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Common saying. Diamond possesses:

  • ultra hardness, (10 on the Mohs scale; 10000 HV on Vicker's Hard Test (iron merely 30-80))
  • hyper thermal conductivity, ($2320~\mathrm{W\, m^{-1}\, K^{-1}}$, or over ten times better than the heatsink in your computer!)
  • extreme pressure resistance, (withstands a crushing 600 gigapascals; or around 2 times the pressure at the center of the earth, enough to snap carbon nano-tubes and graphene or create metallic oxygen or overcome copper's electron degeneracy pressure, making the maximum chamber pressure of a firing pistol seem literally like popping popcorn... I digress)
  • and excellent luster (what do you expect, it's a diamond) combine to make the gemstone coveted by all.

Diamonds are the stuff of awesome.

But do they really exist forever?

Wikipedia notes that,

Diamond is less stable than graphite, but the conversion rate from diamond to graphite is negligible at standard conditions.

Huh. But Wikipedia doesn't mention how long. So how long would it take for this super-material to convert to the stuff I scribble with?

(If you doubt the claims about diamond's seemingly unbelievable properties, check out the link on Wikipedia about diamond and this and this.)

Great point Joe made, that $10^{80}$ is just forever to us puny humans. Being a geek I can't resist the urge to compare the time length $10^{80}$:

  • Makes the entire lifespan of a red dwarf star seem like the Planck second.
  • Enough time for you to sift through all the atoms in the entire UNIVERSE at a rate of one atom per second.
  • Getting $67 worth of US quarters and flipping them, one per second, to get all heads-up.
  • Chance of macroscopic quantum tunneling!! (I don't know precisely how much, but quite large)

...and this is $10^{80}$ seconds I'm talking about...

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    $\begingroup$ More important question what happen when you compress dimond even more. Like 100x more pressure. Does it become a super diamond or black hole dimond or something like that??? $\endgroup$
    – bodacydo
    Commented Jul 20, 2015 at 2:11
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    $\begingroup$ @bodacydo I recently read that it was just discovered that they can get stronger under pressure. SciAm online perhaps, but I'm not sure. It might be used for "inner" anvils in a nested pressure anvil situation, it mentioned. I seem to recall that phase changes are involved, with different atomic bonds. $\endgroup$
    – JDługosz
    Commented Jul 20, 2015 at 3:08
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    $\begingroup$ Not directly related to the question, but the "pressure resistance" and "hardness" do not mean that you can't destroy them. Diamonds are brittle - hit them hard enough with a hammer and they shatter. Also, remember that there is a thing called carbon dioxide. It's made of diamond/graphite and oxygen (and there's no lack of that in our atmosphere). Throw a diamond in a fire and it burns away as CO2. $\endgroup$
    – Gimelist
    Commented Jul 20, 2015 at 8:23
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    $\begingroup$ You can destroy diamonds, so they're not really forever - in fact, diamonds burn, releasing just the usual carbon dioxide you'd expect. The ignition temperature is rather high, but the burning is self-sustaining after ignition. Also, while they are incredibly hard, this also leads to them being rather fragile - sure, they've got great pressure resistance, but they don't respond well to shock. You just need a simple hammer to break a diamond into lots of shards. That said, you'd be hard pressed to find something that scratches a diamond - so for jewelry purposes, they are forever. $\endgroup$
    – Luaan
    Commented Jul 20, 2015 at 8:29
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    $\begingroup$ De Beers needed a slogan for diamonds that expressed both the theme of romance and legitimacy. An N. W. Ayer copywriter came up with the caption "A Diamond Is Forever," [...] Even though diamonds can in fact be shattered, chipped, discolored, or incinerated to ash, the concept of eternity perfectly captured the magical qualities that the advertising agency wanted to attribute to diamonds. Within a year, "A Diamond Is Forever" became the official motto of De Beers. From The Atlantic 1982 $\endgroup$
    – CodesInChaos
    Commented Jul 20, 2015 at 10:59

4 Answers 4

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how long would it take for this super-material to convert to the stuff I scribble with?

No, despite the fact that James Bond said "Diamonds are Forever", that is not exactly the case. Although Bond's statement is a fair approximation of reality it is not a scientifically accurate description of reality.

As we will soon see, even though diamond is slightly less stable than graphite (by ~ 2.5 kJ/mol), it is kinetically protected by a large activation energy.

Here is a comparative representation of the structures of diamond and graphite.

structures of diamond and graphite

(image source: Satyanarayana T, Rai R. Nanotechnology: The future. J Interdiscip Dentistry 2011;1:93-100)

environment of a carbon in diamond and graphite

(image source)

Note that diamond is composed of cyclohexane rings and each carbon is bonded to 2 more carbons external to the cyclohexane ring. On the other hand, graphite is comprised of benzene rings and each carbon is bonded to only 1 carbon external to the benzene ring. That means we need to break 6 sigma bonds in diamond and make about 2 pi bonds (remember it's an extended array of rings, don't double count) in graphite per 6-membered ring in order to convert diamond to graphite.

