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Last time we started talking about the allotropes of carbon, finished graphite and began with diamond.  Tonight we shall continue the diamond saga and maybe move to a third common allotrope.

Last week I was having some connectivity problems and, quite frankly, was ill with a bad cold, so I just did not feel much like writing.  I am better (much) this week and my computer seems to be functioning within design parameters.

Since the part that I wrote about diamond was so short last time, I shall paraphrase it as the start of this piece.  That way you do not have to hit the link to get up to speed.

For two things to be made of nothing but carbon (with some trace impurities), diamond and graphite could not be more different.  Whilst graphite is one of the softest materials known, diamond (specifically laboratory produced diamond, the crystals of some of which approach perfection) is the hardest material known.  Graphite is always opaque and black, where high quality diamonds are water white and clear.  Graphite is only a fair conductor of heat, whilst diamond has the highest thermal conductivity of any known substance, 7.5 times greater than copper!  Why these differences?  We have to look at the physical, and hence, the electronic structure of diamond and contrast them with graphite.

Here is a nice video showing the structures of diamond and graphite.

Note that in diamond, except at the edges, each carbon is connected to four other carbons rather than only three in the case of graphite.  Since carbon requires four bonds, in diamond all of the bonds are single bonds.  There is no aromatic character in diamond, so it is a different kind of supermolecule than is graphite.  Who says that molecules are too small to see?  A diamond is a single molecule!

Not having aromatic character has huge implications.  First, diamonds do not conduct electricity because all of the electrons are localized.  Second, they do not absorb light for the same reason.  Carbon-carbon single bonds are pretty strong, so diamonds are chemically inert for the most part and because there are not layers like in graphite, diamond is rigid and hard.  Because of the rigidity of the bonds, diamond conducts vibrations readily, that this is the cause for its extremely high thermal conductivity.

Aside for gem and investment uses where the beauty, clarity, and large size are important, the most common industrial use for diamonds is as abrasives because of the extreme hardness.  Interestingly, natural diamonds vary in hardness but even the softest diamonds are harder than most other materials.  Actually, there are a lot of diamonds mined (around 135,000,000 carets, or 27,000 kg), but only 20% of those are gem quality.  Most of the rest go for abrasives.  That is not nearly enough to meet demand, so another 110,000 kg are produced artificially for abrasives.

The hardness of diamond has to do with the crystal structure, as said before.  The reason that synthetic diamonds are harder (for the most part) than natural ones is that the crystal structure can be controlled under laboratory conditions to a greater level of perfection than in natural ones, at least for the exotic processes used to produce these special ones.  The 110,000 kg of abrasive diamonds mentioned above are not produced by the processes that are used to make nearly perfect ones, and have a hardness similar to natural stones.  But there is another parameter that can be controlled in synthetic stones that is not available in natural ones.

The carbon isotope ratio in natural diamond is around 98.9% 12C and 1.1% 13C, with some minor variation depending on the source of the carbon in natural diamonds.  In a subtle way, the heavier isotope of carbon acts as a defect in the crystal lattice, making it a bit less stable.  By using carbon sources for synthetic diamond that are depleted in 13C, nearly perfect crystals can be produced, and these are the ones that are the hardest known substance.

Let us look at gem quality diamonds for just a bit, and then get back to more geeky stuff.  The four Cs determine the value of a diamond as a gem.  This stand for carat, clarity, color, and cut.  Carat and cut can be influenced by human interaction, but clarity only marginally so and color only a bit more than clarity.

The carat is standard unit of mass for a diamond.  It is equal to 200 milligrams in the SI unit system.  A point, often used for small gems, is just 2 mg.  It is easy to see 200 mg, but seeing 2 mg required Ted Williams eyes.  Everything being equal, the higher the number of carats, the greater the value of the stone.  Do not confuse carat with karat, the term used for describing the purity of gold.  That karat (sometimes also spelt carat) is 1/24th, so pure gold is 24 karat gold.  Thus, 18 karat gold is 18/24 or 75% gold, the rest being cheaper alloying metals.

Clarity is just that, how clear the diamond appears.  The very best diamonds are just about perfectly clear.  Perfect diamond crystals are, obviously, perfectly clear, but most diamonds have inclusions, materials that are foreign such as graphite particles or other materials in them with reduce their clarity.  As with carat, the more transparent the diamond the move valuable it is.  The only way to manipulate clarity in natural diamonds is to cut off parts of the raw stone that are not nice a clear.

Color has to do with, well, the color of the stone.  In most cases, absolutely colorless diamonds are preferred and command a higher price.  However, some stones are of a beautiful color and command a premium.  These are called fancy diamonds, and I suppose that the Hope Diamond, a lovely blue stone, is the most famous of the bunch.  Intensely colored diamonds are much more valuable than pale ones, so a little bit of color decreases the value of a diamond, but a lot of color increases it significantly.

