Why is graphite more stable than diamond at 25˚C and 1 atm?
Graphite is more stable than diamond at 25˚C and 1 atm because of the difference in their crystal structures. To understand why this is the case, we need to consider the concept of free energy.
The stability of a substance is determined by its free energy, which is a measure of the energy available to do work. At a given temperature and pressure, a substance will have a certain free energy associated with it. The substance with the lower free energy is considered to be more stable.
Now, let's talk about the crystal structures of graphite and diamond. Graphite consists of layers of carbon atoms arranged in a two-dimensional hexagonal lattice. These layers are held together by weak van der Waals forces, allowing them to slide over each other easily. On the other hand, diamond has a three-dimensional structure where each carbon atom is bonded to four neighboring carbon atoms, forming a rigid tetrahedral network.
At ambient conditions (25˚C and 1 atm), the energy required to break the strong covalent bonds in diamond is higher compared to the weak van der Waals forces in graphite. This means that to convert diamond into graphite, one needs to supply energy to break the strong bonds and allow the atoms to rearrange into a more stable configuration. Therefore, graphite is more thermodynamically stable than diamond at these conditions.
To quantitatively compare the stabilities, one can calculate the free energies of graphite and diamond using thermodynamic equations and experimental data for their respective enthalpies and entropies of formation. These values can be found in thermodynamics databases or textbooks. By comparing the free energies, it can be determined which form of carbon is more stable at a given temperature and pressure.
In summary, graphite is more stable than diamond at 25˚C and 1 atm due to the lower energy required to break the weak van der Waals forces in graphite compared to the strong covalent bonds in diamond.