When healthy, Earth’s stratosphere contains a low concentration of ozone (O₃) that absorbs potentially harmful ­ultraviolet (UV) radiation by the cycle shown at right.
Chlorofluorocarbon refrigerants, such as Freon 12 (CF₂Cl₂), are stable in the lower atmosphere, but in the stratosphere they absorb high-energy UV radiation to generate chlorine radicals.


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?

Verified step by step guidance

1. Step 1: Understand the problem. The question involves the destruction of ozone (O3) in the stratosphere by chlorine radicals (Cl⋅) generated from chlorofluorocarbons (CFCs). The goal is to propose two mechanisms that explain how a small number of chlorine radicals can destroy a large number of ozone molecules and determine which mechanism is more likely when the concentration of chlorine atoms is very small.
2. Step 2: Analyze the given reactions. The reactions provided show how chlorine radicals interact with ozone and other species in the stratosphere. Key reactions include: (1) Cl⋅ + O3 → Cl—O⋅ + O2, (2) Cl—O⋅ + O → Cl⋅ + O2, and (3) 2 Cl—O⋅ → Cl—O—O—Cl. These reactions suggest a catalytic cycle where chlorine radicals are regenerated, allowing them to destroy multiple ozone molecules.
3. Step 3: Propose Mechanism 1. In this mechanism, chlorine radicals (Cl⋅) act as a catalyst in a two-step cycle: (1) Cl⋅ reacts with ozone (O3) to form Cl—O⋅ and O2, and (2) Cl—O⋅ reacts with an oxygen atom (O) to regenerate Cl⋅ and produce another O2 molecule. This cycle can repeat, allowing a single Cl⋅ to destroy multiple ozone molecules.
4. Step 4: Propose Mechanism 2. In this mechanism, chlorine radicals form a dimer intermediate (Cl—O—O—Cl) through the reaction 2 Cl—O⋅ → Cl—O—O—Cl. This dimer can then decompose under UV light (hv) to regenerate 2 Cl⋅ and release O2. The regenerated Cl⋅ can then continue to destroy ozone molecules, making this another catalytic cycle.
5. Step 5: Determine the more likely mechanism at low chlorine concentrations. Mechanism 1 is more likely when the concentration of chlorine atoms is very small because it involves a simple two-step cycle that does not require the formation of a dimer intermediate. Mechanism 2, which involves the formation of Cl—O—O—Cl, is less likely at low concentrations due to the lower probability of two Cl—O⋅ radicals colliding to form the dimer.

Key Concepts

Ozone Depletion Mechanism

Ozone depletion occurs when ozone (O3) molecules are broken down by reactive species, such as chlorine radicals (Cl·). These radicals can initiate a chain reaction, where one chlorine atom can destroy thousands of ozone molecules. Understanding this mechanism is crucial for analyzing how a small concentration of chlorine can lead to significant ozone loss in the stratosphere.

Exothermic Reactions

Exothermic reactions release energy, often in the form of heat, during the process of breaking and forming chemical bonds. In the context of ozone depletion, the reactions involving chlorine radicals and ozone are exothermic, which means they can proceed rapidly and contribute to the efficiency of ozone destruction. Recognizing the nature of these reactions helps in understanding their impact on atmospheric chemistry.

Chain Reactions

Chain reactions are processes where the products of a reaction initiate further reactions, leading to a rapid increase in the number of reactants consumed. In the case of chlorine radicals and ozone, one chlorine radical can react with an ozone molecule to produce more radicals, perpetuating the cycle. This concept is essential for explaining why even a small number of chlorine radicals can have a large effect on ozone concentrations.