The most easily manipulated characteristic of a diamond is its cut.  This has to do with how many and how precise the facets (the flat, geometrical surfaces on a stone) are cut.  There is quite a bit of physics behind that, so we might as well get Geeky.

Diamond has a refractive index of 2.148, a very high value.  Technically, the refractive index the the ration of the speed of light in the medium (diamond) compared to the speed of light in a vacuum.  Thus, the speed of light in a diamond is only 46.6% of the speed of light in a vacuum.  That would not be of very much consequence, but the kicker is that this value if for monochromatic light (light of only one wavelength).  Related to the refractive index is the dispersion of light in a diamond.

It turns out that the refractive index in wavelength dependent.  Thus, light of a longer wavelength experiences a lower refractive index than light of a shorter wavelength.  This is how prisms work to separate out the colors of the rainbow.  Diamond as a dispersion of 0.044, a high value.  That is why a well cut diamond shows fire, the separation of the different colors of light.  That is also why cubic zirconia appears to have more fire than diamond, because it has a dispersion value of around 0.06 or a little higher.

To maximize the fire in a diamond, the cut has to be just right.  There are some physical limitations because of mechanical reasons (like not making the stone too thin at the base to survive mounts) that there is no "perfect" cut, but there are several schemes that maximize the prism effect of this remarkable stone, giving it more fire.

Now back to more technical things.  Diamond is, as I said, most valuable as an abrasive, if tonnage is any indication.  Many of you may have a diamond knife sharpener in a kitchen drawer, and they are great.  Just make sure to use either some light oil (or, as I prefer, a tiny drop of dishwashing liquid and water) to keep the steel particles from the knife floated out of the abrasive surface.

Interestingly, even though diamond is extremely hard, it is not tough at all.  In fact, just the opposite is true; diamond is extremely brittle.  Those same sp3 bonds betwixt each carbon that give it its extreme hardness also do not "like" to be disturbed with mechanical shock.  I imagine that more than one reader here has had the misfortune to drop a diamond ring "butter side down" onto a tile floor, only to see the diamond to break.  By the way, NEVER try to verify the identity of a diamond by trying to "cut" (actually just scratch the surface) of glass with it.  Physical trauma can knock off edges of a real diamond, and cheap crystals like quarts and cubic zirconia are hard enough to scratch glass as well.

Now it seems proper to discuss the scientific and experimental uses of diamond.  Remember, carbon is just above silicon on the periodic table, and one would expect that crystalline diamond might have some useful semiconductor application.  Sure enough, it does, but it is extremely difficult to grow synthetic diamonds with the right mix of dopants to for a good integrated circuit.  I think that this will happen as the technology to produce larger and better diamonds improves.  An added benefit is that, because of the high thermal conductivity of diamond, heat dispersion is less of a problem than that for silicon.  The best, most perfect synthetic crystals have a thermal conductivity of 33,200 W cm-1 K-1, thousands of times higher then that of silicon, at only 149 W cm-1 K-1.  Imagine if we would be able to product lots of good diamonds!  Heat sinks would be many times more efficient that those made of aluminum, and computer failure do to heating of the CPU would be pretty much unknown.  Even better, undoped diamond could be grown coaxially with doped diamond, forming a perfect bond betwixt CPU and the heat sink!

I guess that I should stop here, but you know that I will not!  Diamonds are NOT forever!  I mentioned before that graphite is thermodynamically more stable than diamond, and that means that all diamonds will eventually decay into graphite.  That is not a worry in the lifetime of us, or our children, or even our children's children's children.  Nice segue to The Moody Blues, do you think?  Here is why.  At 1.8 kJ mol-1 higher in energy than graphite, diamond must change.  However, at ambient temperatures the probability of undergoing the transition is extremely low.  Here is a typical graphic representation of such a process, and keep in mind that the transition from diamond to graphite has an activation energy (the peak on the graph) thousands if not millions of times higher.  But it will happen.

Photobucket

In this graphic, the energy of the diamond is on the left and the energy of the graphite on the right.  The hump in the middle is called the activation energy, and it is very high (off the chart) for the diamond to graphite transition.  If you heat a diamond to 700 or 800 degrees C the rate of transition is significant, but at room temperature it will take millions of years or longer.

This particular graphic is for a reaction with and without a catalyst (in this case, an enzyme).  The catalyst reduces the activation energy so it increases the rate of a reaction, but does not change the equilibrium since the starting and ending energies are unchanged.

Well, you have done it again!  You have wasted many more einsteins of perfectly good photons reading this rocky piece.  And even though Michele Bachmann realizes that she really is insane when she reads me say it, I always learn much more that I could ever hope to teach in writing this series, so keep those comments, questions, corrections, and other feedback coming.  I shall stay around for Comment Time as long as comments warrant (unless I visit someone this evening) and shall return tomorrow evening for Review Time.  Remember, no science or technology issue is off topic here.

Warmest regards,

Doc, aka Dr. David W. Smith

Crossposted at

The Stars Hollow Gazette,

Docudharma, and

firefly-dreaming

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