Why is nitrogen part of so many explosives?

April 5, 2004 12:04 AM

Why is nitrogen part of so many explosives?

This comment got me to thinking -- there's ammonium nitrate, nitroglycerin, TNT, et cetera...what about nitrogen makes it so reactive in these compounds. Isn't it an inert gas?

(by the way, my knowledge of chemistry is limited to my tenth-grade high school class. Most of which knowledge I've undoubtedly forgotten.)

posted by Vidiot to science & nature (12 answers total)

The volatility of the aforementioned compounds has little to do with the Nitrogen contained within it. Once Nitrogen is bound to other elements within a compound, it takes on entirely different physical and reactive properties.

The classic example given is the chemical reaction between sodium (a soft metal that reacts explosively with water) and chlorine (a poisonous green gas) to form Sodium Chloride (common table salt, which is neither explosive nor poisonous.)
posted by Danelope at 12:13 AM on April 5, 2004

Nitrogen is inert in N2 form, but not in N form. Nitrates also contain a lot of oxygen which helps them to burn stuff. So I hear.
posted by tracicle at 12:45 AM on April 5, 2004

A lot of nitrogen compounds are oxydizing agents - that is, compounds that freely "give up" oxygen. A typical explosive involves a fuel, an oxydizing agent, and something that turns into a gas on explosion, to generate an increase in volume and pressure.

As Danelope said, it's not the Nitrogen that's important, it's the Nitrate - Nitrogen bound with oxygen, often in quite unstable ways. Give it a bit of energy and the nitrogen will throw off the oxygen atoms. The oxygen atoms will react with the fuel to generate great amounts of heat. Other compounds will turn into gasses, creating an expansion and an explosion.
posted by Jimbob at 12:49 AM on April 5, 2004

I was never much of a chemist, so I don't know the details of why certain nitrogen compounds have certain properties. But as a physicist, I can come up with an explanation from an argument involving energy (which is obviously far more elegant, better than chemistry, physics is great, etc. etc.).

Pure nitrogen gas exists as molecules consisting of two nitrogen atoms, and is written as N₂. This is a very stable compound (the N₂ molecule has a triple bond). The stability of the raw state is precisely the reason that nitrogen is often found in explosives.

One way of thinking about stability is to ask, "How much energy would I need to break this bond apart?" It is clear in the case of N₂ that you would need to put in a lot of energy to break the stable triple bond.

Now consider this process in reverse. What if we were trying to somehow take single nitrogen atoms, and combine them into molecules of N₂ gas? Well, if breaking apart the bond requires us to put in a lot of energy, then forming the bond must produce a large amount of excess energy. The nitrogen atoms go from an less stable, higher energy state, to a much more stable, lower energy state, and that difference in energy has to go somewhere!

So it's the formation of the stable N₂ triple bond that provides nitrogen's contribution to the energy released in explosions -- that's the answer to your question. But this leads us to ask, "Why isn't every compound containing nitrogen explosive?"

In different compounds, nitrogen finds itself in different chemical environments (X-NO₂, Y-NO₃, etc.). Almost all of these configurations will be less stable (at least, from the nitrogen atom's point of view) than being in a nice N₂ molecule, so in theory, any of them could release a large amount of energy if we managed to get the nitrogen atoms to break away and form N₂. So how come some compounds are naturally explosive, and other compounds are basically stable?

Every reaction has something called an 'activation energy' associated with it. This is the energy we have to put into the system in order to kick-start the reaction. Even if formation of the final result (N₂) is very energetically favourable (as we have described above), the reaction will still not take place unless we can provide sufficient activation energy.

In the case of stable, non-volatile nitrogen compounds, the activation energy for an 'explosion' reaction is very large. In day-to-day existence, these compounds never get a large enough kick-start, and so they never explode.

Volatile nitrogen compounds, on the other hand, have a comparatively small activation energy -- this is precisely what makes them 'volatile' compounds! It only takes a small kick to get them to react explosively by breaking apart and reforming into N₂, and these are the nitrogen compounds that are found in explosives.
posted by chrismear at 1:10 AM on April 5, 2004

Except that you actually want compounds that require a substantial but easily achievable kick to get them going -- hence semtex and other plastic explosives that you can drop, kick, punch, or slap without it going off, and that require an explosively generated shockwave to set off the explosion.

