There is a weird smattering of organic oxides whose molecules contain a foreign oxygen latched on to an otherwise familiar framework. I’ve written about DMSO before, which is essentially dimethyl sulfide with an extra oxygen atom along for the ride. N-Oxides fit into this group of compounds as well.
Perhaps no molecule better typifies the class than nitrous oxide, N2O. Even the molecular formula sets off neural fireworks: that can’t be right. A central nitrogen flanked by N and O? Something’s wrong here. The central nitrogen seems overworked, while the oxygen seems to be missing a bond. Despite its bizarre structure, nitrous is surprisingly unreactive—I learned this recently while helping out a teacher friend with one of his students’ science fair projects. Thanks to its lack of reactivity at sane temperatures, detecting N2O is a pain.
Synthesizing nitrous, on the other hand, is quite easy. Upon heating, ammonium nitrate breaks down into N2O and two water molecules. The melting point of NH4NO3 is downright eye popping for an ionic compound: 170 ºC. Some rather unsafe methods actually produce nitrous from molten ammonium nitrate at high temperatures.
NH4NO3(l) → N2O(g) + 2 H2O(g)
One must be careful as the dissociation of ammonium nitrate into gaseous nitric acid and ammonia competes with N2O formation (“decomposition”).
NH4NO3(l) → HNO3(g) + NH3(g)
Dissociation is endothermic and decomposition exothermic, so heating can set up an interesting situation where the dissociation reaction can “quench” the heat released by decomposition. When dissociation is suppressed, the decomposition reaction can become a runaway exotherm. Although this document on the safe production of nitrous oxide from ammonium nitrate is written for chemical engineers—they even go to the trouble of writing out “gram-mole”!—it’s an illuminating read with some additional information about the synthesis of nitrous oxide.
I made nitrous for quick delivery using the (awesome) syringe method of Mattson. Decomposing NH4NO3 in aqueous solution is much safer that working with the explosive salt in molten form. A small amount of chloride ion catalyzes decomposition, meaning the reaction can be carried out at lower temperatures and dissociation is a non-issue. The reaction is run in a syringe, so the gas forms right into its storage and delivery vessel with no air contamination.
First, a hot water bath at 95 ºC was prepared in a 600 mL beaker. In a test tube, 1.0 g of NH4NO3 and 0.050 g of NaCl were dissolved in 5 mL of 6 M aqueous nitric acid. The solution was taken up in a 25 mL syringe, the syringe capped with a closed stopcock, and the assembly placed in the water bath. (Although I didn’t lubricate my syringe first, this is recommended to prevent leaks.) Bubbles of slightly yellow gas formed and gradually pushed the plunger up. Once 15 mL of gas had formed (about 15 minutes), the stopcock was opened and most of the liquid was pushed out of the syringe. (A 60 mL syringe and longer reaction time can be used to collect more gas.) The stopcock was closed and the syringe placed nozzle down in a beaker for storage.
With safe technique, this is a painless way to synthesize a little N2O(g) in less than an hour. Now, what to do with it? Perhaps I’ll connect the syringe and stopcock to a pressure gauge and put Boyle’s law to the test!