Editing Ozone (section)

==Reactions==
{{More citations needed section|date=March 2022}}
Ozone is among the most powerful [[oxidizing agent]]s known, far stronger than {{chem2|O2}}. It is also unstable at high concentrations, decaying into ordinary diatomic oxygen. Its [[half-life]] varies with atmospheric conditions such as temperature, humidity, and air movement. Under laboratory conditions, the half-life will average ~1500 minutes (25 hours) in still air at room temperature (24&nbsp;°C), zero humidity with zero air changes per hour.<ref>Half-life time of ozone as a function of air conditions and movement McClurkin, J.D.*#1, Maier, D.E.2. {{doi|10.5073/jka.2010.425.167.326}}</ref>
:<chem>2 O3 -> 3 O2</chem>

This reaction proceeds more rapidly with increasing temperature. [[Deflagration]] of ozone can be triggered by a spark and can occur in ozone concentrations of 10 [[wt%]] or higher.<ref>{{cite journal |title=Explosion properties of highly concentrated ozone gas |year=2005 |last1=Koike |first1=K |last2=Nifuku |first2=M |last3=Izumi |first3=K |last4=Nakamura |first4=S |last5=Fujiwara |first5=S |last6=Horiguchi |first6=S |journal=Journal of Loss Prevention in the Process Industries |volume=18 |page=465 |issue=4–6 |bibcode=2005JLPPI..18..465K |doi=10.1016/j.jlp.2005.07.020 |url=http://www.iitk.ac.in/che/jpg/papersb/full%20papers/K-106.pdf |archive-url=https://web.archive.org/web/20090327085613/http://www.iitk.ac.in/che/jpg/papersb/full%20papers/K-106.pdf |archive-date=2009-03-27}}</ref>

Ozone can also be produced from oxygen at the anode of an electrochemical cell. This reaction can create smaller quantities of ozone for research purposes.<ref>{{cite web |title=Electrochemical Production of High-Concentration Ozone-Water Using Freestanding Perforated Diamond Electrodes |url=https://www.researchgate.net/publication/244688668}}</ref>
:<math chem>\ce{O3_{(g)}{} + 2H+{} + 2e- <=> O2_{(g)}{} + H2O} \quad (E^\circ = \text{2.075 V})</math><ref>{{cite book |title= Quantitative Chemical Analysis |last=Harris |first=Daniel C. |publisher=W. H. Freeman |year=2007 |isbn=978-0-7167-7694-9 |pages=[https://archive.org/details/quantitativechem00harr_464/page/n297 279] |url=https://archive.org/details/quantitativechem00harr_464 |url-access=limited}}</ref>

This can be observed as an unwanted reaction in a [[Hoffman apparatus]] during the electrolysis of water when the voltage is set above the necessary voltage.

===With metals===
Ozone oxidizes most [[metal]]s (except [[gold]], [[platinum]], and [[iridium]]) into [[oxide]]s of the metals in their highest [[oxidation state]]. For example:
:<math chem>\begin{align}
& \ce{Cu + O3 -> CuO + O2} \\
& \ce{2 Ag + O3 -> Ag2O + O2}
\end{align}</math>

===With nitrogen and carbon compounds===
Ozone oxidizes [[nitric oxide]] to [[nitrogen dioxide]]:
: <chem>NO + O3 -> NO2 + O2</chem>
This reaction is accompanied by [[chemiluminescence]].

The {{chem2|NO2}} can be further oxidized to [[nitrate radical]]:
: <chem>NO2 + O3 -> NO3 + O2</chem>
The {{chem2|NO3}} formed can react with {{chem2|NO2}} to form [[dinitrogen pentoxide]] ({{chem2|N2O5}}).