A typical aliphatic C–C bond strength is ~340 kJ/mol and a typical pi bond strength is ~260 kJ/mol. So to break 6 sigma bonds and make 2 pi bonds would require ~((6*340)-(2*260)) ~ 1500 kJ/mol. If the transition state were exactly midway between diamond and carbon (with roughly equal bond breaking and bond making), then we might approximate the activation energy as being half that value or ~750 kJ/mol. Since graphite is a bit more stable than diamond, we can refine our model and realize that the transition state will occur a bit before the mid-point. So our refined model would suggest an activation energy something less than 750 kJ/mol. Had we attempted to incorporate the effect of aromaticity in graphite our estimate would be even lower. In any case, this is an extremely large activation energy, so, as we anticipated the reaction would be very slow.

An estimate (see p. 171) of the activation energy puts the reverse reaction (graphite to diamond; but since, as noted above, the energy difference between the two is very small the activation energy for the forward reaction is almost the same) at 367 kJ/mol. So at least our rough approximation was in the right ballpark, off by about a factor of 2. However, it appears that the transition state is even further from the midpoint (closer to starting material) than we might have guessed.

This activation energy tells us that at 25 °C, it would take well over a billion years to convert one cubic centimeter of diamond to graphite.

Note 04/17/20: As mentioned in a comment, the original "estimate" link became defunct and was replaced today with a new estimate link. However the original article and estimate can can still be seen on the Wayback Machine and it estimates the activation energy to be 538.45 kJ/mol, reasonably close to our estimate.

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    $\begingroup$ I always wondered about that: if mineral diamonds are as much as a billion years old, why are they not mixed state due to some atoms rebinding? If it happens on the surface of the crystal only, then we'd find diamonds encrusted in graphite in the matix, rather than directly in the matrix. It does not seem that diamonds have decayed before being mined. $\endgroup$
    – JDługosz
    Commented Jul 20, 2015 at 3:14
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    $\begingroup$ Taking the calculations shown in the link containing the estimated activation energy and extending them to $\mathrm{298\ K}$ (which is a simple but coarse approximation), then the conversion of diamond to graphite takes place on times of the order of $\mathrm{10^{80}\ yr}$; this is several orders of magnitude higher than the characteristic timescale of solar-mass black hole evaporation ($\approx \mathrm{10^{67}\ yr}$)!!! Compare also with the current age of the Universe, at a measly 13.8 billion years ($\mathrm{1.38×10^{10}\ yr}$). $\endgroup$
    – Nicolau Saker Neto
    Commented Jul 20, 2015 at 3:35
  • $\begingroup$ @JDługosz because they spent most of that time deep in the Earth, where they are indeed stable. It's only recent (geologically speaking) that they were transported to the surface in volcanic eruptions. $\endgroup$
    – Gimelist
    Commented Jul 20, 2015 at 8:21
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    $\begingroup$ Seems to me that $10^{80}$ years is sufficiently long to be synonymous with "forever"... $\endgroup$
    – Joe
    Commented Jul 20, 2015 at 14:43
  • $\begingroup$ The link provided under "estimate" is now defunct; is there an alternative source one can check out for more information? $\endgroup$
    – user39221
    Commented Aug 31, 2018 at 3:48
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I add an other answer : it depends what kind of people you are, if you are as crazy as them, of course not !

2012 Christmas Lectures - Burning a Diamond

But ron's answer is better sure ! ;)

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Well, it seems that diamond is not forever because upon exposure to sunlight it loses atoms.

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    $\begingroup$ This article seems pretty interesting. Could you give a more in-depth summary of it? We tend to avoid link only answers on this site. $\endgroup$
    – Tyberius ♦
    Commented Sep 5, 2017 at 13:58
  • $\begingroup$ Not upworthy answer, but +1 for the interesting link $\endgroup$
    – electronpusher
    Commented Oct 14, 2019 at 17:04
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Diamonds burn, but the temperature at which they burn depends on whether or not the diamonds are in contact with air. The temperature of diamond ignition in pure oxygen is 690º C to 840º C.  In a stream of oxygen gas, diamonds burn at a low red heat initially. They will gradually rise in temperature and reach a white heat. Then, the diamonds will burn uninterruptedly with a pale blue flame, even after the oxygen heat source is removed. The diamond crystals will gradually decrease in size and finally disappear. The flame at the last moment will flicker brightly and then disappear, leaving not a trace of ash or residue.

For this to take place in an air mixture, the heat must be kept directly applied on the diamonds at all times. If removed, the diamonds will not continue to burn, due to oxygen being diluted with nitrogen that does not support combustion

The most famous gems on Earth have new competition in the form of a planet made largely of diamond, astronomers say.

The alien planet, a so-called "super-Earth," is called 55 Cancri e and was discovered in 2004 around a nearby star in our Milky Way galaxy. After estimating the planet's mass and radius, and studying its host star's composition, scientists now say the rocky world is composed mainly of carbon (in the form of diamond and graphite), as well as iron, silicon carbide, and potentially silicates.

At least a third of the planet's mass is likely pure diamond, info

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  • $\begingroup$ It seems from the Wikipedia article that the composition of 55 Cancri is hypothetical: "much of which may be in the form of diamond", so "likely" seems a stretch. Still, pretty cool. $\endgroup$
    – Buck Thorn ♦
    Commented Apr 18, 2020 at 8:46

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