Atomic chlorine, Cl, is a free radical and an efficient X catalyst. As we shall see in detail in the next chapter, its concentration in the stratosphere was greatly increased over the twentieth century by the release at the Earth’s sur- face of synthetic chlorine-containing gases. These gases were commercially produced in great volume because they are efficient, nonflammable refrigerants and propellants. However, they are so stable that they eventually rise from ground level to the stratosphere where they decompose, yielding atoms of chlorine.
The unintended consequence of this catalytic destruction process is the ozone hole—the massive destruction of ozone that now occurs annually above the South Pole. Thus the stratosphere, even though it lies far above the Earth’s surface, has not escaped our ecological footprint. The process that produces the ozone hole, detailed here and in Chapter 2, is highly complex, and requires a systems approach for its analysis. In recent times, scientists and engineers have attempted to better anticipate the long-range environmental consequences of new products and processes, having learned their lesson from the ozone hole and other environmental disasters.
However, synthetic gases are not the only suppliers of chlorine to the ozone layer. There always has been some chlorine in the stratosphere as a result of the slow upward migration of the methyl chloride gas, CH3Cl (also called chloromethane), produced at the Earth’s surface, mainly in the oceans by the interaction of chloride ion with decaying vegetation. Recently another large source of methyl chloride, from tropical plants, has been dis- covered; this may be the missing source of the compound for which scientists had been searching.
Only a portion of the methyl chloride molecules are destroyed in the troposphere. When intact molecules of it reach the stratosphere, they are
As discussed in Table 0 -1 on pages xxiii–xxiv, this systems- approach emphasis on environmental, economic, and social consequences is known as the triple bottom line.
photochemically decomposed by UV-C or attacked by OH radicals. In either case, atomic chlorine, Cl, is eventually produced:
CH3Cl 9:Cl CH3 or
OH CH3Cl 9: Cl other products
Chlorine atoms are efficient X catalysts for ozone destruction by Mechanism I:
Cl O39: ClO O2
ClO O 9: Cl O2 overall O3 O 9: 2 O2
Each chlorine atom can catalytically destroy many tens of thousands of ozone molecules in this manner. At any given time, however, the great majority of stratospheric chlorine normally exists not as Cl or as the free radical chlorine monoxide, ClO, but as a form that is not a free radical and that is inactive as a catalyst for ozone destruction. The two main catalytically inactive (or reservoir) molecules containing chlorine in the stratosphere are hydrogen chloride gas, HCl, and chlorine nitrate gas, ClONO2.
The chlorine nitrate is formed by the combination of chlorine monoxide and nitrogen dioxide; after a few days or hours, a given ClONO2 molecule is photochemically decomposed back to its components, and thus the catalyti- cally active ClO is re-formed.
ClO NO2!1 ClONO2
However, under normal circumstances, more chlorine exists at steady state as ClONO2 than as ClO.
The other catalytically inactive form of chlorine, HCl, is formed when atomic chlorine abstracts a hydrogen atom from a molecule of stratospheric methane:
Cl CH49: HCl CH3
This reaction is slightly endothermic, so its activation energy is nonzero, and it therefore proceeds at a slow but significant rate (see Box 1-1). (The methyl free radical, CH3, does not operate like the X catalysts since it combines with an oxygen molecule and is finally degraded to carbon dioxide by reactions discussed in Chapter 3.) Eventually, each HCl molecule is reconverted to the active form, i.e., atomic chlorine, by reaction with the hydroxyl radical:
OH HCl 9: H2O Cl
Again, usually much more chlorine exists as HCl than as atomic chlorine at any given time under normal steady-state conditions.
UV-C
sunlight
0!
Catalytic Processes of Ozone Destruction 29
Cl O3 ClO O2
O
HCl
CH4 OH
ClONO2
NO2 light
When the first predictions concerning stratospheric ozone depletion were made in the 1970s, it was not realized that about 99% of stratospheric chlorine usually is tied up in the inactive forms. When the existence of inac- tive chlorine was discovered in the early 1980s, the predicted amounts of stratospheric ozone loss in the future were lowered appreciably. As we shall see, however, there are conditions under which inactive chlorine can become temporarily activated and massively destroy ozone, a discovery which was not made until the late 1980s.
Although there has always been some chlorine in the stratosphere due to the natural release of CH3Cl from the surface, in recent decades the amount has been completely overshadowed by much larger quantities of chlorine released into air during the production or use of synthetic chlorine- containing gaseous compounds. Most of these substances are chlorofluorocar- bons (CFCs); their nature, usage, and replacements for them will be discussed in detail in Chapter 2.
