[pnacze199204]

Why is diamond so chemically unreactive? As is the case with hardness, does it have to do with the arrangement of the atoms of the mineral's crystal lattice? Why, for example, is sulphuric acid unable to dissolve it, but sodium nitrate does?

[rolnor]

Do you know what a diamond is made of? Its chemical composition?

[Corribus]

A simple answer is that the carbon atoms in diamond are arranged in almost perfect tetrahedral lattice. This makes them very stable. C-C bonds are (in isolation) also reasonably strong.

A more nuanced explanation would require examination of the diamond surface. Solids predominantly react at their surfaces, so it is the surface that really determines the speed with which a hunk of something will change into something else. As you may imagine, even though carbon atoms are arranged in perfect tetrahedral arrangement, that situation ends at the edge of the diamond. Atoms on the surface have what are sometimes called "dangling bonds". This is a spot where an atom on the surface wants to be bonded to another atom, but (because it's on the edge) isn't. That can make for an unhappy situation, and it is usually where reactions take place. (If you've seen the videos of penguins in Antarctica huddling in a big mass to keep warm, imagine for a moment how the ones on the edge feel and that should given you a nice idea of what's going on.)

Surface science is complicated and the reactivity depends on a lot of things, not least of which is the specific crystal plane that is exposed to the environment. I am no expert in the surface science of diamond, but a brief look at the literature seems to indicate that the "dangling bonds" on the surface carbon atoms of diamond are usually passivated by hydrogen atoms. C-H bonds are also incredibly strong. With a strong enough base you could probably rip those off and start to do some chemistry at the dangling bond sites on the diamond surface - most likely reacting with atmospheric oxygen, but bearing in mind that any such chemistry leading to destruction of bulk diamond would also require breaking a lot of C-C bond underneath, so such decomposition would still be extraordinarily slow.

Which brings up one other point worth mentioning: the distinction between kinetic and thermodynamic stability. We think of diamond as very stable, but it's not the most stable allotrope of carbon, thermodynamically speaking (at standard temperature and pressure). The standard state of carbon is actually graphite. What this means is that carbon atoms in diamond are actually happier as graphite than they are as diamonds. By implication, if you left a diamond on your table and waited long enough, You'd come back to find a lump of graphite instead.

Why doesn't this just happen, then? Well, rearranging the carbon atoms in diamond to their more preferred graphite orientation would require breaking lots of carbon bonds so that you can make new carbon bonds in a different orientation. That breaking of carbon bonds takes a lot of energy. So even if you get that energy back (and more), it is still a (very) slow process to get things going. Which is why diamonds are forever.

One might also wonder why diamonds don't just react with oxygen in the atmosphere.  Consider the reaction of complete oxidation of bulk diamond to carbon dioxide in the presence of oxygen:

C (diamond) + O₂ (g) --> CO₂ (g)

At room temperature, the standard molar change in enthalpy for this reaction is -395.4 kJ/mol. This is how much heat is either released to or sucked up from the environment during the reaction. The negative value tells us the reaction is quite exothermic, that a lot of heat is released during the reaction, and that the products are more thermodynamically stable than the reactants.

The standard molar entropy change is +6.2 J/K/mol. The reaction results in a slight gain of entropy. Pretty much a wash here because one molecule of gas is produced for every molecule of gas that's consumed.

And the standard Gibbs energy change is -397.3 kJ/mol. The negative value here tells us that the reaction is spontaneous.This is particularly relevant, because it signals that under standard conditions, diamond is quite happy to be obliterated and converted into carbon dioxide. The driving force is indeed pretty large, given how exothermic the reaction is. And yet, we don't see this happen because* the requirement of breaking lots of very stable carbon-carbon bonds is just too steep a hill to climb. There's not enough energy hanging around at room temperature in order to get this to happen efficiently. I guess that if you had some magic microscope where you could zoom in and examine each carbon atom on the surface of a diamond, you'd probably occasionally see one with an oxygen atom bonded to it, but it just doesn't happen enough to manifest a significant change to the macroscopic properties of the diamond.

Ultimately diamonds ARE chemically stable, but not for the reason you might think. Carbon atoms in diamond would rather be somewhere nicer, but they're just too lazy to go there. In that way, they are very much NOT like penguins.

