Editing Ozone (section)

===Ozone decomposition===

====Types of ozone decomposition====
Ozone is a toxic substance,<ref>{{cite journal |last=Menzel |first=D. B. |title=Ozone: an overview of its toxicity in man and animals |date=1984 |journal=Journal of Toxicology and Environmental Health |volume=13 |issue=2–3 |pages=183–204 |issn=0098-4108 |bibcode=1984JTEH...13..181M |pmid=6376815 |doi=10.1080/15287398409530493}}</ref><ref name="EPA-2022">{{cite web |publisher=United States Environmental Protection Agency |title=Ozone Generators that are Sold as Air Cleaners |date=28 February 2022 |website=EPA |url=https://www.epa.gov/indoor-air-quality-iaq/ozone-generators-are-sold-air-cleaners#info-sources |access-date=28 February 2022 |url-status=live |archive-url=https://web.archive.org/web/20220209015459/https://www.epa.gov/indoor-air-quality-iaq/ozone-generators-are-sold-air-cleaners |archive-date=9 February 2022}}</ref> commonly found or generated in human environments (aircraft cabins, offices with photocopiers, laser printers, sterilizers, ...).

The [[catalysis|catalytic]] decomposition of ozone is very important to reduce pollution. This type of decomposition is the most widely used, especially with solid catalysts, and it has many advantages such as a higher conversion with a lower temperature. Furthermore, the product and the catalyst can be instantaneously separated, and this way the catalyst can be easily recovered without using any separation operation. The most-used materials in the catalytic decomposition of ozone in the gas phase are [[manganese dioxide]], transition metals such as Mn, Co, Cu, Fe, Ni, or Ag, and noble metals such as Pt, Rh, or Pd.

[[Radical (chemistry)|Free radicals]] of [[chlorine]] (Cl{{sup|'''·'''}}), formed by the action of ultraviolet radiation on chlorofluorocarbons (CFCs) and sea salt, are known to catalyze the breakdown of ozone in the atmosphere.

There are two other possibilities for decomposing ozone in the gas phase:
* Thermal decomposition, in which the ozone is decomposed using only the action of heat. The problem is that this type of decomposition is very slow with temperatures below 250&nbsp;°C. However, the decomposition rate can be increased working with higher temperatures but this would involve a high energy cost.
* Photochemical decomposition, which consists of radiating ozone with ultraviolet radiation (UV) and it gives rise to oxygen and radical peroxide.<ref>{{cite thesis |last=Roca Sánchez |first=Anna |title=Estudio cinético de la descomposición catalítica de ozono |date=2015-09-01 |url=https://riunet.upv.es/handle/10251/54140}}</ref>

====Kinetics of ozone decomposition into molecular oxygen====
The uncatalyzed process of ozone decomposition in the gas phase is a complex reaction involving two [[Elementary reaction|elementary reactions]] that finally lead to molecular oxygen,<ref>{{Cite web |last1=Flowers |first1=Paul |last2=Theopold |first2=Klaus |last3=Langley |first3=Richard |last4=William R. Robinson |first4=PhD |date=2019-02-14 |title=12.6 Reaction Mechanisms - Chemistry 2e {{!}} OpenStax |url=https://openstax.org/books/chemistry-2e/pages/12-6-reaction-mechanisms |access-date=2025-05-02 |website=openstax.org |language=English}}</ref> and this means that the reaction order and the [[rate equation|rate law]] cannot be determined by the stoichiometry of the overall reaction.

Overall reaction: <chem>2 O3 -> 3 O2</chem>

Rate law (observed): <math chem>V = \frac{K_{obs} \cdot [\ce{O3}]^2}{[\ce{O2}]}</math>

where <math> K_{obs} </math> is the observed [[Reaction rate constant|rate constant]] and <math> V </math> is the reaction rate. From the rate law above it can be determined that the partial order respect to molecular oxygen is −1 and respect to ozone is 2; therefore, the global reaction order is 1.

The first step is a unimolecular reaction wherein one molecule of ozone decomposes into two products (molecular oxygen and oxygen). The oxygen atom from the first step is a [[reactive intermediate]] because it participates as a reactant in the second step, which is a bimolecular reaction because there are two different reactants (ozone and oxygen) that give rise to molecular oxygen.

Step 1: Unimolecular reaction &nbsp; &nbsp;<chem>O3 -> O2 + O</chem>

Step 2: Bimolecular reaction &nbsp; &nbsp; <chem>O3 + O -> 2 O2</chem>

These two steps have different reaction rates and rate constants. The reaction rate laws for each of these steps are shown below:
:<math chem>V_1 = K_1 \cdot [\ce{O3}] \qquad V_2 = K_2 \cdot [\ce{O}] \cdot [\ce{O3}]</math>

The following mechanism allows to explain the rate law of the ozone decomposition observed experimentally, and also it allows to determine the reaction orders with respect to ozone and oxygen, with which the overall reaction order will be determined. 

The first step is assumed reversible and faster than the second reaction, which means that the slower [[Rate-determining step|rate determining step]] is the second reaction. This step determines the rate of product formation, and so <math> V=V_2 </math>. However, this equation depends on the concentration of oxygen (intermediate), which does not appear in the observed rate law. Since the first step is a rapid equilibrium, the concentration of the intermediate can be determined as follows:
:<math chem>K_{eq} = \frac{K_1}{K_{-1}} = \frac{[\ce{O2}] \cdot [\ce{O}]}{[\ce{O3}]}</math>
:<math chem>[\ce{O}] = \frac{K_1 \cdot [\ce{O3}]}{K_{-1} \cdot [\ce{O2}]}</math>

Then using these equations, the formation rate of molecular oxygen is as shown below:
:<math chem>V={K_2 \cdot K_1 \cdot [\ce{O_3}]^2 \over K_{-1} \cdot [\ce{O_2}]}</math>

The mechanism is consistent with the rate law observed experimentally if the rate constant ({{math|''K''<sub>obs</sub>}}) is given in terms of the individual mechanistic steps' rate constants as follows:<ref>{{cite journal |last1=Batakliev |first1=Todor |last2=Georgiev |first2=Vladimir |last3=Anachkov |first3=Metody |last4=Rakovsky |first4=Slavcho |last5=Zaikov |first5=Gennadi E. |title=Ozone decomposition |date=June 2014 |journal=Interdisciplinary Toxicology |volume=7 |issue=2 |pages=47–59 |issn=1337-6853 |pmid=26109880 |doi=10.2478/intox-2014-0008 |pmc=4427716}}</ref>
:<math chem>V={K_\text{obs} \cdot [\ce{O_3}]^2 \over [\ce{O_2}]} </math>

where <math> K_\text{obs}={K_{2} \cdot K_{1} \over K_{-1}}</math>