So why do so many explosives (and rocket fuel) use aluminum or, if we're talking about Ariane, aluminium?
posted by ROU_Xenophobe at 6:15 AM on April 5, 2004

Chrismear: That pretty much sums it up from an organic chemistry POV, too. Though, the underlying reason WHY the nitrogenous group causes such instability has to do with electrons' orbitals and resonance structures. I can't really explain in detail without drawing out (though I will try), but essentially electron withdrawing and donating groups cause the N part to gain significant partial positive or negative charges leading to high-energy resonance structures (contrary to what you learned in HS or even freshman chemistry, electrons are not confined to "their element," but rather float around the molecule creating aggregate partial positive or negative charges due to their electronegativity, or ability to hold an negative charges- Oxygen (the 2nd highest electronegativity behind fluorine), which is coupled with N is explosives so much). Resonance structures are the visual representations of the multiple ways electrons can be an orbit about a molecule. The buildup of positive and negative charges due to the electron-withdrawing and electron-donating groups lead to greater and greater buildups on certain elements destabilizing the resonance structures, thus more explosive when the required activation energy is met.

If nobody understands what I just said, thats why I will never be writing an organic chemistry textbook.

ROU- I don't know the exact chemistry behind those reactions, but I do know that aluminum (which subs as Potassium Permanganate) is used at a catalyst to get reactions just started or in conjunction with other molecules to, simply put, add on or take away parts of another molecule.
posted by jmd82 at 7:21 AM on April 5, 2004

Oxygen (the 2nd highest electronegativity behind fluorine)

So if you went through the undoubtedly exceedingly unpleasant processes and made something like TNT but with flourine where the oxygens should be, you'd make bigger booms?
posted by ROU_Xenophobe at 7:54 AM on April 5, 2004

ROU -- Flourine can only carry one single bond (as opposed to oxygen's two), so you can't use it in place of oxygen.
posted by LittleMissCranky at 8:13 AM on April 5, 2004

Also, I'm impressed with these explanations. Nice to see some science types around here.
posted by LittleMissCranky at 8:14 AM on April 5, 2004

LMC: exactly...questions like this make me realize how chemistry and the world we live in is simply amazing. Think about it: everything you see is made from 92 naturally occurring elements (probably thinking of another random number (radioactive isotopes?), but go with me). With a few exceptions, every elements reacts in such a specific way that for most applications, another element will just not do. Go to the left or right on the periodic, you change the number of bonds you can made. Go up or down on the table and you have significantly changed the radius of the elements. No, only only that element will do, such as in this case with explosives where you cannot simply sub out oxygen for something else.
OK, even if I did just come from a biochemistry exam, am I being a total whacko or does anyone see where I'm coming from and find this just COOL?
posted by jmd82 at 10:55 AM on April 5, 2004

ROU: The aluminum is there to react with the oxygen in the other fuel component, typically ammonium perchlorate, (ClH₄NO₄) in solid fuel rockets. The formation of aluminum oxide releases a lot of heat.

Another very exothermic reaction that involves the formation of aluminum oxide is the thermite reaction,
2Al+Fe₂O₃->Al₂O₃+2Fe
which generates a lot of heat, but not an explosion since there aren't any gaseous products.
posted by yarmond at 1:04 PM on April 5, 2004

ROU - This is the periodic table showing electronegativities (see jmd82's explanation). You'll see that electronegativities increase as you move up and to the right. Oxygen is at the top of its column and thus there isn't anything that would make a good substitute for it. You could (perhaps) substitute it with something from below with less effect as you moved down the table.

Jmd82's very good answer moves well into the area of BSc chemistry. One of the key difficulties with chemistry for the 'layman' is that you get told a different version of how bonds form and how electrons and electron shells work at different levels of schooling (At least this is my experience in the UK- at school up to 16 you get the simple 2:8:8:18 model, after that you get a simple version of the s,p,f,d shells and finally at BSc level they start on the more complex cloud models (which alas are now at the edge of my memory)). This makes it difficult to explain the intricacies of how reactions work even to interested parties. (Hope that's not too patronising)
posted by biffa at 3:07 PM on April 5, 2004