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 need to determine the total number of valence electrons for each atom, in this case Carbon (C), Fluorine (F), and three Chlorine atoms (Cl). Each atom contributes the following number of valence electrons: - Carbon: 4 valence electrons - Fluorine: 7 valence electrons - Chlorine: 7 valence electrons (each) Since there are three chlorines in the molecule, we get a total of 4 + 7 + 3(7) = 25 valence electrons.

Carbon should be placed in the center as it has the lowest electronegativity compared to other atoms. It's important to remember that carbon's preferred configuration has four bond pairings. The fluorine atom and the three chlorine atoms will be bonded surrounding the central carbon. The molecular structure, ignoring electrons will now look like this: C(FCl_3)

First, we will try creating single bonds between the central carbon atom and each of the surrounding atoms. FORMULA: C-F, C-Cl, C-Cl, C-Cl Once these single bonds are in place, we can subtract the electrons used in these single bonds from our total valence electrons. Each bond has used two electrons (one from each atom), bringing the total valence electrons used to 8, leaving us with 17 valence electrons.

We need to distribute the remaining electrons (17) around the fluorine and three chlorine atoms. We start by giving the fluorine a filled octet (8 electrons) by adding six more electrons around it. We will also place six electrons each around the three chlorine atoms; these will be in the pairs as well. After all the electrons have been distributed, we have used all 25 valence electrons, and the Lewis structure of CFC-11 is as follows: F | Cl-C-Cl | Cl

Chlorofluorocarbons (CFCs) like CFC-11 are responsible for the depletion of stratospheric ozone. One of the main chemical characteristics of CFCs that allow them to deplete stratospheric ozone is related to their stability. CFCs are chemically unreactive, which allows them to persist in the atmosphere, especially in the lower layers like the troposphere. They can be transported to the higher layers (stratosphere) via atmospheric circulation, where they are exposed to high-energy ultraviolet (UV) radiation. The strong bonds between Carbon and Fluorine, as well as Carbon and Chlorine, contribute to their stability and chemical inertness. In the stratosphere, the UV radiation breaks the Chlorine-Fluorine or the Chlorine-Carbon bond in the CFC molecule, releasing the reactive Chlorine radical. CFCl3 (CFC-11) + UV radiation → CFCl2 + Cl The reactive Chlorine radical then reacts with ozone (O3), breaking it down into oxygen gas (O2) and a chlorine oxide radical. Cl + O3 → ClO + O2 The chlorine oxide radical can then react with an oxygen atom, regenerating the chlorine radical and involving it in another reaction that destroys more ozone. ClO + O → Cl + O2 This chain reaction makes the CFCs able to effectively deplete stratospheric ozone, which has detrimental environmental effects. In summary, the stability of CFCs enables them to persist and rise to the stratosphere, where UV radiation causes them to release reactive chlorine radicals, which then take part in an ozone-depleting chain reaction.