*There's another reason for the slow speed as well. The electrons in oxygen in its most stable form are arranged in such a way that their spins are unpaired. Because they are unpaired, a molecule of oxygen has a relatively large total electron spin angular momentum. None of the other reactants or products have unpaired electrons. So, the reaction violates, in a way, the conservation of momentum. Getting the reaction to go also requires enough energy to force the electrons in oxygen to pair so they have no net angular momentum, which creates a kinetic barrier. This is, by the way, why all things made of carbon don't just spontaneously combust in Earth's atmosphere.

References: thermodynamic data for oxidation of diamond was calculated from values reported here: https://www2.chem.wisc.edu/deptfiles/genchem/netorial/modules/thermodynamics/table.htm

[rolnor]

I think what you write is true for all elements, non will react and start to burn in air/oxygen? Some oxidize, like the alkali metals, but even they do not start to burn? Its simply to high energy of activation.
Edit: One exception would be white phosphorous, but that requires some heat, like the touch of a hand or finger
Edit: A interesting comparison is fluorine, many elements and substances catch fire in fluorine-gas. The bonds between fluorine atoms in fluorine-gas are not so stable compared to the bonds between the oxygen atoms in oxygen-gas. Compare this with ozone, ozone will react at room temperature with charcoal to form carbon dioxide. Oxygen gas would be very reactive because of its extreme electro-negativity, but it's hold back by the relatively strong bonds in oxygen-gas. So maybe diamond would oxidize slowly in ozone? You could check this if you weigh a diamond very carefully with a scale and then kept it in a concentrated stream of ozone for a week, and then weigh it again. Maybe you would lose a few nano-grams. Maybe in a 100% stream of ozone you could get charcoal to catch fire?

[Corribus]

Reactions with oxygen are frequently favored enthalpically because oxygen forms strong bonds with other elements but the (sigma) bond between two oxygen atoms is fairly weak due to electron pair repulsion. However, these reactions may be unfavorable from an entropic standpoint, at least in the case of oxidation of noncarbonaceous solids, when there is a net reduction in the number of gas molecules as the reaction proceeds. Regardless, with some exceptions, reaction with O₂ is thermodynamically favorable in many cases. This does not imply that these reactions are fast, of course, as many have large activation barriers due to strong bonds that must be broken in the substrate or spin angular momentum selection rules that drastically reduce the probability of individual reaction events being completed. In non-carbonaceous substrates, some of these barriers are reduced, either because the substrate is also paramagnetic, heavy atoms relax the selection rules (that's just conjecture on my part), or metal atoms involved, as they often do, catalyze otherwise unfavorable reactions.

Ozone is a far more vigorous oxidant than O₂. Ozone is diamagnetic, so spin states are no longer a kinetic barrier to reaction.

Also note that while O₂ is paramagnetic in its ground state, at higher temperatures a substantial portion of paramagnetic O₂ is converted into diamagnetic singlet O₂, which has none of the impediments that hinder it's less reactive analog. Maybe it helps to consider some real numbers here.

At room temperature, virtually all oxygen exists in the triplet state. (The amount of singlet oxygen is on the order of 1 part in 10¹⁷.) In a typical combustion engine, the temperature might be about 2500 °C (2773 K). The energy gap between singlet and triplet oxygen is about -94 kJ/mol. We can use the Boltzmann distribution to estimate the fraction of all oxygen that exists in the singlet and triplet states at 2773 K. That calculation comes out to be about 1.7% singlet oxygen and the remainder triplet oxygen. This serves to explain why gasoline does not spontaneously combust under air under normal conditions, but reacts quite vigorously in the presence of an ignition source. (Aerosolization of the fuel and compression of the gas also has a lot to do with it, not to mention all that energy around to help surmount the more convention Arrhenius reaction barrier.)

Anyway, point is, to get back to the original question, at room temperature most oxygen is paramagnetic and just reacts very sluggishly with organic substrates - even relatively volatile ones like gasoline, to say nothing of diamonds. Crank the temperature of a pure oxygen atmosphere up to 2500 degrees, or swap out the oxygen for ozone, or both, and things start to change because some of those kinetic barriers disappear. While I guess it's unlikely you'd see the diamond spontaneously explode under those conditions, you would almost certainly start to see some slow but observable combustion of the diamond surface. Hold it there long enough and the diamond would probably disappear, particularly if it had some impurities that could help speed the combustion along.

[Borek]

Crank the temperature of a pure oxygen atmosphere (...) you would almost certainly start to see some slow but observable combustion of the diamond surface. Hold it there long enough and the diamond would probably disappear, particularly if it had some impurities that could help speed the combustion along.

