At normal room temperature and pressure, graphite is (slightly) more stable than diamond. But the melting point is not a good indicator of this. Melting points are determined by the bonding structure of the solid and any potential liquid (indeed there may not be a liquid phase at some pressures). They don't necessarily tell you whether a different bonding structure is more thermodynamically stable. The thermodynamic stability of different carbon structures depends on subtle tradeoffs among several factors in the bonding of the various solids (allotropes) that can be formed with different bonding arrangements. There is no easy way to tell which is the most thermodynamically stable other than by measurement or some very high-powered calculations about the quantum mechanics of the bonding.
The different allotropes have radically different structures (tetrahedral in diamond, a ball of hexagons and pentagons in buckminsterfullerene and flat plates of hexagons in graphite). Changing one into another requires a lot of bond breaking which isn't going to happen much at room temperature. Interconversion is possible at high pressures and temperatures when carbon is dissolved in liquid rock or metal and the carbon allotrope that forms is more likely to be whatever is stable at that whatever the pressure is (diamond is more stable at higher pressure which is why the earth or industry needs very high pressures to create diamond).
Being a form of carbon, diamond oxidizes in the air if heated over 700 °C. In absence of oxygen, e.g. in a flow of high-purity argon gas, diamond can be heated up to about 1700 °C. Its surface blackens but can be recovered by re-polishing. At high pressure (~20 GPa) diamond can be heated up to 2500 °C, and a report published in 2009 suggests that diamond can withstand temperatures of 3000 °C and above at high pressure. Diamonds are carbon crystal that forms deep within the earth under high temperatures and extreme pressures. At surface air pressure (one atmosphere), diamonds are not as stable as graphite, and so the decay of diamond is thermodynamically favorable (deltaH = −2 kJ / mol). Diamonds are definitely not forever. However, owing to a very large kinetic energy barrier, diamonds are metastable; they will not decay into graphite under normal conditions.
The conversion of C(diamond) right arrow C(graphite) is thermodynamically spontaneous. However, at room temperature and pressure one does not observe diamond converting to graphite. The most reasonable explanation of these data is that the conversion of graphite to diamond is rapid. spontaneity does not imply anything about the rate of reaction. diamond and graphite are at equilibrium. diamond is thermodynamically stable.
f) Why is diamond so hard? g) Why is graphite used as a lubricant? h) Why is the melting point of ice (0C, I atm pressure) so high compared to solid H,S (-82 C, 1 atm pressure)? i) Why is it necessary to state the pressure in 0 above?
8. Given the relative costs, making diamond from graphite would appear to be an economically sound proposition. The standard free energy difference for the solid-solid phase change C(s, graphite)C(s, diamond) is Δ๔ : μdiamond-graphite-2.900 kJ mol-1 (at 25°C), The densities of these materials are ρ(diamond): 3.51 g cm-3 and ρ(graphite): 2.26 g cm-3. Estimate the pressure required to convert graphite to diamond
1) Why don't diamonds convert to graphite under ambient pressures, even though diamonds are thermodynamically less stable than graphite?