A few weeks ago I tried to make nitrocellulose, with mixed results:

A large portion of common explosives— including TNT (trinitrotoluene) and dynamite (nitroglycerin)—have the key prefix “nitro” in them. This is not a coincidence; nitrogen groups, when added to benign organic compounds, tend to suddenly make them more reactive. The full chemistry behind this is complex, but a good thing to keep in mind is that nitrogen, when added to a compound consisting predominantly of carbon, has the net effect of destabilizing the electronic structure of the molecules. Carbon has four electrons, and so it likes to form four bonds with neighboring molecules—either in the form of four single bonds, two doubles, or a triple and a single. In the chemical theory of Lewis structures, this property is the basis of the “octet rule”: if we count each single bond as two electrons “paired together” with a covalent bond(with one electron coming from the carbon, and the other from whatever it bonds to), then carbon is content when it has eight electrons around it—four from itself, and four from the various things to which it bonds. Nitrogen, on the other hand, has five electrons when it is not bonded to any other atoms. But rather than always forming five bonds, nitrogen and most other atoms prefer to stick to the octet rule, and so two of nitrogen’s electrons form an internal bond within the molecule(called a lone pair), the members of which don’t bond to anyone outside the molecule. As a result, nitrogen usually ends up forming three single bonds or a single and a double bond, with its lone pair of electrons allowing it to still have eight total electrons. While there are some great exceptions to this covalent bonding scheme(much to the dismay of high schoolers everywhere), the Lewis octet rule for drawing covalent molecules provides great insight into the relative stability of chemical compounds.

While the Lewis structure rules give us excellent insight into how organic molecules bond, they don’t actually tell us what the molecule looks like at the atomic level. As quantum mechanics has predicted(and modern molecular imaging routines like NMR have verified), the structure of an organic molecule is heavily influenced by the addition strange atoms like nitrogen. Nitrogen has relatively high electronegativity, which means that it tends to attract electrons to itself. The derivation of why nitrogen and other gases(particularly diatomic gases like halogens, which tend to covalently bond in pairs to form compounds like O2 and N2) is rather complex(see Linus Pauling’s derivation, or consult Wikipedia’s summary of it), but it essentially involves the idea that certain atoms just accept electrons in bonds much more readily—the reasons involve nuclear shielding, dissociation energy, and a litany of other complex quantum effects. But the relevant result is that, even in large organic molecules the electrons tend to be slightly delocalized—they like to move into parts of the molecule beyond their bonding pair, where they traditionally would be expected to stay. The electronegative nitrogen atom tends to pull all of them very slightly towards itself, ensuring that every electron pair has slight asymmetry due to the pull of the nitrogen group. While this pull varies based in distance in the molecule, it has the net effect of destabilizing bonds throughout the molecule. This allows the organic molecule to fall apart much more easily, making the decomposition reaction explosively exothermic.

Thus nitration of organic compounds with nitric acid is a powerful way of making normal compounds become explosive. The reaction itself is often carried out in an ice bath with nitric acid, and it usually must be carefully monitored to prevent the nitrogen groups from substituting in for too many carbons in the compound(an effect known as “runaway nitration,” which either creates a more dangerous product or merely an inert product), as well as to ensure that the compound does not detonate prematurely. The above video shows the combustion of “guncotton,” which is the result of nitrating the cellulose polymers that comprise ordinary household cotton. For obvious reasons, this compound is known as “nitrocellulose.”

A key thing to notice is that the guncotton dissolves in acetone—ordinary cotton does not do this quite so readily. This is because the electronegative nitrogen succeeds in making parts of the cellulose molecules polar—because they attract some of the electrons towards themselves, the molecules are left with a positive and a negative end that was not present before they had nitrogen attachments. This polarity allows the highly polar acetone molecules to easily pick apart the polymer chains and dissolve the cotton. If the acetone is then allowed to evaporate, the polar molecules recombine into solid nitrocellulose in a manner similar to when sugar water evaporates and leaves behind rock candy.


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