Write the Lewis structures for the following reactive species found to contribute to the destruction of the ozone layer and indicate which are radicals: (a) chlorine monoxide, \(\mathrm{ClO}\); (b) dichloroperoxide, \(\mathrm{Cl}-\mathrm{O}-\mathrm{O}-\mathrm{Cl}\); (c) chlorine nitrate, \(\mathrm{ClONO}_{2}\) (the central \(\mathrm{O}\) atom is attached to the \(\mathrm{Cl}\) atom and to the \(\mathrm{N}\) atom of the \(\mathrm{NO}_{2}\) group); (d) chlorine peroxide, \(\mathrm{Cl}-\mathrm{O}-\mathrm{O}\).

Valence electrons are the outermost electrons of an atom, playing a crucial role in chemical bonding and reactivity. Every atom, depending on its position in the periodic table, has a characteristic number of valence electrons. For instance, chlorine (found in Group 17) has 7 valence electrons, while oxygen (in Group 16) has 6 valence electrons. When calculating the total valence electrons
in molecules or chemical species, you simply sum up the valence electrons of each constituent atom.

In the exercise, this step is essential for drawing the correct Lewis structures for ozone-depleting substances. Whether a molecule has an even or odd number of valence electrons can indicate the presence of radicals - species with at least one unpaired electron, which can be particularly reactive and thus play into the depletion of the ozone layer.

Chemical radicals are atoms, molecules, or ions that have unpaired valence electrons, making them highly reactive. They are capable of initiating chain reactions that can cause significant changes in a substance's chemical composition. In the exercise, you learned how to identify radicals by looking at the unpaired electrons in the Lewis structures. For example, chlorine monoxide (ClO) and chlorine peroxide (Cl-OO) both have an unpaired electron, indicating that they are radicals. These radicals are particularly noteworthy as they Participate in catalytic cycles that result in the depletion of the ozone layer, shielding Earth from harmful ultraviolet radiation.

The ozone layer is a critical part of Earth's atmosphere that absorbs most of the sun's harmful ultraviolet radiation. Ozone depletion occurs when certain chemicals, typically man-made substances like chlorofluorocarbons (CFCs), reach the stratosphere and break down ozone molecules, reducing the layer’s protective capacity.

The compounds you've encountered in the exercise are associated with ozone depletion due to their reactive chlorine radicals. These radicals, once free in the stratosphere, can catalyze the breakdown of ozone, turning it into ordinary oxygen and consequently allowing more ultraviolet light to reach the Earth’s surface. Understanding the role that these substances play helps in designing measures to protect the ozone layer and reduce the production of harmful chemicals.

A Lewis structure is a graphical representation that shows the bonds between atoms of a molecule and the lone pairs of electrons that may exist. It is a clue to the molecular geometry and can help predict the reactivity of the molecule. Drawing Lewis structures involves distributing valence electrons to form bonds and complete the octets (or duets for hydrogen) for each atom, or to show the presence of unpaired electrons in the case of radicals.

In the given exercise, the Lewis structures reveal which substances are radicals and provide insight into their potential to participate in reactions leading to the depletion of the ozone layer. For example, in the structure of chlorine nitrate, we see that all the electrons are paired, indicating that it is not a radical, whereas with chlorine peroxide, the presence of an unpaired electron classifies it as a radical. Drawing clear and correct Lewis structures is fundamental in understanding the behavior of these ozone-depleting substances.