Solid [[nitronium perchlorate]] can be made from {{chem2|NO2, ClO2}}, and {{chem2|O3}} gases:
: <chem>NO2 + ClO2 + 2 O3 -> NO2ClO4 + 2 O2</chem>

Ozone does not react with ammonium [[salt (chemistry)|salts]], but it oxidizes [[ammonia]] to [[ammonium nitrate]]:
: <chem>2 NH3 + 4 O3 -> NH4NO3 + 4 O2 + H2O</chem>

Ozone reacts with [[carbon]]<!--graphite?--> to form [[carbon dioxide]], even at room temperature:
: <chem>C + 2 O3 -> CO2 + 2 O2</chem>

===With sulfur compounds===
Ozone oxidizes [[sulfide]]s to [[sulfate]]s. For example, [[lead(II) sulfide]] is oxidized to [[lead(II) sulfate]]:
: <chem>PbS + 4 O3 -> PbSO4 + 4 O2</chem>

[[Sulfuric acid]] can be produced from ozone, water and either elemental [[sulfur]] or [[sulfur dioxide]]:
:<math chem>\begin{align}
& \ce{S + H2O + O3 -> H2SO4} \\
& \ce{3 SO2 + 3 H2O + O3 -> 3 H2SO4}
\end{align}</math>

In the [[gas phase]], ozone reacts with [[hydrogen sulfide]] to form sulfur dioxide:
: <chem>H2S + O3 -> SO2 + H2O</chem>

In an [[aqueous]] solution, however, two competing simultaneous reactions occur, one to produce elemental sulfur, and one to produce [[sulfuric acid]]:
:<math chem>\begin{align}
& \ce{H2S + O3 -> S + O2 + H2O} \\
& \ce{3 H2S + 4 O3 -> 3 H2SO4}
\end{align}</math>

===With alkenes and alkynes===
{{Main|Ozonolysis}}
Alkenes can be oxidatively cleaved by ozone, in a process called [[ozonolysis]], giving alcohols, aldehydes, ketones, and carboxylic acids, depending on the second step of the workup.

[[File:General reaction equation of ozonolysis.svg|class=skin-invert-image|frameless|upright=2.2|center|General reaction equation of ozonolysis]]

Ozone can also cleave alkynes to form an [[acid anhydride]] or [[diketone]] product.<ref>{{cite book |last=Bailey |first=P. S. |chapter=Chapter 2 |title=Ozonation in Organic Chemistry |volume=2 |publisher=Academic Press |location=New York, NY |year=1982 |isbn=978-0-12-073102-2}}</ref> If the reaction is performed in the presence of water, the anhydride hydrolyzes to give two [[carboxylic acid]]s.

:[[File:Ozonolysis-alkyne.png|class=skin-invert-image|none|450px]]

Usually ozonolysis is carried out in a solution of [[dichloromethane]], at a temperature of −78&nbsp;°C. After a sequence of cleavage and rearrangement, an organic ozonide is formed. With reductive workup (e.g. [[zinc]] in [[acetic acid]] or [[dimethyl sulfide]]), ketones and aldehydes will be formed, with oxidative workup (e.g. aqueous or alcoholic [[hydrogen peroxide]]), carboxylic acids will be formed.<ref name="OrgChem">{{cite book
|title=Organic Chemistry, 9th Edition
|chapter=Chapter 8 Alkenes and Alkynes – Part II: Addition Reactions and Synthesis
|author1=Solomons, T.W. Graham |author2=Fryhle, Craig B.
|name-list-style=amp |publisher=Wiley
|year=2008
|isbn=978-0-470-16982-7
|page=344
}}</ref>

===Other substrates===
All three [[atom]]s of ozone may also react, as in the reaction of [[tin(II) chloride]] with [[hydrochloric acid]] and ozone:
: <chem>3 SnCl2 + 6 HCl + O3 -> 3 SnCl4 + 3 H2O</chem>
Iodine perchlorate can be made by treating [[iodine]] dissolved in cold [[anhydrous]] [[perchloric acid]] with ozone:
: <chem>I2 + 6 HClO4 + O3 -> 2 I(ClO4)3 + 3 H2O</chem>
Ozone could also react with potassium iodide to give oxygen and iodine gas that can be titrated for quantitative determination:<ref>{{cite journal |last1=Al-Baarri |first1=A. N. |last2=Legowo |first2=A. M. |last3=Abduh |first3=S. B. M. |last4=Mawarid |first4=A. A. |last5=Farizha |first5=K. M. |last6=Silvia |first6=M. |title=Production of Ozone and the Simple Detection using Potassium Iodide Titration Method |date=June 2019 |journal=IOP Conference Series: Earth and Environmental Science |language=en |volume=292 |issue=1 |page=012062 |bibcode=2019E&ES..292a2062A |issn=1755-1315 |doi=10.1088/1755-1315/292/1/012062 |doi-access=free |s2cid=198344024}}</ref>
: <chem>2KI + O3 + H2O -> 2KOH + O2 + I2</chem>

