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To draw the Lewis structure, we first need to count the total number of valence electrons in the CFC-11 molecule. The number of valence electrons for each atom is as follows: - Carbon (C) has 4 valence electrons - Fluorine (F) has 7 valence electrons - Chlorine (Cl) has 7 valence electrons Considering there are three Cl atoms, the total number of valence electrons in the CFC-11 molecule will be: Total valence electrons = 4 (C) + 7 (F) + 3 * 7 (Cl) = 4 + 7 + 21 = 32

We need to determine the central atom in the molecule, which would typically be the atom with the lowest electronegativity (unless it is hydrogen). In CFC-11, the central atom is the Carbon (C).

We will now connect the atoms by single bonds to form a skeletal structure. The carbon atom will be at the center, bonded to the fluorine and the three chlorine atoms. Cl | Cl - C - F | Cl

Now, we should distribute the remaining valence electrons as lone pairs to complete the octet for each atom. We used 8 valence electrons in the single bonds (2 electrons for each bond). There are 24 valence electrons left to be distributed. We add 6 electrons (3 lone pairs) to each of the three chlorine atoms and 6 electrons (3 lone pairs) to the fluorine atom. Cl : : : | : : Cl - C - F : : Cl : :

Verify that all atoms satisfy the octet rule; meaning they have 8 electrons in their valence shell. Carbon has 4 bonds (4 * 2 = 8 electrons), Fluorine has 1 bond and 3 lone pairs (2 + 6 = 8 electrons), and each Chlorine atom has 1 bond and 3 lone pairs (2 + 6 = 8 electrons). All atoms satisfy the octet rule.

CFC-11 can effectively deplete ozone in the stratosphere due to the formation of highly reactive chlorine radicals. The Cl-C bond in CFC-11 is weaker than the other C-X (X = halogen) bonds due to the large size and lower electronegativity of chlorine, making it easier for UV radiation to break the bond. When CFC-11 absorbs UV radiation, a high-energy photon causes the dissociation of a chlorine atom from the molecule: \(\mathrm{CFCl}_{3} + \textrm{UV} \longrightarrow \mathrm{CFCl}_{2} + \mathrm{Cl}^*\) The chlorine radical (Cl*) reacts with ozone, removing an oxygen atom, and forming chlorine monoxide (ClO) and molecular oxygen (O2). \(\mathrm{Cl}^* + \mathrm{O}_{3} \longrightarrow \mathrm{ClO} + \mathrm{O}_{2}\) Chlorine monoxide can then react with an oxygen atom, regenerating the chlorine radical and forming molecular oxygen. \(\mathrm{ClO} + \mathrm{O} \longrightarrow \mathrm{Cl}^* + \mathrm{O}_{2}\) As a result, the chlorine radicals can continuously destroy thousands of ozone molecules before being removed from the stratosphere by other chemical reactions, leading to a severe depletion of stratospheric ozone.