Red-orange NO2 absorbs sunlight(Figure 1) and produces ozone by the sequence of reactions shown below.

Figure 1. The absorption spectrum of NO2, a component of photochemical smog.

          hn (l < 410 nm)
         NO2 ---> NO + O
          O + O2 ---> O3

Contrary to the impression given in many introductory chemistry texts, these reactions alone cannot explain photochemical smog. In the virtual experiment, it was found that measurement of the rate constant for the gas phase reaction:

          2NO + O2 ---> 2NO2

obeys the rate law, rate = k[NO]²[O2], and the rate constant k_298K = 2 x 10^-38 cm⁶/molec² s.

At low [NO], the reaction becomes very slow. The half-lives calculated under pseudo second order conditions ([O2] = 4.92 x 10¹⁸ molec/cm³) at atmospheric pressure are as follows:

The reaction between oxygen and NO is so slow under ambient atmospheric conditions ([NO] < 50 ppbv) that oxidation by ozone dominates. If this occurs there is no net production of ozone.

         NO + O3 ---> NO2 + O2

Other oxidants present in the unpolluted troposphere (OH, OOH) can oxidize NO (e.g., NO + OOH ---> NO2 + OH). In the absence of volatile organic compounds, OH and OOH primarily derive from photochemical reactions of ozone (e.g., O3 + sunlight ---> O2 + O and then O + H2O ---> 2 OH). When NO is reoxidized by ozone (or an oxidant immediately derived from it), there is no net ozone production. Only when NO is reoxidized with oxidants other than ozone or its immediate products is there a net increase in tropospheric ozone levels.

     The second key to urban ozone production is the presence of volatile organic compounds (VOCs) and carbon monoxide. These compounds are byproducts of incomplete combustion, and VOCs also are produced by the evaporation or leakage of fuels. In addition, resinous trees, such as pine, emit significant quantities of natural VOCs, called terpenes. This exacerbates the pollution problem in areas such as the southeastern U.S. Oxidants for NO in urban smog include peroxy organic radicals, formed by the oxidation of carbon monoxide and VOCs. The chemistry begins by forming an organic free radical with the potent trace atmospheric oxidant, OH. Methane is the best understood hydrocarbon system:

     CH4 + OH ---> CH3 + H2O

     CH3 + O2 ---> CH3OO

     CH3OO + NO ---> CH3O + NO2

     CH3O + O2 ---> HCHO + HO2

       hn (l <330 nm)
     HCHO ---> HCO + H

     HCO + O2 ---> CO + HOO

      H + O2 ---> HOO

     CO + OH ---> H + CO2 (note how CO oxidation can enter the cycle)

     H + O2 ---> HOO

Each peroxy radical produced in the above sequence (shown in boldface) can reoxidize NO to NO2, (shown in the equation below) while forming additional OH to maintain the radical chain reactions. The peroxy species in this scheme are formed directly from O2 instead of from ozone.

     NO + HOO ---> NO2 + OH

     It is the synergistic effect of NOx, CO, and volatile hydrocarbons that produces photochemical smog. That is why automobile emission control regulations mandate the reduction of carbon monoxide and VOC emissions, as well as NOx. One approach to meeting these regulations is the three-way automotive catalyst. It mitigates carbon monoxide and VOC emissions (by their catalytic oxidation to CO2 and H2O). The NOx emissions are controlled by their catalytic reduction to N2. The difficult task of simultaneously reducing NO and oxidizing the VOC and CO exhaust components still poses problems in the technology for catalytic converters. Present automotive catalysts do not solve the NOx reduction problem well. As a result, automobiles are deliberately run in a fuel rich mode to provide excess reducing agent in the exhaust stream for NOx; however, this results in a 10% decrease in fuel efficiency. Excess CO and VOCs are ultimately removed by catalytic oxidation.

   The time evolution of urban smog can also be easily understood based on the above chemistry. First NO appears as an initial combustion product, as the morning rush hour begins. Oxidation to NO2 then occurs, followed by photochemical production of ozone. Ozone levels do not become highly elevated until afternoon. When the sun sets and transportation and photochemical activity diminish, the reactive gases (ozone and NO_x)react or are removed by other deposition processes. For example, NO2 reacts with moisture and oxidants to yield nitric acid aerosols, which deposit as "acid rain." Nighttime chemical reactions may leave an aerosol residue in the air, but ozone levels return to normal by midnight. Visit the South Coast Air Quality Management District for daily plots of smog components in the Los Angeles Basin.

Stratospheric Chemistry of Nitrogen Oxides. It is ironic that NOx catalyzes formation of toxic ozone in the troposphere, but helps destroy the protective stratospheric ozone layer. Because NO and NO2 are ultimately oxidized to HNO3 and deposit within several days, they do not enter the stratosphere directly. It takes 1-2 years for diffusion and convection to transport gases to the stratosphere from sea level. Only gases with atmospheric lifetimes of many months survive the transport time. Nitrous oxide (t1/2 = 120 yrs) is such as gas.

   High energy radiation in the stratosphere causes harmless photodissociation of N2O. The oxygen atom is formed in an electronically excited state denoted ¹D .

       hn (l <330 nm)
     N=N=O ---> N2 + O(1D)

The O(¹D) present in the stratosphere, produced mainly from photolysis of O3, reacts as follows:

      N2O + O(¹D) ---> 2NO (62%)
                             ---> N2 + O2 (38%)

A free radical chain reaction catalytically destroys ozone by a mechanism similar to that of chlorine radicals produced by the photodestruction of CFCs.

     NO + O3 ---> NO2 + O2 (k_220K = 3.5 x 10^-15 cm³/molec sec)
      NO2 + O ---> NO + O2 (k_220K = 9.3 x 10^-12 cm³/molec sec)
__________________________________________________
      O3 + O ---> 2O2

Catalysis is important because the rate constant for the direct reaction between O and O3 (k_220K = 6.8 x 10^-16 cm³/molec sec) is slower than that for the two step catalytic process. Paul Crutzen shared the 1995 Nobel Prize in Chemistry largely for his work showing the role of nitrogen oxides in stratospheric ozone depletion. Supersonic jets fly at high altitudes to limit air friction, and inject NO from combustion exhaust directly into the lower stratosphere. For this reason, their use has been limited. An increase in the atmospheric concentration of nitrous oxide poses a similar threat to the stratosphere.

References:

1)Graedel, T. E.; Crutzen, P. J. Atmospheric Change : an Earth System Perspective; W.H. Freeman: New York, 1993.
2)Silver, R. G.; Sawyer, J. E.; Summers, J. C. Catalytic Control of Air Pollution : Mobile and Stationary Sources; American Chemical Society: Washington, 1992.
3)Gray, H. B.; Simon, J. D.; Trogler, W. C. Braving the Elements; University Science Books: Sausalito, 1995.
4)Wayne, R. P. Chemistry of Atmospheres; Second ed.; Oxford University Press: Oxford, 1991.
5) View today's stratospheric ozone levels from a NASA satellite. Note that the Dobson Unit (DU) used to measure ozone is the thickness in hundredths of a millimeter of the column of ozone above the Earth if it were compressed to 1 atm pressure at 0^o C. Thus, 300 DU corresponds to an ozone layer only 3 mm thick.

[Table of Contents]

[Introduction]