I recall reading as a kid about an experiment with diamonds - done long ago, perhaps even before "alchemy" became "chemistry" (as if the dividing line was sharp) - someone tried to melt two smaller diamonds (heating with a sun/lens?) to produce a larger one, but instead of combining they disappeared. But I don't remember the source, so I am not sure it wasn't an urban legend.

[rolnor]

Crank the temperature of a pure oxygen atmosphere (...) you would almost certainly start to see some slow but observable combustion of the diamond surface. Hold it there long enough and the diamond would probably disappear, particularly if it had some impurities that could help speed the combustion along.

I recall reading as a kid about an experiment with diamonds - done long ago, perhaps even before "alchemy" became "chemistry" (as if the dividing line was sharp) - someone tried to melt two smaller diamonds (heating with a sun/lens?) to produce a larger one, but instead of combining they disappeared. But I don't remember the source, so I am not sure it wasn't an urban legend.

Here you can see this IRL: https://www.youtube.com/watch?v=WWpm6_Y7ASI

[pnacze199204]

So, if we converted that 20% of the oxygen in our air into ozone, it would mean that at a standard temperature virtually everything would oxidise because there would be no kinetic barrier, right? I don't have a chemistry background so some of the issues are hard for me to understand. Thank you in advance for your reply!

[Borek]

So, if we converted that 20% of the oxygen in our air into ozone, it would mean that at a standard temperature virtually everything would oxidise because there would be no kinetic barrier, right?

Define "virtually everything" - such general statements are almost never safe to use, there is always plenty of fine prints and details.

But yes, plenty of things that we consider stable would catch a fire almost immediately.

[pnacze199204]

I mean, for example, concrete, rubber, wood, paper, etc.
The ozonosphere starts at an altitude of 10 km, so why then don't planes or rockets flying into space explode? Is the concentration of ozone in the ozonosphere too low for this to happen? 

[Borek]

Check amount of ozone in ozonosphere. It is not what you think it is.

I mean, for example, concrete, rubber, wood, paper, etc.

In concrete everything is already oxidized, it doesn't burn.

[Corribus]

So, if we converted that 20% of the oxygen in our air into ozone, it would mean that at a standard temperature virtually everything would oxidise because there would be no kinetic barrier, right?

Virtually every reaction has a kinetic barrier. Ozone will oxidize many things because it is thermodynamically favorable, but it still takes time. Ozone is more reactive than oxygen, but it doesn't happen instantly.

The ozonosphere starts at an altitude of 10 km, so why then don't planes or rockets flying into space explode?

Pressure and temperature, for two. Also in many cases even if a material may be oxidized readily, oxidation only happens at the material surface where the oxidizing agent can interact with the material. Typically an oxidized layer is formed, which can protect the pristine material underneath from further oxidation because the oxidizing agent cannot readily permeate through the oxidized layer.

You might imagine a layer of rust on the surface of iron. Rust is iron oxide. But the iron underneath cannot become oxidized until the rust flakes off, exposing fresh iron. In the case of iron, this happens pretty easily because iron oxide and metallic iron do not adhere well to each other. This is not always the case - in some metals, the oxide layer is bound very tightly to the surface and so the material may persist indefinitely, protected by the oxidized layer. Examples you may be familiar with are the green verdigris that covers bronze/copper statues and the tarnish that covers silver platters and flatware. (To be clear, these are not usually formed only through reaction with oxygen but involve other atmospheric gasses as well). Those layers may only be nm to microns thick. In the case of tarnish, easily removed by a polish, but the layer will be formed again in due time. Every time you polish, you are removing some of the silver, so in principle if you polished enough time, your silver spoons would vanish.

[pnacze199204]

Gasoline, diamond and many other materials in the presence of oxygen, oxidize slowly at room temperature unless they have access to a heat source. How is this the case with ozone? Would a higher temperature be needed here, too, to overcome the activation energy to start the chemical reaction?

[Hunter2]

It is the same, but ozone is more instable and start earlier a chemical reaction with organic material as normal oxygen.Activation energy is more lower.

[rolnor]

The concentration has nothing to do with energy of activation? If something does not react in 1% ozone, it does not react in 100% ozone. If you want very strong ozone I think you can electrolyse ammonium persulfate or potassium persulfate I think.