Radicals in the Ozone

Lewis’ octet theory is used to predict most formulas and molecular structures for molecular compounds. The theory centers on the resemblance to noble gas ns²np⁶ valence octets. Therefore, it is most useful in accounting for the formulas of compounds of the representative elements (particularly those in periods 1 and 2), whose valence electrons are also s and p electrons. The octet rule is not as useful when dealing with compounds involving the transition elements, most of which involve d or f orbitals when bonding. In addition, there are also some exceptions to the Lewis theory among the representative elements. The exceptions fall into three categories:

(1.) Electron Deficient Species are compounds that do not have enough electrons to achieve octets around all atoms, even though they are stable molecules. These types of compounds typically contain beryllium (Be), boron (B), or aluminum (Al). For example, BF₃ is an electron-deficient compound. Each fluorine has a complete octet, but the boron has only six electrons. Any compound in which one or more atoms do not have complete octets are referred to as electron deficient. Electron-deficient molecules typically react with species containing lone pairs of electrons, acquiring octets by forming coordinate covalent bonds.

(2.) Species with Expanded Octets are elements in the third period and below that can accommodate more than an octet of electrons. Elements such as silicon (Si), phosphorus (P), sulfur (S), chlorine (Cl), bromine (Br), and iodine (I) obey the octet rule in many cases, but in other circumstances, they can form more bonds than the rule allows. For example, BrF₅ is an example of a compound with an expanded octet. As you can see, the bromine forms five bonds and has a lone pair of electrons, which means it has 12 electrons in its valence shell. An atom such as the bromine that has more than an octet is said to have expanded its valence shell. This can only happen if the valence shell has enough orbitals to accommodate the extra electrons. This means that it is impossible to expand the valence shell of an atom in the second period because there is no such thing as a 2d orbital.

(3.) Free Radicals contain at least one unpaired electron. Most molecules or complex ions discussed in general chemistry courses are demonstrated to have pairs of electrons. However, there are some stable molecules that contain an odd number of electrons, which are called free radicals. These free radicals can play many important roles in applied chemistry fields including biology, medicine, and astrochemistry. Free radicals are usually very reactive compared to other molecules with all paired electrons. Specifically, they tend to combine with other molecules to pair their unpaired electron with a partner of opposite spin. Since the majority of molecules have all paired electrons, these reactions usually produce another free radical. Here we will take a closer look at free radicals that form and react in the ozone layer of the stratosphere.

The general mechanism in which radicals degrade the ozone can be summarized as:

Water in the stratosphere is quickly decomposed by electromagnetic radiation into hydrogen and hydroxide radicals:

\[\ce{H2O + hv -> H \cdot + HO\cdot}\nonumber\]

Then the hydroxide radical will react with ozone to form a third radical, the peroxide radical:

\[\ce{HO\cdot + O3 -> HOO\cdot + O2}\nonumber\]

Methyl chloride, CH₃Cl, is both a naturally produced and man-made chemical that forms chlorine radicals, which are some of the most reactive substances responsible for catalytic ozone decomposition. Methyl chloride is broken down by light to form methyl and chlorine radicals:

\[\ce{CH3Cl + hv -> CH3\cdot + Cl\cdot} \nonumber\]

Besides the general mechanism shown above, there are two additional paths by which ozone is degraded by ClO_x substances:

\[\ce{HOO \cdot + ClO \cdot -> HOCl + O2 }\nonumber\]

\[\ce{HOCl + hv -> HO\cdot + Cl\cdot }\nonumber\]

\[\ce{Cl\cdot + O3 -> Cl\cdot + O2 }\nonumber\]

\[\ce{HO \cdot + O3 -> HOO\cdot + O2 }\nonumber\]

Net: \[\ce{2O3 -> 3O2 }\nonumber\]

The second pathway involves a bromine radical:

\[\ce{BrO\cdot + ClO\cdot -> Br\cdot + Cl\cdot + O2 }\nonumber\]

\[\ce{Br\cdot + O3 -> BrO\dot + O2 }\nonumber\]

\[\ce{Cl\cdot + O3 -> ClO\cdot + O2 }\nonumber\]

Net: \[\ce{2O3 -> 3O2 }\nonumber\]

This degradation of the ozone found in the stratosphere is bad because ozone protects humans by absorbing some UV radiation from the sun. This UV radiation can cause mutations that lead to cancer.

From ChemPRIME: 7.1: Exceptions to the Octet Rule

REFERENCES:

[1] s-owl.cengage.com/ebooks/vini...Sect8-2-c.html

[2] http://chemwiki.ucdavis.edu/Theoretical_Chemistry/Chemical_Bonding/Lewis_Theory_of_Bonding/Violations_of_the_Octet_Rule

[3] Debra Boehmler (2012). Air. The Ozone Layer. (ELMS Content Unit). Available from: </webapps/blackboard/execute/launcher?type=Course&id=_825831_1&url=>. [Accessed 1 Jun 2012 ].

Contributors and Attributions

• Ed Vitz (Kutztown University), John W. Moore (UW-Madison), Justin Shorb (Hope College), Xavier Prat-Resina (University of Minnesota Rochester), Tim Wendorff, and Adam Hahn.