In A Nutshell
Carbon has a lot of different crystal structures. Two of them—diamonds and graphite—are always switching personalities, Jekyll & Hyde–style.
The Whole Bushel
This switching is very odd, because the two minerals couldn’t be more different.
Diamond is transparent and the hardest natural substance known. Graphite, an opaque mineral, is so soft you can pull a chunk of it out of an outcrop and write with it (though pencil manufacturers mix it with clay before wrapping it up in painted wood and sticking an eraser and a big “No. 2” on it).
Diamond is an excellent electrical insulator; graphite is a terrific conductor. Diamonds are abrasives; graphite is a lubricant . . . well, you get the idea.
Graphite, the Dr. Jekyll half, has its carbon atoms layered in two-dimensional sheets. Each sheet has carbon-carbon bonds, which are very strong, but the sheets are only held together by weak stacking interactions that let them slide past each other. Diamond, a very attractive Mr. Hyde, has its carbon atoms arranged in a tetrahedron. Now those powerful carbon bonds can back one another up in three dimensions. They’re ready to handle anything that comes their way.
Structure is destiny in the world of crystals. The structures of diamond and graphite are different because of different conditions where the carbon crystallized. Diamond forms deep within the Earth. Its three-dimensional carbon crystals stand up well to the high-temperature, high-pressure conditions down there.
No one is quite sure how the carbon that turns into diamonds got into the planet’s mantle in the first place. It could have been inorganic and present since Earth formed, or it might once have been living matter that turned into graphite and was carried down via the subduction of a tectonic plate.
Let’s leave the diamonds buried in Earth’s mantle for a bit and take a closer look at that graphite.
Graphite always forms at the planet’s surface, where geologic processes carbonize the remains of living matter. Its two-dimensional shape is quite stable under the relatively low temperatures and pressures up here. If you turn up pressure and temperature, though, either by subducting graphite down into the bowels of the Earth or by an asteroid impact, graphite becomes unstable and turns into diamond.
It’s now time for the diamonds to reach the surface. That happens in a properly dramatic way for such a treasure, in an explosive kimberlite eruption.
Once up here, diamonds undergo the same fish-out-of-water experience graphite had when its formerly comfortable pressure and temperature conditions changed. They become unstable and start to break down into graphite.
Luckily for us, that process is extremely slow. It takes a lot of energy to undo diamond’s powerful carbon-carbon bonds. Sorry—you’re not going to be able to write a love letter with that diamond ring or pin for many thousands of years. Well, not unless you’re in a hurry to get some graphene, the latest “miracle substance.”
Researchers are trying to make graphene by zapping diamond with a laser. It would be insane to try that with, say, the Hope diamond, but there are plenty of relatively inexpensive natural and synthetic industrial diamonds out there to play with.
Diamond and graphite are both crystallized carbon, but the way their atoms are arranged makes a big difference in their overall properties. It’s relatively easy to turn 2-D graphite into diamond, but 3-D diamond is only mostly forever. It will turn into graphite given enough heat and pressure . . . or time.
Show Me The Proof
Scientific American: How can graphite and diamond be so different if they are both composed of pure carbon?
Smithsonian: Diamonds Unearthed
Popular Mechanics: How Does an Asteroid Impact Make Diamonds?
Kansas Geological Survey: What are Kimberlites?
Royal Society of Chemistry: Carbon
CNN: ‘Miracle material’ graphene one step closer to commercial use
Lasers for Science Facility Programme: Turning diamond to graphite