pictures above.