One of the concerns about the use of Freons is that they will migrate to the upper atmosphere, where chlorine atoms can be generated by the following reaction: $$ \mathrm{CCl}_{2} \mathrm{~F}_{2}(g) \stackrel{\mathrm{h}}{\longrightarrow} \mathrm{CF}_{2} \mathrm{Cl}(g)+\mathrm{Cl}(g) $$ on Chlorine atoms can act as a catalyst for the destruction of ozone. The activation energy for the reaction $$ \mathrm{Cl}(g)+\mathrm{O}_{3}(g) \longrightarrow \mathrm{ClO}(g)+\mathrm{O}_{2}(g) $$ is \(2.1 \mathrm{~kJ} / \mathrm{mol}\). Which is the more effective catalyst for the destruction of ozone, \(\mathrm{Cl}\) or \(\mathrm{NO}\) ? (See Exercise 75.)

Activation energy is a critical concept in chemical reactions, including those occurring in atmospheric chemistry. It refers to the minimum amount of energy required for reactants to transform into products during a chemical reaction. Think of it like the initial push you need to get a ball rolling up a hill before it can freely roll down the other side. In the context of ozone depletion, the activation energy determines how easily certain catalytic reactions can occur.

From the textbook exercise, it's clear that both chlorine (Cl) and nitrogen oxide (NO) can act as catalysts that lower the activation energy required for the decomposition of ozone (O₃). Comparing their respective activation energies gives us insight into which catalyst is more effective at facilitating this reaction. Since a lower activation energy means the reaction can occur more readily and at a faster rate, the textbook example shows that NO, with an activation energy of 1.8 kJ/mol, is a more efficient catalyst than Cl, with an activation energy of 2.1 kJ/mol. This subtlety highlights the importance of activation energy in determining reaction efficiency.

Atmospheric chemistry deals with chemical reactions that take place in Earth's atmosphere. It is a complex field because the atmosphere is a dynamic and sensitive system influenced by various factors like solar radiation, temperature, and human activities. The reactions occurring in the atmosphere can lead to the formation or depletion of atmospheric gases, which in turn can have significant impacts on climate and living organisms.

For example, ozone, a molecule composed of three oxygen atoms (O₃), plays a crucial role in protecting life on Earth by absorbing harmful ultraviolet radiation from the sun. However, certain reactions catalyzed by man-made or naturally occurring substances can lead to its destruction, commonly referred to as 'ozone depletion.' This process has serious environmental consequences, such as increased UV radiation reaching the Earth's surface, which can lead to higher rates of skin cancer and cataracts in humans, as well as negative effects on ecosystems.

Freons, a common term for a class of chlorofluorocarbons (CFCs), have been widely used as refrigerants, propellants in aerosol sprays, and in various other applications. However, when they are released into the atmosphere, they can drift upwards to the stratosphere where they are broken down by UV light, releasing chlorine atoms as depicted in the reaction from the textbook exercise: CCl₂F₂(g) → CF₂Cl(g) + Cl(g).

Once free, these chlorine atoms can catalyze reactions that deplete the ozone layer, like the one in which a single Cl atom can convert multiple molecules of ozone into oxygen without being consumed itself. This makes chlorine a potent catalyst for ozone destruction. Freons, therefore, are significant contributors to ozone layer depletion, leading to international regulatory efforts like the Montreal Protocol, which are aimed at phasing out their use and restoring the ozone layer.