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{{about|the compound ClO|the oxoanion with the formula {{chem|ClO|-}}|hypochlorite|the molecule Cl<sub>2</sub>O|Dichlorine monoxide}}
{{about|the compound ClO|the oxoanion with the formula ClO<sup>−</sup>|hypochlorite|the molecule Cl<sub>2</sub>O|Dichlorine monoxide}}
{{single source|date=February 2024}}

{{Chembox
{{Chembox
| Watchedfields = changed
| Watchedfields = changed
| ImageFileL1 = Chlorosyl.svg
| ImageFileL1 = Chlorosyl.svg
| ImageFileR1 = Chlorine-monoxide-3D-vdW.png
| ImageFileR1 = Chlorine-monoxide-3D-SF.png
| PIN = Chlorine monoxide
| PIN = Chlorine monoxide
| SystematicName = Chlorooxidanyl
| SystematicName = Chlorooxidanyl
| OtherNames = Chlorine(II) oxide
| OtherNames = Chlorine(II) oxide
|Section1={{Chembox Identifiers
|Section1={{Chembox Identifiers
| Abbreviations = ClO(.)
| Abbreviations = ClO<sup>•</sup>
| CASNo_Ref = {{cascite|correct|CAS}}
| CASNo = 12301-79-0
| CASNo = 7791-21-1
| UNII_Ref = {{fdacite|correct|FDA}}
| UNII = 0EQ5I4TK19
| PubChem = 166686
| PubChem = 166686
| ChemSpiderID = 145843
| ChemSpiderID = 145843
| ChemSpiderID_Ref = {{Chemspidercite}}
| ChEBI = 29314
| ChEBI = 29314
| MeSHName = Chlorosyl
| MeSHName = Chlorosyl
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| InChIKey = NHYCGSASNAIGLD-UHFFFAOYSA-N}}
| InChIKey = NHYCGSASNAIGLD-UHFFFAOYSA-N}}
|Section2={{Chembox Properties
|Section2={{Chembox Properties
| Cl=1 | O=1
| Cl=1 | O=1
| Appearance =
| Appearance =
| Density =
| Density =
| MeltingPt =
| MeltingPt =
| BoilingPt =
| BoilingPt =
| Solubility = }}
| Solubility = }}
|Section3={{Chembox Hazards
|Section3={{Chembox Hazards
| MainHazards =
| MainHazards =
| FlashPt =
| FlashPt =
| Autoignition = }}
| AutoignitionPt = }}
|Section4={{Chembox Thermochemistry
|Section4={{Chembox Thermochemistry
| DeltaHf = 101.8 kJ/mol<ref name="holleman_wiberg"/>
| DeltaHf = 101.8 kJ/mol<ref name="holleman_wiberg"/>
| DeltaHc =
| DeltaHc =
| Entropy =
| Entropy =
| HeatCapacity =
| HeatCapacity =
}}
}}
}}
}}


'''Chlorine monoxide''' is a chemical [[Radical (chemistry)|radical]] with the chemical formula ClO. It plays an important role in the process of [[ozone depletion]]. In the [[stratosphere]], [[chlorine]] atoms react with [[ozone]] molecules to form chlorine monoxide and [[oxygen]].
'''Chlorine monoxide''' is a chemical [[Radical (chemistry)|radical]] with the chemical formula ClO<sup>•</sup>. It plays an important role in the process of [[ozone depletion]]. In the [[stratosphere]], [[chlorine]] atoms react with [[ozone]] molecules to form chlorine monoxide and [[oxygen]].


:Cl· + {{chem|O|3}} → ClO· + {{chem|O|2}}
:Cl<sup>•</sup> + O<sub>3</sub> → ClO<sup>•</sup> + O<sub>2</sub>


This reaction causes the depletion of the ozone layer.<ref name="holleman_wiberg">{{cite book | title = Inorganic chemistry | author1 = Egon Wiberg | author2 = Nils Wiberg | author3 = Arnold Frederick Holleman | publisher = Academic Press | year = 2001 | isbn = 0-12-352651-5 | page = 462}}</ref> This reaction can go on and the ClO· radicals can go on to react as such:
This reaction causes the depletion of the [[ozone layer]].<ref name="holleman_wiberg">{{cite book | title = Inorganic chemistry | author1 = Egon Wiberg | author2 = Nils Wiberg | author3 = Arnold Frederick Holleman | publisher = Academic Press | year = 2001 | isbn = 0-12-352651-5 | page = 462}}</ref> The resulting ClO<sup>•</sup> radicals can further react:


:ClO· + O· → Cl· + {{chem|O|2}}
: ClO<sup>•</sup> + O<sup>•</sup> → Cl<sup>•</sup> + O<sub>2</sub>


regenerating the chlorine radical. In this way, the overall reaction for the decomposition of ozone is catalyzed by chlorine, as ultimately chlorine remains unchanged. The overall reaction is:
regenerating the chlorine radical. In this way, the overall reaction for the decomposition of ozone is catalyzed by chlorine, as ultimately chlorine remains unchanged. The overall reaction is:


