When healthy, Earth's stratosphere contains a low...

When healthy, Earth's stratosphere contains a low concentration of ozone $\left(\mathrm{O}_{3}\right)$ that absorbs potentially harmful ultraviolet (UV) radiation by the cycle shown at right. Chlorofluorocarbon refrigerants, such as Freon $12\left(\mathrm{CF}_{2} \mathrm{Cl}_{2}\right)$, are stable in the lower atmosphere, but in the stratosphere, they absorb high-energy UV radiation to generate chlorine radicals. $$ \mathrm{CF}_{2} \mathrm{Cl}_{2} \stackrel{h \nu}{\longrightarrow} \cdot \mathrm{CF}_{2} \mathrm{Cl}+\mathrm{Cl} $$ The presence of a small number of chlorine radicals appears to lower ozone concentrations dramatically. The following reactions are all known to be exothermic (except the one requiring light) and to have high rate constants. Propose two mechanisms to explain how a small number of chlorine radicals can destroy large numbers of ozone molecules. Which of the two mechanisms is more likely when the concentration of chlorine atoms is very small? $$ \begin{aligned} &\mathrm{Cl}-\mathrm{O}-\mathrm{O}-\mathrm{Cl} \stackrel{h \nu}{\longrightarrow} \mathrm{O}_{2}+2 \mathrm{Cl} \cdot \mathrm{Cl}-\mathrm{O} \cdot+\mathrm{O} \longrightarrow \mathrm{O}_{2}+\mathrm{Cl} \cdot \\ &\mathrm{Cl} \cdot+\mathrm{O}_{3} \longrightarrow \mathrm{Cl}-\mathrm{O} \cdot+\mathrm{O}_{2} \quad 2 \mathrm{Cl}-\mathrm{O} \cdot \longrightarrow \mathrm{Cl}-\mathrm{O}-\mathrm{O}-\mathrm{Cl} \end{aligned} $$

When healthy, Earth's stratosphere contains a low concentration of ozone $\left(\mathrm{O}_{3}\right)$ that absorbs potentially harmful ultraviolet (UV) radiation by the cycle shown at right.
Chlorofluorocarbon refrigerants, such as Freon $12\left(\mathrm{CF}_{2} \mathrm{Cl}_{2}\right)$, are stable in the lower atmosphere, but in the stratosphere, they absorb high-energy UV radiation to generate chlorine radicals.
$$
\mathrm{CF}_{2} \mathrm{Cl}_{2} \stackrel{h \nu}{\longrightarrow} \cdot \mathrm{CF}_{2} \mathrm{Cl}+\mathrm{Cl}
$$
The presence of a small number of chlorine radicals appears to lower ozone concentrations dramatically. The following reactions are all known to be exothermic (except the one requiring light) and to have high rate constants. Propose two mechanisms to explain how a small number of chlorine radicals can destroy large numbers of ozone molecules. Which of the two mechanisms is more likely when the concentration of chlorine atoms is very small?
$$
\begin{aligned}
&\mathrm{Cl}-\mathrm{O}-\mathrm{O}-\mathrm{Cl} \stackrel{h \nu}{\longrightarrow} \mathrm{O}_{2}+2 \mathrm{Cl} \cdot \mathrm{Cl}-\mathrm{O} \cdot+\mathrm{O} \longrightarrow \mathrm{O}_{2}+\mathrm{Cl} \cdot \\
&\mathrm{Cl} \cdot+\mathrm{O}_{3} \longrightarrow \mathrm{Cl}-\mathrm{O} \cdot+\mathrm{O}_{2} \quad 2 \mathrm{Cl}-\mathrm{O} \cdot \longrightarrow \mathrm{Cl}-\mathrm{O}-\mathrm{O}-\mathrm{Cl}
\end{aligned}
$$

Key Concepts

Catalytic Cycles

Catalytic cycles are reaction mechanisms in which a catalyst—in this case, reactive radicals—is regenerated in subsequent steps so that a single catalyst molecule can facilitate the transformation of many substrate molecules. Within atmospheric chemistry, such cycles allow trace species like chlorine radicals to repeatedly engage in reactions that deplete ozone, amplifying their overall effect despite their low concentrations.

Chain Reaction Mechanisms

Chain reactions involve a series of steps where an initial reactive species generates further reactive intermediates that continue to propagate the reaction cycle. In the context of ozone destruction, the formation and subsequent reactions of radicals create cycles of propagation and regeneration, enabling a single radical to convert many ozone molecules into other products before termination occurs.

Free Radical Chemistry

Free radicals are highly reactive species due to their unpaired electrons, making them effective participants in atmospheric reactions. Their ability to rapidly react with ozone and other species underpins both the initiation and the propagation steps in the catalytic chain mechanisms. This concept is central to understanding how small concentrations of a radical can lead to significant changes in atmospheric composition.

Photochemical Activation

Photochemical activation refers to the process by which molecules absorb photons, typically from UV radiation, to become chemically excited and form reactive intermediates such as radicals. In the stratosphere, this process is critical for generating the chlorine radicals from stable compounds, thereby initiating the chain reactions responsible for ozone depletion.

Concentration Effects in Kinetic Mechanisms

Kinetic mechanisms are heavily influenced by the concentration of the reactive species involved. When radical concentrations are very low, bimolecular interactions that lead to radical termination (such as dimer formation) become less likely, favoring chain propagation pathways where radicals react with abundant reactants like ozone. This principle explains why certain catalytic cycles are more effective in environments with low concentrations of active radicals.

Transcript

0:00 Okay.

00:02 Recent years, ozone in the stratosphere has been depleted at very fast rates because of chloro -florocarbon decompositions.

00:11 The chloro -floric carbon, such as cfcl3, is first decomposed by this reaction, then the chlorine radical from here undergoes this reaction to deplete the ozone.

00:26 We're asked to rate the overall reaction for the last two steps.

00:30 So you can see this disappears and this disappears.

00:37 End up with 03 plus o yields, whoops, and i could write 202, but i'm going to write 02 plus o2.

00:55 Next, what are the roles of cl and clo? cl is a catalyst, and this is because it's present as an initial reactant and as a final product.

01:44 The clo is an intermediate because it's present, as first a reactant, or first a product, then a reactant.

02:25 Why is the fluorine radical not important to this mechanism? and this is, i don't know if the answer is actually so simple.

02:59 Okay, let me see my reactions here.

03:05 Okay, so the cf bond is much stronger than the c -l bond.

03:20 And that is because fluorine has a higher electronegativity than chlorine.

03:44 Thus, the cl radical will form, but the fluorine stays bonded to the carbon.

04:08 And finally, d, i think it's finally, i don't know, i shouldn't say finally.

04:16 One suggestion is to add hydrocarbons to ethane.

04:20 Will this work? to reduce the concentration of the chlorine radicals is to add hydrocarbons, such as ethane to the stratosphere.

04:57 Will this work? how will this work? if we add for, let's see what happens in this reaction.

05:08 Let me write it down.

05:10 See if i can think of what's going to happen here.

05:13 Okay, so that will bond.

05:15 It'll create a hcl plus c2h5, which you know just gives us like more acid up there.

05:24 That's probably not too fabulous.

05:30 And then we're going to draw potential energy versus reactionist progress for the uncatalized 030 plus two or yields 202.

06:05 And then we're going to use thermodynamic data to figure out if this is my uncatalized.

06:17 Let me read this whole thing...

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