The gas 'A' is having very low reactivity reaches to stratosphere. It is non-toxic and non-flammable

Q: Why doesn't the gas react in the lower atmosphere (troposphere)?

CFCs are highly stable and have very low reactivity due to strong carbon-halogen bonds. There are no significant sinks (like rainout or oxidation) for them in the troposphere, allowing them to slowly drift upward.

Q: Why does the C-Cl bond break instead of the C-F bond?

The carbon-chlorine ($C-Cl$) bond has a lower bond dissociation enthalpy compared to the carbon-fluorine ($C-F$) bond. Therefore, UV radiation provides enough energy to break the weaker $C-Cl$ bond first.

Q: What is homolytic cleavage?

Homolytic cleavage is a type of bond breaking where each atom involved in the bond takes one electron from the shared pair, resulting in the formation of highly reactive species called free radicals.

Q: Why are free radicals so reactive?

Free radicals have an unpaired electron, making them highly unstable. They desperately seek another electron to pair up, which drives them to react rapidly with other molecules, such as ozone.

Q: How does one chlorine radical destroy so much ozone?

The chlorine radical acts as a catalyst. It reacts with ozone to form chlorine monoxide ($\dot{C}lO$) and oxygen. The $\dot{C}lO$ then reacts with a free oxygen atom to regenerate the chlorine radical, allowing the cycle to repeat thousands of times.

Question

The gas 'A' is having very low reactivity reaches to stratosphere. It is non-toxic and non-flammable but dissociated by UV-radiations in stratosphere. The intermediates formed initially from the gas 'A' are

Options:

(A) $\dot{C}lO + \dot{C}F_2Cl$

(B) $\dot{C}lO + \dot{C}H_3$

(C) $\dot{C}H_3 + \dot{C}F_2Cl$

(D) $\dot{C}l + \dot{C}F_2Cl$

Sigma Percentile · Accepted Answer

### The Mystery of Gas 'A'

The problem introduces us to a mysterious gas 'A' with a very specific set of properties. It is described as non-toxic, non-flammable, and possessing very low reactivity.

Because of this extreme stability, it doesn't react or break down in the lower atmosphere (the troposphere). Instead, it slowly drifts upward, eventually reaching the stratosphere.

These characteristics are the classic hallmarks of **Chlorofluorocarbons (CFCs)**, widely used in the past as refrigerants and aerosol propellants. A common example is Freon-12, which has the chemical formula $CF_2Cl_2$.

### The Stratospheric Journey

Once these CFC molecules reach the stratosphere, the environment changes drastically. They are no longer protected by the Earth's thick lower atmosphere.

Here, they are exposed to intense, high-energy **ultraviolet (UV) radiation** from the Sun. This radiation packs enough energy to initiate chemical reactions that were impossible at ground level.

### The Photochemical Cleavage

When a UV photon strikes a $CF_2Cl_2$ molecule, it targets the weakest link in the structure. The carbon-chlorine ($C-Cl$) bond is significantly weaker than the carbon-fluorine ($C-F$) bond.

The UV energy causes the $C-Cl$ bond to undergo **homolytic cleavage**. This means the bond breaks evenly, with each atom taking one electron from the shared pair.

$$ CF_2Cl_2(g) \xrightarrow{UV} \dot{C}l(g) + \dot{C}F_2Cl(g) $$

This reaction generates two highly reactive species: a **chlorine free radical** ($\dot{C}l$) and a **chlorodifluoromethyl free radical** ($\dot{C}F_2Cl$).

Looking at our options, we can confidently identify that the initial intermediates formed are indeed $\dot{C}l$ and $\dot{C}F_2Cl$. Therefore, **Option (d)** is the correct answer.

### The Chain Reaction of Destruction

While the question ends there, the chemistry certainly does not. The chlorine free radical ($\dot{C}l$) we just formed is the true villain of this story.

It acts as a powerful catalyst, aggressively attacking ozone ($O_3$) molecules in the stratosphere and breaking them down into ordinary oxygen ($O_2$).

$$ \dot{C}l(g) + O_3(g) \rightarrow \dot{C}lO(g) + O_2(g) $$

$$ \dot{C}lO(g) + O(g) \rightarrow \dot{C}l(g) + O_2(g) $$

Notice how the chlorine radical is regenerated at the end of this cycle. Because it is not consumed, a single chlorine atom can destroy up to 100,000 ozone molecules before it is finally removed from the stratosphere!

This devastating chain reaction is the primary cause of **ozone layer depletion**, making it a crucial concept to master for your exams.

Animated Solution for Chemistry - Environmental Chemistry: The gas 'A' is having very low reactivity reaches to stratosphere. It is non-toxic and non-flammable but dissociated by UV-radiations in stratosphere. The intermediates formed initially from the gas 'A' are

Select Answer:

Visualized Solution

Identifying Gas 'A'

Dissociation of CFCs by UV Radiation

Identifying the Initial Intermediates

The Way Forward (Ozone Depletion)

The Sigma Insight: Atmospheric Pollution

The Mystery of Gas 'A'

The Stratospheric Journey

The Photochemical Cleavage

The Chain Reaction of Destruction

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