As with methyl chloride, large quantities of methyl bromide, CH3Br, are also produced naturally and some of it eventually reaches the stratosphere, where it is decomposed photochemically to yield atomic bromine. Like chlo- rine, bromine atoms can catalytically destroy ozone by Mechanism I:
Br O39: BrO O2
BrO O 9: Br O2
In contrast to chlorine, almost all the bromine in the stratosphere remains in the active free-radical forms Br and BrO, since the inactive forms, hydrogen bromide, HBr, and bromine nitrate, BrONO2, are efficiently decomposed photochemically by sunlight. In addition, the formation of HBr from the attack of atomic bromine on methane is a slower reaction than the analogous process involving atomic chlorine, since it is much more endo- thermic and therefore has a higher activation energy:
Br CH49: HBr CH3
A lower percentage of stratospheric bromine exists in inactive form than does chlorine because of the slower speed of this reaction and because of the efficiency of the photochemical decomposition reactions. For that reason, stratospheric bromine is more efficient at destroying ozone than is chlo- rine (by a factor of 40 to 50), but there is much less of it in the strato- sphere, so overall it is less important.
Reaction summary
The Steady-State Analysis of Atmospheric Reactions BOX 1-3
The Steady-State Approximation
If we know the nature of the creation and destruction reaction steps for a reactive sub- stance, we can sometimes algebraically derive a useful equation for its steady-state concentration.
As a simple example, consider the forma- tion and destruction of oxygen atoms above the stratosphere. As mentioned before, the atoms are formed by the photochemical dis- sociation of molecules of diatomic oxygen:
O29: 2 O (i)
The atoms re-form diatomic oxygen when two of them collide simultaneously with a third molecule, M, which can carry away most of the energy released by the newly formed O2 molecule:
O O M 9: O2 M (ii) Recall from introductory chemistry that the rates of the individual steps in reaction mechanisms can be calculated from the con- centrations of the reactants and the rate constant, k, for the step. Thus the rate of reac- tion (i) equals ki [O2]. The rate constant ki here incorporates the intensity of the light impinging upon the molecular oxygen. Thus since two O atoms are formed for each O2 molecule that dissociates,
rate of formation of O atoms 2 ki[O2] The rate of destruction of oxygen atoms by reaction (ii) is
rate of destruction of O atoms 2 kii[O]2[M]
where we square the oxygen atom concentra- tion because two of them are involved as reactants in the step.
The net rate of change of O atom concen- tration with time equals the rate of its forma- tion minus the rate of its destruction:
rate of change of [O] 2 ki[O2] 2 kii[O]2[M]
When atomic oxygen is at a steady state, this net rate must be zero, and thus the right-hand side of the equation above must also be zero.
As a consequence, it follows that kii[O]2[M] ki[O2]
By rearrangement of this equation, we obtain a relationship between the steady-state con- centrations of O and of O2:
[O]ss2/[O2]ss ki/(kii[M])
We see now why the ratio of oxygen atoms to diatomic molecules increases as we go higher and higher above the stratosphere: it is because the air pressure drops, and therefore so does [M], so the O2 re-formation rate decreases.
When molecules such as HCl and HBr eventually diffuse from the strato- sphere back into the upper troposphere, they dissolve in water droplets and are subsequently carried to lower altitudes and are transported to the ground by rain. Thus, although the lifetime of chlorine and bromine in the strato- sphere is long, it is not infinite and the catalysts are eventually removed.
However, the average chlorine atom destroys about 10,000 molecules of ozone before it is removed!
Review Questions 12–16 are based upon material in the preceding section.
Catalytic Processes of Ozone Destruction 31
Steady-State Analysis of the Chapman Mechanism
After this introduction, we now are ready to apply the steady-state analysis to the Chap- man mechanism described by Figure 1-8. The four reactions of concern are shown again below. Notice that the recombination of O atoms, i.e., reaction (ii) above, is not includ- ed because its rate in the mid- and low-strato- sphere is not competitive with other reac- tions, since the oxygen atom concentration is small there.
