One unsolved puzzle in the global nitrogen budget is the increasing atmospheric concentration of nitrous oxide. Measurements of the amount of N2O entrapped in polar ice show that the global concentration of nitrous oxide remained constant at about 285 ppbv until the mid 1700s. Then its concentration increased to 315 ppbv at present. The increasing concentration of nitrous oxide (0.3 %/yr), its long lifetime (~150 yr), its contribution as a greenhouse gas (global warming potential (GWP) 310 times that of carbon dioxide), and its role as a stratospheric ozone sink make it an important trace atmospheric component. Present estimates for the global budget are shown in Table I. A detailed U.S. emission summary is available.

   It is important to recognize which facts are reliable and which ones are not well defined in environmental problems. One thing that can be done with high accuracy is the measurement of atmospheric gas concentrations. For nitrous oxide such measurements have been made around the world and highly accurate data is available for the past 20 years. Concentrations from ice core samples provide a record of atmospheric gas concentrations as long as 100,000 years ago, but the data are somewhat less reliable than contemporary atmospheric measurements and require several assumptions about the equilibrium with atmospheric gases at the time of entrapment in the ice, the chemical stability of the gas trapped in ice for thousands of years, the rate of diffusion in the ice, and other issues. Such concerns are less problematic for nitrous oxide because of its low solubility in water and its low reactivity. Much uncertainity arises, however, when total budgets are constructed. For example, most estimates of sources of nitrous oxide in Table I have large estimated ranges.

   The discovery that the nitric acid oxidation of cyclohexanone/cyclohexanol to adipic acid produced significant amounts of nitrous oxide led to a voluntary change by industry to eliminate these emissions. This eliminates about 10 - 15% of the global increase, but a complete understanding of the nitrous oxide budget is needed. Ironically, the industrial sources are best understood, because point source emissions are easily quantified. Challenges remain in defining natural environmental biological emissions, as well as identifying new atmospheric reactions.

Table I. Nitrous Oxide Sources and Sinks.

Nitrous Oxide	Tg N per yr*
Natural Sources oceans wet forests dry savannas forests grasslands Anthropogenic Sources agriculture biomass burning industrial combustion adipic acid production nitric acid production	1 - 5 2.2 - 3.7 0.5 - 2.0 0.1 - 2.0 0.5 - 2.0 1.8 - 5.3 0.2 - 1.0 0.4 - 0.9 0.4 - 0.6 0.1 - 0.3
Sinks stratospheric removal	9 - 16
Net Increase	3.1 - 4.7

* 1 Tg (terragram) = 10¹² g, values from Houghton, J. T.; Callander, B. A.; Varney, S. K. Climate Change 1992. The Supplementary Report to the IPCC Scientific Assessment., Cambridge University Press: Cambridge, 1992 and Houghton, J. T.; Meira Filho, L. G.; Bruce, J.; Lee, H.; Callander, B. A.; Haites, E.; Harris, N.; Maskell, K., Cambridge University Press: Cambridge, 1994.

Historical Perspective on Ammonium Nitrate (NH4NO3). One of the largest uncertainties in Table I arises from the contribution of fertilizers to nitrous oxide emissions. There is debate about the increase in global nitrification and denitrification processes because of the increasing application of soluble nitrogen fertilizers. About 66 billion lb. of NH4NO3 are produced annually (about half as the solid). Its explosive properties have caused problems since its introduction as a fertilizer. This arises because the ammonium ion (N oxidation state = -3) and nitrate ion (N oxidation state = +5) lie at opposite extremes of the nitrogen redox series. In the solid they are poised to undergo an exothermic internal redox reaction. This occurs so rapidly once it is started that an explosion results. However, ammonium nitrate is difficult to detonate and it was once common practice to break up piles of the caked salt with explosives. In 1921 about 9 million lb. of a solidified mixture of ammonium nitrate and ammonium sulfate detonated in Oppau, Germany, while being broken apart with explosives. It resulted in the death of over 500 people. The next major loss of life occurred on April 16, 1947, in Texas City, Texas. Two adjacent ships, containing a total of 6.5 million lbs of ammonium nitrate, caught fire, detonated, and caused over 600 fatalities and 300 injuries. A war correspondent for the Associated Press likened the destruction to Nagasaki, victim of the second atomic bomb. Later that year on July 28 in Brest, France another shipload of 6.6 million lbs of ammonium nitrate exploded and killed 21 persons. Accidents still occur, such as the Dec 13, 1994 explosion at the Terra Industries' fertilizer plant in Iowa.

   The wide availability of ammonium nitrate has often led to its misuse as a crude terrorist explosive (e.g., the 1995 Oklahoma City bombing). At about 220^o C decomposition begins and NH4NO3 may explode above 350^o C.

      NH4NO3 ---> N2O + 2H2O ---> N₂ + 1/2 O₂ + 2H₂O

At lower temperatures N2O is the primary product, but under explosive conditions decomposition continues to yield N2 and oxygen. The excess O2 is itself an oxidant. Adding fuel oil to ammonium nitrate causes rapid combustion with the excess O2, and doubles the explosive yield. That is why the terrorist explosives are often called fertilizer/fuel oil bombs.

References:

1)Silver, R. G.; Sawyer, J. E.; Summers, J. C. Catalytic Control of Air Pollution : Mobile and Stationary Sources; American Chemical Society: Washington, 1992.
2)Thiemens, M. H.; Trogler, W. C. Science 1991, 251, 932-934.
3)Trogler, W. C. J. Chem. Educ. 1995, 72, 973-977.
4)Gray, H. B.; Simon, J. D.; Trogler, W. C. Braving the Elements; University Science Books: Sausalito, 1995.