Draw the Lewis structure for the chlorofluorocarbon CFC-11, \(\mathrm{CFCl}_{3}\). What chemical characteristics of this substance allow it to effectively deplete stratospheric ozone?

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.