O29: 2 O (1)
O O2 M 9: O3 M (2) O39: O2 O (3) O3 O 9: 2 O2 (4) Noting that O is produced or consumed in all four reactions, we obtain four terms in its overall rate expression and assume it is in a steady state:
rate of change of [O] 2 rate1 rate2 rate3 rate4 0 (A) Other useful information about concentra- tions can be obtained by considering the steady- state expression for the ozone concentration:
rate of change of [O3] rate2 rate3 rate4
0 (B) If we add together the expressions for the rates of change in [O] and in [O3], i.e., equa- tions (A) and (B) above, we find that the rates for reactions 2 and 3 cancel, and we obtain
2 rate1 2 rate4 0
Using the expressions for these two rates in terms of reactant concentrations, we find
2 k1[O2] 2 k4[O3][O] 0 or
[O3][O] k1[O2]/k4 (C) Another useful expression can be obtained by subtracting equation (B) from (A). We obtain
2 rate1 2 rate2 2 rate3 0
which by rearrangement and cancellation becomes
rate3 rate2 rate1
It is known from experiment that rate2 (and rate3) are much larger than rate1, so the latter can be neglected here, giving simply
rate3 rate2
Using the expressions for these two reaction rates in terms of the concentrations of their reactants,
k3[O3] k2[O] [O2][M]
Rearranging this equation, we can solve for the ratio of ozone to atomic oxygen:
[O3]/[O] k2[O2][M]/k3 (D) Equations (C) and (D) give us two equations in the two unknowns, [O] and [O3]. Multiply- ing their left sides together and equating the result to the product of their right sides elimi- nates [O] and leaves us with an equation for the ozone concentration:
[O3]2 [O2]2[M]k1k2/k3k4
or, taking the square root of both sides, we obtain an expression for the steady-state con- centration of ozone in terms of the diatomic oxygen concentration:
[O3]ss/[O2]ss [M]0.5(k1k2/k3k4)0.5 (E) (continued on p. 32)
Thus the steady-state ratio of ozone to diatom- ic oxygen depends on the square root of the air density through [M]. The ratio is also pro- portional to the square root of the product of the rate constants for the reactions, 1 and 2, in which atomic oxygen and then ozone are produced, and inversely proportional to the square root of the product of the ozone destruction reaction rate constants. Substitu- tion of numerical values for the rate constants k and for [M] into equation (E) predicts the correct order of magnitude for the ozone/
diatomic oxygen ratio, i.e., about 104 in the mid-stratosphere. Ozone never is the main oxygen-containing species in the atmosphere, not even in “the ozone layer.”
Equation (E) predicts that the concentra- tion of ozone relative to that of diatomic oxygen should fall slowly as we climb in the atmosphere, given that it is proportional to the square root of the air density, through the [M] dependence. This occurs because the for- mation reaction of ozone, through step 2, will slow down as [M] declines. This decline with increasing altitude is observed in the upper stratosphere and above. Below about 35 km, however, the more important change in the terms of equation (E) involves k1, and conse- quently the [O3]/[O2] ratio is not simply pro- portional to [M]0.5.
The rate constant k1 incorporates the intensity of sunlight capable of dissociating diatomic oxygen into its atoms. Since the UV-C sunlight required ( 242 nm) is successively filtered by absorption as the light beam descends toward the Earth’s sur- face, the value of k1 declines especially rap- idly in the low stratosphere and below. Thus the concentration of ozone predicted by applying the steady-state analysis to the Chapman mechanism successfully predicts that the ozone concentration will peak in the stratosphere. However, as discussed
above, the actual peak of ozone concentra- tion (⬃25 km, above the equator) occurs rather lower in the stratosphere than the altitude of maximum production (⬃40 km) because horizontal air movement transports ozone downward.
Substitution of equation (E) into (C) allows us to deduce an expression for the steady-state concentration of free oxygen atoms:
[O]ss (k1k3/k2k4)0.5/[M]0.5
Thus the concentration of atomic oxygen is predicted to increase with altitude, as [M]
declines—as in our previous analysis for the upper atmosphere—and as k1 and k3 increase, since UV light intensity increases with increasing altitude. Indeed, atomic oxygen dominates over ozone at high altitudes, whereas below about 50 km, ozone is always dominant.
The production of ozone through reaction (2) is critically dependent upon the supply of free oxygen atoms in reaction (1). The rate of oxygen atom production, in turn, is highly dependent upon the intensity of UV-C sun- light. As we have noted, this intensity falls sharply as we descend through the strato- sphere. The UV-C light intensity also depends strongly upon latitude, being strongest over the equator and declining continuously toward the poles. Thus ozone production is greatest over the equator.
The qualitative behavior of the variation of ozone concentration with altitude predicted by equation (E) is correct, but the predicted amounts of ozone exceed the observed—by about a factor of two near the peak concentra- tion. Scientists eventually found that they had underestimated the rate of the ozone destruc- tion reaction (4) by about a factor of four, since there are catalysts in the stratosphere that greatly speed up the overall reaction.
33
Review Questions
Review Questions
Test your knowledge of some of the factual information in this chapter. If the answer to a question is not obvious to you, use the Index to find the subtopic involved and review that material.
1. Which three gases constitute most of the Earth’s atmosphere?
2. What range of altitudes constitutes the troposphere? the stratosphere?
3. What is the wavelength range for visible light?
Does ultraviolet light have shorter or longer wavelengths than visible light?
4. Which atmospheric gas is primarily responsible for filtering sunlight in the 120⫺220-nm region?
Which, if any, gas absorbs most of the Sun’s rays in the 220⫺320-nm region? Which absorbs primarily in the 320⫺400-nm region?
