From various natural sources—including fires, lightning, anaerobic biological decay, and emissions from volcanoes—our atmosphere regularly receives inputs of many gases including the partially oxidized compounds carbon mon- oxide, CO, nitric oxide, NO, and sulfur dioxide, SO2, and several simple compounds of hydrogen combined with another element in a highly reduced form, such as ammonia, NH3, hydrogen sulfide, H2S, and methane, CH4.
The concentrations of these gases do not build up in clean air because there are not only sources for them but also sinks, which result in their continual destruction. For the gases mentioned above, the destruction pro- cesses are oxidation reactions that occur in air. However, none of the gases reacts directly with diatomic oxygen molecules because the activation energy for such processes is too high. Rather, their reactions begin when they are attacked by the hydroxyl free radical, OH, even though the concentration of this species in air is exceedingly small, a few million molecules per cubic centimeter on average (see Problem 3-1). Its concentration has remained con- stant in air over the last few decades at least.
The presence of an unpaired electron makes most free radicals, including OH, very reactive. The Lewis structure for the hydroxyl free radical is
9H O
In clean tropospheric air, as in the stratosphere, the hydroxyl radical is produced when a small fraction of the excited oxygen atoms resulting from the photochemical decomposition of trace amounts of atmospheric ozone, O3, react with gaseous water to abstract one hydrogen atom from each H2O molecule:
O39: O2 O*
O* H2O 9: 2 OH
The average tropospheric lifetime of a hydroxyl radical is only about one second, since it reacts quickly with one or other of many atmospheric gases.
Recall that a free radical has one electron in the outermost shell of one of its atoms that neither participates in a bond to another atom nor is part of a nonbonding electron pair.
UV-B
See Additional Problem 2 for the lifetime calculation.
PROBLEM 4
The average outdoor concentration of carbon monoxide, CO, is about 1000 g m3. What is
this concentration expressed on the ppm scale? On the molecules cm3 scale? Assume that the outdoor temperature is 17°C and that the total air pressure is 1.04 atm.
PR PR PR
PROBOBOBOBLELELEL M M MM4
Because the lifetime of hydroxyl radicals is short and sunlight is required to form more of them, the OH concentration drops quickly at nightfall.
PROBLEM 3-1
In one study, the concentration of OH in air at the time was found to be 8.7 106 molecules per cubic centimeter. Calculate its molar concentration, and its concentration in parts per trillion, assuming that the total air pressure
is 1.0 atm and the temperature is 15°C. ●
In its reaction with otherwise-stable gases whose molecules contain multiple bonds, OH adds itself to them, thereby forming a larger free radi- cal. For example, hydroxyl adds to carbon monoxide molecules, forming the transient free radical HOCO:
OH CO 9: HOCO
Most collisions of OH and CO molecules are ineffective in promoting a reac- tion. Consequently, the average lifetime of a carbon monoxide molecule in air is a month or two.
Molecular oxygen reacts quickly with transient free radicals such as HOCO once they are formed, thereby involving itself in the oxidation process. In the present case, O2 abstracts a hydrogen atom from the free radi- cal, thereby forming the hydroperoxy free radical, HOO, and the fully oxi- dized product CO2, carbon dioxide:
O2 HOCO 9: HOO CO2
The hydroperoxy radical produced in the atmospheric oxidation of car- bon monoxide, and indeed of most molecules, is in turn converted back to the hydroxyl radical by its oxidation of nitric oxide, NO, which is present in adequate concentration for this purpose in all but the very clean- est air:
HOO NO 9: OH NO2
The general cycle of OH/HOO formation and consumption in the atmo- spheric oxidation of various molecules is summarized by the diagram below;
in the case of some organic molecules, sunlight is required for intermediate steps in the mechanism:
OH HOO
NO
stable gas, O2 (light) PROBLEM3-1
H!O!O
"
O C"O
"
N O O!N"O C#O
H!O!C"O
Introduction 75
Although suspected for decades of playing a pivotal role in air chemistry, the presence of OH in the troposphere was confirmed only relatively recently since its concentration is so very small. The great importance of the hydroxyl radical to tropospheric chemistry arises because it, not O2, initiates the oxi- dation of almost all reduced gases. Without OH and its related reactive spe- cies HOO, most naturally occurring gases, and pollutant gases such as the unburned hydrocarbons emitted from vehicles would not be efficiently removed from the troposphere.
The reactions that OH initiates correspond to a flameless, ambient- temperature “burning” of the reduced gases of the lower atmosphere. If these gases were to accumulate, the atmospheric composition would be quite dif- ferent, as would the forms of life that would be viable on Earth. Interestingly, hydroxyl is unreactive toward molecular oxygen—in contrast to the behavior of O2 with many free radicals—and to molecular nitrogen, thus it survives long enough to react with so many other species.
Within a few minutes, most of the nitrogen dioxide, NO2, produced in the OH/HOO cycle during the daytime absorbs UV-A from sunlight (see its spectrum in Figure 3-1) and photochemically decomposes to nitric oxide and atomic oxygen (this also occurs in the stratosphere, as mentioned in Chapter 1):
NO2 UV-A 9: NO O
From the viewpoint of the nitrogen oxides, the cycle of NO oxidation by HOO and the reduction of NO2 by sunlight are summarized on the next page:
Indeed, OH has been called the “tropospheric vacuum cleaner” or “detergent.”
Light with wavelength shorter than 394 nm has sufficient energy to decompose NO2 by this reaction.
100
10
300 500
Wavelength (nm)
600 1
Absorption cross section (10–20cm2)
400
Visible
UV-B UV-A
Ultraviolet
FIGURE 3-1 Absorption spectrum for gaseous NO2. [Source: J. H. Seinfeld and S. N. Pandis, “Atmospheric Chemistry and Physics,” John Wiley & Sons, New York, 1998.]
O
NO NO2
sunlight UV-A HOO
The oxygen atoms produced in this cycle quickly react with molecular oxy- gen to form ozone. As is the case in the stratosphere, this reaction is the only source of ozone in the troposphere:
O O2!: O3
In summary, stable gases in the air that are not already fully oxidized react directly with OH, rather than O2, even though it is present in tiny concentration. The OH is originally produced from reaction of O* from ozone photodecomposition, the ozone having been created from the oxy- gen atom produced by photochemical decomposition of NO2. After it is used for reaction initiation, the OH is transformed into HOO, which is recycled back to OH by reaction with NO.
Urban Ozone: The Photochemical Smog Process