“A diamond is the hardest natural substance on earth, but if it is placed in an oven and the temperature is raised to about 763 degrees Celsius (1405 degrees Fahrenheit), it will simply vanish, without even ash remaining. Only a little carbon dioxide will have been released.”
So do they melt at 3550 degrees, or vanish at 763?
I suspect that the 763C thing is diamond combustion, (even if there may not be a visible flame or ash residue.) And the carbon dioxide released will have to be several times the weight of the diamond, though that may not ‘look’ like much.
The melting point, then, would only be reached if you could heat diamond that hot in the absence of oxygen it could burn in.
Another question is whether the pure carbon would actually be a melted ‘liquid’ at 3551C, or if it would have sublimated into a gas. (Sublimation being used here for the direct state change of a solid into gas or gas into solid without an intermediate liquid state.) This would obviously depend on the pressure involved.
I took a geology class a couple of years ago and the prof mentioned that a diamond melts at 3950K (as in Kelvin) not celsius as from your link. Plus this was most likely done at a higher pressure then just 1atm on preview I see chrisk beat me to it
OP said 3550 celcius, which would be 3823 K… not really too far from your value of 3950. Reminds me of the old joke in astronomy/astrophysics, where an old hand mentions a star’s core temperature as being twelve million degrees or something like that, and a newbie asks if that’s in K or C. The response comes back ‘either!’ - a difference of 273 doesn’t really matter on that kind of scale.
According to this source(1), if you put diamond in an oven, it’ll oxidize (burn) at roughly 1300 degrees Farenheit, at atmospheric oxygen concentrations (if surface oxygen coverage is directly proportional to partial pressure). In this case, it means that:
1: The Oxidation of Diamond, J.Y. Howe, L.E. Jones, and A.N. Cormack.
According to this phase diagram for carbon (scroll down), at one bar of pressure (approximately atmospheric pressure), the carbon would have a liquid phase. It also indicates that the melting point of graphite is 4200K, and that of diamond is 4500K.
These figures differ from those quoted in the OP. I would note that the phase diagram in the OP seems to agree more with the figures in my link. In addition, I see that the OP’s link indicates that graphite “sublimes” [sublimates], which does not agree with either phase diagram.
(The phase diagram also indicates that diamond is not thermodynamically stable at room temperature and pressure; however, the diagram does not show that it takes a very long time to revert back to graphite. For this reason, diamond is referred to as “metastable.”)
P.S. For what it’s worth, I suspect that it is difficult to experimentally measure such high melting points. I’ve worked with graphite furnace atomic absorption spectroscopy, and the graphite is used as a platform for the vaporization of various metals present in small concentrations in water. The graphite heats up to several thousand deg C in this instrument, but is nowhere near its melting point. In fact, graphite is used because it has such an extremely high melting point.
As chrisk notes, you would require greater than ambient pressure to see carbon in liquid state in an oxygen-nitrogen atmosphere. (In a vacuum or an inert, noble gas atmosphere you could probably turn it liquid at 1 bar of pressure, but carbon is otherwise very reactive.) Since diamond isn’t a common engineering material, I can’t verify the temperatures ~1000K sounds about right for sublimation or burning of diamond. (Kelvins = degrees Celsius + 273.15)
Producing diamonds artificially requires very intense pressures, on the order of hundreds of bars of pressure, and temperatures between 3000-5000K. So it’s correct, if incomplete, to say that diamond “melts” (i.e. becomes liquid) at 3550K and some corresponding pressure on the PV table. The big trick of manufacturing artificial diamonds–aside from generating the requisite temperatures and pressures–is preventing it from undergoing transform to a different and more stable allotropic state like graphite or amorphous form carbon while it cools back to ambient pressure. There are forms that have even more remarkable mechanical and thermal properties than diamond (carbon nanotubes, fullerite C60, aggregated nanorods) but they aren’t found in quantity in nature.
I thought Kimberlite was a great source of diamonds because during the Kimberlite formation, diamonds (and other geologic goodies) came out of the melt–implying that carbon was mixed in the magma in a liquid form. Any geologists want to correct my way-out-of-date memory?
Nope - diamonds are xenocrysts - deep crust material carried up by the kimberlite from the diamond forming zone. I think that the reason kimberlite is a good source of diamonds is more to do with the eruption speed and temperature - too long in the hot stuff and the sparkler is a whiff of CO2 (although there may not be too much free O2 in kimberlite - it looks pretty reducing).
