Chlorine monoxide: Difference between revisions
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{{about|the compound ClO|the oxoanion with the formula ClO<sup>−</sup>|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}} |
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|Section2={{Chembox Properties |
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| Cl=1 | O=1 |
| Cl=1 | O=1 |
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|Section3={{Chembox Hazards |
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|Section4={{Chembox Thermochemistry |
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| DeltaHf = 101.8 kJ/mol<ref name="holleman_wiberg"/> |
| DeltaHf = 101.8 kJ/mol<ref name="holleman_wiberg"/> |
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:Cl<sup>•</sup> + O<sub>3</sub> → ClO<sup>•</sup> + O<sub>2</sub> |
:Cl<sup>•</sup> + O<sub>3</sub> → ClO<sup>•</sup> + O<sub>2</sub> |
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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: |
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: |
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: ClO<sup>•</sup> + O<sup>•</sup> → Cl<sup>•</sup> + O<sub>2</sub> |
: ClO<sup>•</sup> + O<sup>•</sup> → Cl<sup>•</sup> + O<sub>2</sub> |
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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: |
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2. Why the "+" on the products side ? was something missing? --> |
2. Why the "+" on the products side ? was something missing? --> |
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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. |
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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 [[ |
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]]. |
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== References == |
== References == |
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[[Category:Chlorine oxides]] |
[[Category:Chlorine oxides]] |
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[[Category:Free radicals]] |
[[Category:Free radicals]] |
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[[Category:Diatomic molecules]] |
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{{Inorganic-compound-stub}} |
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Latest revision as of 09:17, 13 September 2024
 | This article relies largely or entirely on a single source. (February 2024) |
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| Names | |||
|---|---|---|---|
| Preferred IUPAC name
Chlorine monoxide | |||
| Systematic IUPAC name
Chlorooxidanyl | |||
| Other names
Chlorine(II) oxide
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| Identifiers | |||
3D model (JSmol)
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| Abbreviations | ClO• | ||
| ChEBI | |||
| ChemSpider | |||
| MeSH | Chlorosyl | ||
PubChem CID
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| UNII | |||
CompTox Dashboard (EPA)
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| Properties | |||
| ClO | |||
| Molar mass | 51.45 g·mol−1 | ||
| Thermochemistry | |||
Std enthalpy of
formation (ΔfH⦵298) |
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).
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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]- ^ a b Egon Wiberg; Nils Wiberg; Arnold Frederick Holleman (2001). Inorganic chemistry. Academic Press. p. 462. ISBN 0-12-352651-5.