===Combustion===
Ozone can be used for [[combustion]] reactions and combustible gases; ozone provides higher temperatures than burning in [[dioxygen]] ({{chem2|O2}}). The following is a reaction for the combustion of [[carbon subnitride]] which can also cause higher temperatures:
: <chem>3 C4N2 + 4 O3 -> 12 CO + 3 N2</chem>

Ozone can react at cryogenic temperatures. At {{convert|77|K}}, atomic [[hydrogen]] reacts with liquid ozone to form a hydrogen [[superoxide]] [[radical (chemistry)|radical]], which [[dimerizes]]:<ref name="Horvath M. 1985. pg 44">{{cite book |year=1985 |title=Ozone |pages=44–49 |author1=Horvath M. |author2=Bilitzky L. |author3=Huttner J. |isbn=978-0-444-99625-1 |publisher=Elsevier}}</ref>
:<math chem>\begin{align}
& \ce{H + O3 -> HO2 + O} \\
& \ce{2 HO2 -> H2O4}
\end{align}</math>

===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>

===Reduction to ozonides===
Reduction of ozone gives the [[ozonide]] anion, {{chem2|O3-}}. Derivatives of this anion are explosive and must be stored at cryogenic temperatures. Ozonides for all the [[alkali metal]]s are known. {{chem2|KO3, RbO3}}, and {{chem2|CsO3}} can be prepared from their respective superoxides:
: <chem>KO2 + O3 -> KO3 + O2</chem>
Although {{chem2|KO3}} can be formed as above, it can also be formed from [[potassium hydroxide]] and ozone:<ref>{{Housecroft2nd|page=439}}</ref>
: <chem>2 KOH + 5 O3 -> 2 KO3 + 5 O2 + H2O</chem>
{{chem2|NaO3}} and {{chem2|LiO3}} must be prepared by action of {{chem2|CsO3}} in liquid {{chem2|NH3}} on an [[ion-exchange resin]] containing {{chem2|Na+}} or {{chem2|Li+}} ions:<ref>{{Housecroft2nd|page=265}}</ref>
: <chem>CsO3 + Na+ -> Cs+ + NaO3</chem>
A solution of [[calcium]] in ammonia reacts with ozone to give [[ammonium ozonide]] and not calcium ozonide:<ref name="Horvath M. 1985. pg 44"/>
: <math chem>\begin{align}
\ce{3 Ca + 10 NH3 + 6 O3 ->\ } & \ce{Ca*6NH3 + Ca(OH)2 + Ca(NO3)2} \\
& + \ce{2 NH4O3 + 2 O2 + H2}
\end{align}</math>

===Applications===
Ozone can be used to remove [[iron]] and [[manganese]] from [[water]], forming a [[precipitate]] which can be filtered:
:<math chem>\begin{align}
& \ce{2 Fe^2+{} + O3 + 5 H2O -> 2 Fe(OH)_3{(s)}{} + O2{} + 4 H+} \\
& \ce{2 Mn^2+{} + 2 O3 + 4 H2O -> 2 MnO(OH)_2{(s)}{} + 2 O2{} + 4 H+}
\end{align}</math>
Ozone oxidizes dissolved [[hydrogen sulfide]] in water to [[sulfurous acid]]:
: <chem>3 O3 + H2S -> H2SO3 + 3 O2</chem>

These three reactions are central in the use of ozone-based well water treatment.

Ozone detoxifies [[cyanide]]s by converting them to [[cyanate]]s.
: <chem>CN- + O3 -> CNO- + O2</chem>

Ozone completely decomposes [[urea]]:<ref>{{cite book |year=1985 |title=Ozone |pages=259, 269–270 |author1=Horvath M. |author2=Bilitzky L. |author3=Huttner J. |isbn=978-0-444-99625-1 |publisher=Elsevier}}</ref>
:<chem>(NH2)2CO + O3 -> N2 + CO2 + 2 H2O</chem>