Nitrogen Oxides in the Atmosphere


Nitrogen is one of the few elements that readily forms strong multiple bonds. There is a very strong triple bond between nitrogen atoms in molecular nitrogen. The N=N bond energy of azides is less than half the energy of a triple bond and N-N bonds are quite weak. Single N-O bonds are considerably weaker than N-H bonds. A summary of N-X bond energies (in kJ/mol) is below. Due to the remarkable strength of the bond in N2, the molecule is chemically and photochemically unreactive.



Ammonia and other volatile, organic amines are released from the decomposition of plant and animal materials. These readily oxidize so most of the nitrogen compounds in the atmosphere, apart from molecular nitrogen, are nitrogen oxides. The oxidation state of nitrogen varies from I to V in the oxides.
    (a) N2O, or nitrous oxide, and some NO, or nitric oxide, are produced by bacteria in the soil and in the oceans. Nitrous oxide is relatively unreactive in the troposphere and is the main source of nitric acid and other nitrates in the stratosphere. This is important in the chemistry related to ozone destruction in the stratospheric ozone layer. Nitrous oxide is relatively non-toxic and is used as an anesthetic and as a propellant for whipped cream in cans.

    (b) Nitric oxide (NO) results from the combination of O2 and N2 in lightning strikes. It is a radical and so is very reactive in the atmosphere. Its unpaired electron can be removed through oxidation to NO+ or the molecule can be reduced to NO-. In biology, NO is important as a signaling agent. Reaction of NO with oxygen produces nitrogen dioxide.

    (c) Nitrogen dioxide (NO2) is a brown gas, responsible for the color of photochemical smog. This is produced by the combustion of fossil fuels and biomass as well as by the oxidative decomposition of ammonia in the atmosphere. It is a radical and so is reactive in the atmosphere. Nitrogen dioxide is in equilibrium with its colorless dimer, N2O4.

    (d) Dinitrogen tetraoxide (N2O4) has a very weak N-N bond and is in equilibrium with NO2. It is a diamagnetic substance. It reacts with water to form nitrous and nitric acids and it oxidizes many metals.

    (e) Other neutral nitrogen oxides include N2O5 and N2O3.

    (f) Nitrogen trioxide (NO3) is a radical with an unpaired electron on one of the oxygen atoms. Its reactivity is similar to that of the hydroxy radical (HO).

    (g) With the exception of the unreactive nitrous oxide (NNO), the neutral nitrogen oxides are commonly grouped as NOx. These are toxic to humans and other organisms and are important in the production of ground-level ozone and photochemical smog.





Reactions of NOx Radicals

Three of the nitrogen oxides are radicals. NO and NO2 are nitrogen-centered radicals.

The more reactive NO3, like OH, has an unpaired electron on oxygen.

NO typically reacts with organic peroxy radicals to form NO2 and an alkoxy radical. The NO2 can add to alkyl peroxy radicals. In the example below, acetaldehyde is oxidized and the resulting peroxyacetyl radical reacts with NO2 to form peroxyacetic nitric anhydride. PAN is a very irritating molecule that is toxic to humans, animals, and plants.



Nitrogen trioxide is an important oxidant at night. It is photolyzed by sunlight. Below is a summary of some of the more important reaction types involving NO3.




Thermodynamics

Most of the nitrogen oxides are unstable with respect to molecular nitrogen and molecular oxygen and have a positive value for the Gibbs free energy change of formation. The table above give the standard heats of formation from the elements (H0). The relationship of the free energy (G) to the enthalpy (H) and entropy (S) is given by the equation:

G = H - TS


For the formation of NO from N2 and O2, the entropy change is small and positive.

N2 + O2 2 NO.........G0 = 86.7 kJ/mol


The Gibbs free energy change relates to the equilibrium constant, K. R is the gas constant 8.314 J/K and the temperature, T, is measured in Kelvin.

G = -RT ln K = -5.7 log K (in kJ at 298 K)


At equilibrium, that is under thermodynamic control, the concentration of NO in the atmosphere should be 10-15.5 atm but actual concentrations of NO are significantly higher than this (approximately 10-10 atm).

So, even though thermodynamics says that NO is unstable with respect to N2 and O2, it doesn't spontaneously break down to those elements. There is a high activation barrier that must be overcome between reactants and products. Nitric oxide is kinetically stable but thermodynamically unstable.

Nitrogen dioxide is a brown gas that is in rapid equilibrium with its colorless dimer. This is a reaction that is controlled by thermodynamics because the activation barrier between product and reactant is small. The equilibrium is temperature dependent as you can see from the picture below (see: University of Michigan website for more on the demonstration.)



At low temperature the Gibbs free energy is negative and the equilibrium favors product, the dimer. At higher temperature, Gibbs free energy is positive and favors the equilibrium favors the reactant.

In summer, the air temperature is higher. There is a greater concentration of NO2 that can be photolyzed to produce ozone. Colder temperatures in the winter tie up some of the nitrogen dioxide as the dimer and ozone production is reduced.


Kinetics

The rate of a reaction typically depends on the activation energy of the slowest step in the reaction, the concentration of reactants that come together in that step, temperature, and pressure.

In the troposphere, the concentration of nitrogen oxides is low. We can talk about the concentration of gases in the atmosphere in terms of partial pressure. Near the Earth's surface, the pressure is 1 atmosphere. Molecular nitrogen has a partial pressure of 0.78 atm in dry air, O2 of 0.21 atm. Everything else is 0.1 atmosphere combined. The concentration of gases can be expressed as ppm or part per million (1 part in 10-6 of an atmosphere). If the concentration of NO in a location is 10-4 ppm, its concentration in atmospheres = 10-4 ppm x 10-6 atmosphere/ppm = 10-10 atm.

Consider the decomposition of NO into elemental nitrogen and oxygen.



The reaction is favorable as written because the Gibbs free energy change is negative. The equilibrium constant favors products and the reaction should proceed until the concentration of NO is reduced to 10-15.5 atm.


However, the reaction does not proceed as written. This is because the activation energy is high, higher than the energy available at ambient temperature and because of the already low concentration of NO. The probability that 2 molecules of NO will collide with sufficient energy and the correct orientation to form molecular nitrogen and oxygen is very low.

The rate of the reaction is very low, even though it is favored energetically.


Now consider the dimerization of NO2 to form N2O4. In this case, the enthalpy of the reaction is small and negative and the Gibbs free energy is strongly influenced by entropy and temperature. Entropy favors the reactants (2 molecules, less ordered state) over reactants.



At low temperature, the TS term is not very important and the negative enthapy term means that the reaction proceeds as written. Here the activation energy is very small. This is typical of radical coupling reactions.

At higher temperature, the Gibbs free energy is dominated by the entropy term and the back reaction is favored. This reaction also has a small activation energy because the N-N bond that is broken in this step is very weak.



Professor Patricia Shapley, University of Illinois, 2010