5. What is the name given to the finite packets of light absorbed by matter?
6. What are the equations relating photon energy E to light’s frequency and wavelength ?
7. What is meant by the expression photochemically dissociated as applied to stratospheric O2?
8. Write the equation for the chemical reaction by which ozone is formed in the stratosphere.
What are the sources for the different forms of oxygen used here as reactants?
9. Write the two reactions that, aside from the catalyzed reactions, contribute most significantly to ozone destruction in the stratosphere.
10. What is meant by the phrase excited state as applied to an atom or molecule? Symbolically, how is an excited state signified?
PROBLEM 1
Consider the following 3-step mechanism for the production and destruction of excited oxygen atoms, O*, in the atmosphere:
O29: O ⫹ O*
O* ⫹ M 9: O ⫹ M O* ⫹ H2O 9: 2 OH
Develop an expression for the steady-state concentration of O* in terms of the concen- trations of the other chemicals involved.
PROBLEM 2
Perform a steady-state analysis for d(Cl]/dt and for d[ClO]/dt in the following mechanism:
Cl29: 2 Cl (1)
Cl ⫹ O39: ClO ⫹ O2 (2) 2 ClO 9: 2 Cl ⫹ O2 (3) ClO ⫹ NO29: ClONO2 (4) Obtain expressions for the steady-state con- centrations of Cl and ClO, and hence for the rate of destruction of ozone.
PROBLEM 3
Perform a steady-state analysis on the 3-step reaction mechanism below. Assume that both ozone and atomic oxygen are in a steady state, and derive an expression for the ratio [NO2]/[NO].
NO29: NO ⫹ O
O ⫹ O29: O3
NO ⫹ O39: NO2 ⫹ O2
PR PR PR
PRROBOBOBOBO LELELELEM M M MM1
light
PR PR PR
PRROBOBOBOBO LELELELEM M MMM2
PR PR PR
PRROBOBOBOBLELELELEM M M MM3
Additional Problems
The problems given within the chapter, and the more elaborate ones given here, are designed to test your problem-solving abilities.
1. A possible additional mechanism that could exist for the creation of ozone in the high stratosphere begins with the creation of
(vibrationally) excited O2 and ground-state atomic oxygen from the absorption of photons with wavelengths less than 243 nm. The O2* reacts with a ground-state O2 molecule to produce ozone and another atom of oxygen. What is the net reaction from these two steps? What do you predict is the fate of the two oxygen atoms, and what would be the overall reaction once this fate is included?
2. In the nonpolluted atmosphere, an important mechanism for ozone destruction in the lower stratosphere is
OH O39: HOO O2 HOO O39: OH 2 O
2
Does this pair of steps correspond to Mechanism I?
If not, what is the overall reaction?
3. A proposed mechanism for ozone destruction in the late spring over northern latitudes in the lower stratosphere begins with the photochemical decomposition of ClONO2 to Cl and NO3, followed by photochemical decomposition of the latter to NO and O2. Deduce a catalytic ozone destruction cycle, requiring no atomic oxygen,
that incorporates these reactions. What is the overall reaction?
4. Deduce possible reaction step(s), none of which involve photolysis, for Mechanism II following the X O39: XO O2 step such that the sum of all the mechanism’s steps does not destroy or create any ozone.
5. As will be discussed in Chapter 2, atomic chlorine is produced under ozone-hole conditions by the dissociation of diatomic chlorine, Cl2. Given that diatomic chlorine gas is the stablest form of the element, and that the Hf° value for atomic chlorine is 121.7 kJ mol1, calculate the maximum wavelength of light that can dissociate diatomic chlorine into the monatomic form. Does such a wavelength correspond to light in the visible or the UV-A or the UV-B region?
6. Under conditions of low oxygen atom
concentration, the radical HOO can react reversibly with NO2 to produce a molecule of HOONO2:
HOO NO29: HOONO2
(a) Deduce why the addition of nitrogen oxides to the lower stratosphere could lead to an increase in the steady-state ozone concentration as a consequence of this reaction.
(b) Deduce how the addition of nitrogen oxides to the middle and upper stratosphere could decrease the ozone concentration there as a consequence of other reactions.
11. Explain why the phrase ozone layer is a misnomer.
12. Define the term free radical, and give two examples relevant to stratospheric chemistry.
13. What are the two steps, and the overall reaction, by which X species such as ClO catalytically destroy ozone in the middle and upper stratosphere via Mechanism I?
14. What is meant by the term steady state as applied to the concentration of ozone in the stratosphere?
15. Explain why, atom for atom, stratospheric bromine destroys more ozone than does chlorine.
16. Explain why ozone destruction via the reaction of O3 with atomic oxygen does not occur to a significant effect in the lower stratosphere.