When Superman wants to super impress Lois Lane, he takes a lump of coal and squeezes it in his super fist until it becomes a diamond. Which is super.
Unfortunately, it’s not a scientifically accurate analogy for the creation of diamonds in nature. So when journalist Stephen Ornes’ 6-year-old son, Sam, asks how coal, which is black, can turn into diamonds, which are clear, there are actually a couple of issues we have to address. First, we need to know where diamonds actually come from. Then, even though diamonds aren’t coal, you’re still left with the basic question Sam is trying to get at—why can pure carbon be black under some circumstances and clear under others? Turns out, the answer has a lot to do with why life, itself, is based on carbon.
Coal is the compressed remains of ancient plants, dinosaur swamps sitting in the palm of your hand. But there are diamonds that are older than terrestrial plants. That fact alone should tell you that diamonds are not actually made from compressed coal. Instead, diamonds are probably formed deep in the Earth—much further down than the levels at which we find coal—where heat and pressure fuse atoms of carbon together into crystalline structures. Later, those crystals get vomited up from the depths with the help of volcanic vents. (You can read more about where diamonds really come from in a post I wrote back in 2012.)
It’s important to make the distinction between diamonds and coal because, if you don’t, then Sam’s question earns a misleadingly simple answer. Diamonds and coal are different colors because coal isn’t pure carbon. The stuff is loaded with impurities: Hydrogen, sulfur, mercury, and more. There’s a reason you don’t want to live next door to a coal-fired power plant and that reason is all the nasty stuff that gets released when the carbon in coal burns.
But that doesn’t mean pure carbon always looks like diamonds. As an example, George Bodner, professor of chemical education at Purdue University, points to carbon black—the black stuff you see when you burn something in the flame of a candle. Another good example, this one from David McMillin, a Purdue professor of inorganic chemistry, is graphite. Like diamond, graphite is carbon. Unlike diamond, it’s a shimmery, silvery black. So what gives?
This is where things get complicated, because the differences between diamonds and carbon black, or diamonds and graphite, happen at the molecular level.
Think about the illustration of an atom—the big ball of a nucleus surrounded by a cloud of electrons whirling through shells designated by energy level. An atom of carbon has six electrons. Two in the lowest shell, closest to the nucleus, and four in the second shell. The lowest shell can only hold two electrons, so, for carbon, that shell is full and stable—an old married couple with a minivan and a cat. But the second shell can hold eight electrons, and carbon only has half that number. That means the electrons in carbon’s outer shell are on the market. They can attract electrons from other atoms, swap and share, binding the atoms together and forming new molecules.
Once that happens, an idea called molecular orbital theory comes into play, because becoming part of a molecule seems to change how electrons go about their business. You can’t think of a molecule of two atoms as a couple of nuclei planets, each with its proprietary electron satellites still distinctly circling. Instead, the electrons of both atoms merge to the point that, when we talk about orbits, we’re talking about molecular orbits now, not atomic ones.
There are two types of molecular orbits, pi bonds and sigma bonds, and each of those has a bond and an antibond. (You can imagine them as twins, one of whom has an inherently evil moustache.) It’s the difference in bonding that makes diamonds clear and other forms of pure carbon black.
Diamonds are entirely constructed from sigma bonds. When two carbon atoms come together to form diamond, the electrons are snugly held, right in between the nuclei. The sigma bond is a tight bond. In molecular chemistry, the tightest bonds happen at the lowest orbitals … the lowest energy levels. So if your bond is very low energy, then its evil twin—the antibond—must be the opposite. Very, very high energy.
Why does this make the diamond clear? The secret is in that big difference between the bond and the antibond. When a photon of light energy slams into a stable material, it can pass through it, be absorbed, or be scattered back in the direction it came from. The net energy (or wavelength) of that photon is a critical factor. When a bunch of atoms are as tightly joined as the ones in a diamond, the photon has to have a lot of energy to be absorbed and excite an electron into an antibonding level; it’s like throwing a bowling ball at a brick wall. A molecule of diamond is like the wall. And by the time you get out the heavy construction equipment and hit that wall with enough force to take a piece out of it … well, that little piece is also going to contain a lot more energy. In this case of the photon is outside the relatively low-energy spectrum of visible light.
So it’s not really that diamonds are clear—that they don’t absorb any of the light that hits them. It’s that our eyes can’t see the colors of really high energy photons. “If you looked at it with UV eyes, you’d see something different,” McMillin said.
In graphite, on the other hand, one quarter of the bonds are pi bonds. In a pi bond, the electrons have a little bit more leeway, like toddlers on a tether. They’re still tightly held, but the nuclei don’t confine them so much and they roam more through the material. And the difference between the bond and the antibond is less extreme. If a sigma bond is a brick wall, the pi electrons are more like bowling pins. Relatively low energy photons can energize them. In fact, graphite virtually absorbs every colored photon in the visible spectrum. None come through or scatter back toward us and we therefore see black. (It’s worth noting that this absorption isn’t like a black hole, where energy has almost no chance of escaping. Instead, in graphite, the energy is absorbed, but then exits again in a changed state—as much smaller bundles of heat energy.)
The difference between sigma-bonded diamonds—which throw off photons outside the spectrum of visible light—and pi-bonded graphite—which absorbs all colors of visible light is extreme. The fact that both are carbon is pretty important, because it means that carbon is extremely versatile. And that, George Bodner said, is what makes carbon such a great element to build life around. “You need strong bonds because you want this thing held together. But there are also times when you want it to, under right conditions, to open up or react. Carbon is so good at that, better than anybody else. And life on this planet evolved around that.”
Published 10:29 am Thu, Apr 17, 2014
About the Author
Maggie Koerth-Baker is the science editor at BoingBoing.net. From August 2014-May 2015, she will be a Nieman-Berkman Fellow at Harvard University. You can follow Maggie's adventures in the Ivory Tower by subscribing to The Fellowship of Three Things newsletter.