Diamond is known for its extreme hardness, which derives from its unique atomic structure: a dense, three-dimensional crystalline lattice where each carbon atom is covalently bonded to four others. The question of whether a diamond can be melted is complex, as this transition does not occur under everyday conditions. True melting, the conversion from a solid to a liquid state, is possible, but it requires a specific combination of immense heat and pressure.
The Default Outcome Turning to Graphite
When a diamond is heated under normal atmospheric pressure, it undergoes a phase change rather than melting. Diamond is metastable at standard pressure and temperature; the thermodynamically stable form of carbon is graphite. Heating provides the energy needed to overcome the activation barrier preventing this change. If oxygen is present, the diamond will burn, converting to carbon dioxide gas between \(690^\circ\text{C}\) and \(840^\circ\text{C}\).
If oxygen is excluded, the diamond transforms into graphite, a process called graphitization, which begins around \(1500^\circ\text{C}\) to \(1900^\circ\text{C}\). In an inert atmosphere, the carbon atoms rearrange from the four-bonded tetrahedral structure of diamond to the three-bonded, layered hexagonal structure of graphite. Since diamond transforms into this different solid phase before reaching the temperature required for true melting, the liquid state is bypassed entirely under standard pressure conditions.
Defining Diamond’s True Melting Point
To achieve true melting, the solid must transition directly into a liquid carbon phase without first converting to graphite or subliming. The theoretical melting point for diamond is estimated to be around \(4000^\circ\text{C}\) to \(4500^\circ\text{C}\). This temperature is far higher than the graphitization temperature.
However, high temperature alone is insufficient because the carbon would turn into graphite. To maintain the diamond structure long enough to melt it, the material must be subjected to enormous pressure, which stabilizes the denser diamond structure. The necessary pressure is approximately \(10\text{ GPa}\), about \(100,000\) times the atmospheric pressure at sea level.
Under these combined conditions, carbon atoms gain enough thermal energy to break the strong covalent bonds, but the high pressure forces them to remain in close proximity, resulting in a liquid carbon phase. This phase is highly transient and challenging to observe. Some experiments suggest that at ultra-high pressures exceeding \(1\text{ TPa}\) and temperatures above \(50,000^\circ\text{C}\), the liquid carbon may exhibit metallic properties.
The Carbon Phase Diagram Explained
The behavior of carbon under extreme heat and pressure is best understood by examining its phase diagram, which maps the stability regions of its different forms. The diagram shows how temperature and pressure determine whether carbon exists as diamond, graphite, a liquid, or a gas. At low pressure, the graphite phase is thermodynamically stable, while diamond is a metastable solid.
The diamond region only becomes the stable phase at very high pressures, confirming why pressure is necessary to prevent graphitization. The point where the three phases—diamond, graphite, and liquid carbon—can coexist in equilibrium is known as the triple point. Estimates place this triple point at temperatures between \(4500\text{ K}\) and \(5000\text{ K}\) (about \(4227^\circ\text{C}\) to \(4727^\circ\text{C}\)) and pressures near \(12\text{ GPa}\).
Any combination of heat and pressure below this triple point results in the conversion of diamond to graphite. Pushing the system into the liquid region requires surpassing both the temperature and pressure thresholds of this triple point. High pressure compresses the carbon system, making the dense liquid phase thermodynamically accessible before conversion to vapor or graphite.
Practical Implications of Extreme Heat
The conditions required to melt a diamond are largely a theoretical and laboratory phenomenon. Specialized equipment, such as a diamond anvil cell or shock-wave generators, is necessary to achieve the immense pressures and temperatures involved. These experiments help scientists understand the behavior of materials in extreme environments, such as planetary interiors or during hypervelocity impacts.
The industrial relevance of these extreme conditions is tied to the synthesis of diamonds. Manufacturing synthetic diamonds involves placing a carbon source under high pressure and temperature to push it into the diamond-stable region of the phase diagram, mimicking geological processes.
For the general public, the melting point of diamond is not a practical concern. A common house fire rarely exceeds \(1200^\circ\text{C}\), which is high enough to cause a diamond to graphitize or burn if exposed to air, but not high enough to approach its true melting temperature.