:O<sup>•</sup> + O<sub>3</sub> → 2 O<sub>2</sub><!-- The reaction was :O· + O
:O· + {{chem|O|3}} → + 2{{chem|O|2}}
3 → + 2O
2. Why the "+" on the products side ? was something missing? -->


This has been a significant impact of the use of [[Chlorofluorocarbon|CFC]]s on the upper stratosphere. The nonreactive nature of CFC's allows them to pass into the stratosphere, where they undergo photo-dissociation to form Cl radicals. These then readily form chlorine monoxide, and this cycle can continue until two [[Radical (chemistry)|radicals]] react to form [[dichlorine dioxide]], terminating the radical reaction. Because the concentration of CFCs in atmosphere is very low, the probability of a terminating reaction is exceedingly low, meaning each radical can decompose many thousands of molecules of ozone.
There has been a significant impact of the use of [[Chlorofluorocarbon|CFC]]s on the upper stratosphere, although many countries have agreed to ban the use of CFCs. The nonreactive nature of CFCs allows them to pass into the stratosphere, where they undergo photo-dissociation to form Cl radicals. These then readily form chlorine monoxide, and this cycle can continue until two [[Radical (chemistry)|radicals]] react to form [[dichlorine monoxide]], terminating the radical reaction. Because the concentration of CFCs in atmosphere is very low, the probability of a terminating reaction is exceedingly low, meaning each radical can decompose many thousands of molecules of ozone.

Even though the use of CFCs has been banned in many countries, CFCs can stay in the atmosphere for 50 to 500 years. This causes many chlorine radicals to be produced and hence a significant amount of ozone molecules are decomposed before the chlorine radicals are able to react with chlorine monoxide to form [[dichlorine monoxide]].


== References ==
== References ==
{{reflist}}
{{reflist}}

{{Chlorine compounds}}
{{Chlorine compounds}}
{{oxygen compounds}}
⚫
{{DEFAULTSORT:Chlorine Monoxide}}


⚫
{{DEFAULTSORT:Chlorine Monoxide}}
[[Category:Chlorine oxides]]
[[Category:Chlorine oxides]]
[[Category:Free radicals]]
[[Category:Free radicals]]
[[Category:Diatomic molecules]]
{{Inorganic-compound-stub}}

Latest revision as of 09:17, 13 September 2024

This article is about the compound ClO. For the oxoanion with the formula ClO−, see hypochlorite. For the molecule Cl2O, see Dichlorine monoxide.
Chlorine monoxide
Names
Preferred IUPAC name
Chlorine monoxide
Systematic IUPAC name
Chlorooxidanyl
Other names
Chlorine(II) oxide
Identifiers
3D model (JSmol)
Abbreviations ClO•
ChEBI
ChemSpider
MeSH Chlorosyl
UNII
  • InChI=1S/ClO/c1-2
    Key: NHYCGSASNAIGLD-UHFFFAOYSA-N
  • [O]Cl
Properties
ClO
Molar mass 51.45 g·mol−1
Thermochemistry
101.8 kJ/mol[1]
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

Chlorine monoxide is a chemical radical with the chemical formula ClO•. It plays an important role in the process of ozone depletion. In the stratosphere, chlorine atoms react with ozone molecules to form chlorine monoxide and oxygen.

Cl• + O3 → ClO• + O2

This reaction causes the depletion of the ozone layer.[1] The resulting ClO• radicals can further react:

ClO• + O• → Cl• + O2

regenerating the chlorine radical. In this way, the overall reaction for the decomposition of ozone is catalyzed by chlorine, as ultimately chlorine remains unchanged. The overall reaction is:

O• + O3 → 2 O2

There has been a significant impact of the use of CFCs on the upper stratosphere, although many countries have agreed to ban the use of CFCs. The nonreactive nature of CFCs allows them to pass into the stratosphere, where they undergo photo-dissociation to form Cl radicals. These then readily form chlorine monoxide, and this cycle can continue until two radicals react to form dichlorine monoxide, terminating the radical reaction. Because the concentration of CFCs in atmosphere is very low, the probability of a terminating reaction is exceedingly low, meaning each radical can decompose many thousands of molecules of ozone.

Even though the use of CFCs has been banned in many countries, CFCs can stay in the atmosphere for 50 to 500 years. This causes many chlorine radicals to be produced and hence a significant amount of ozone molecules are decomposed before the chlorine radicals are able to react with chlorine monoxide to form dichlorine monoxide.

References

[edit]
  1. ^ a b Egon Wiberg; Nils Wiberg; Arnold Frederick Holleman (2001). Inorganic chemistry. Academic Press. p. 462. ISBN 0-12-352651-5.
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