Chlorine monoxide: Difference between revisions
No edit summary |
No edit summary |
||
| Line 48: | Line 48: | ||
:O· + {{chem|O|3}} → + 2{{chem|O|2}} |
:O· + {{chem|O|3}} → + 2{{chem|O|2}} |
||
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 |
This has been a significant impact of the use of [[Chlorofluorocarbon|CFC]]s on the upper stratosphere, but many counties in the world has agreed to ban the use of CFCs so as to preserve the ozone layer. 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 monoxide]], terminating the radical reaction. However, the concentration of CFCs in atmosphere is very low. Hence the probability of a terminating reaction is exceedingly low, meaning each radical can decompose many thousands of molecules of ozone. Chlorine radicals also remain in the for about 500 years, hence it can decompose thousands or tens of thousands of molecules of ozone. |
||
So even if the use of CFCs is completely banned, the Chlorine radicals will continue to decompose ozone molecules until it reacts with chlorine monoxide to form [[dichlorine monoxide]] or until it reaches the end of its life. |
|||
== References == |
== References == |
||
Revision as of 13:10, 11 April 2016
| |||
| 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 | ||
PubChem CID
|
|||
CompTox Dashboard (EPA)
|
|||
| |||
| |||
| 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).
| |||
The 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· + O
3 → ClO· + O
2
This reaction causes the depletion of the ozone layer.[1] This reaction can go on and the ClO· radicals can go on to react as such:
- ClO· + O· → Cl· + O
2
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· + O
3 → + 2O
2
This has been a significant impact of the use of CFCs on the upper stratosphere, but many counties in the world has agreed to ban the use of CFCs so as to preserve the ozone layer. 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 radicals react to form dichlorine monoxide, terminating the radical reaction. However, the concentration of CFCs in atmosphere is very low. Hence the probability of a terminating reaction is exceedingly low, meaning each radical can decompose many thousands of molecules of ozone. Chlorine radicals also remain in the for about 500 years, hence it can decompose thousands or tens of thousands of molecules of ozone.
So even if the use of CFCs is completely banned, the Chlorine radicals will continue to decompose ozone molecules until it reacts with chlorine monoxide to form dichlorine monoxide or until it reaches the end of its life.
References
- ^ a b Egon Wiberg; Nils Wiberg; Arnold Frederick Holleman (2001). Inorganic chemistry. Academic Press. p. 462. ISBN 0-12-352651-5.
 | This inorganic compound–related article is a stub. You can help Wikipedia by expanding it. |