Editing Ozone depletion (section)

== Ozone cycle overview ==
[[File:Ozone cycle.svg|thumb|upright=1.7|The ozone cycle]]
Three forms (or [[allotropy|allotropes]]) of [[oxygen]] are involved in the [[ozone-oxygen cycle]]: oxygen atoms (O or atomic oxygen), oxygen gas ({{chem|O|2}} or diatomic oxygen), and ozone gas ({{chem|O|3}} or triatomic oxygen).<ref>{{Cite web |date=1999-07-30 |title=Ozone |url=https://earthobservatory.nasa.gov/features/Ozone/ozone_2.php |access-date=2022-04-06 |website=earthobservatory.nasa.gov |language=en}}</ref> Ozone is formed in the stratosphere when oxygen gas molecules photodissociate after absorbing UVC photons. This converts a single {{chem|O|2}} into two atomic oxygen [[Radical (chemistry)|radicals]]. The atomic oxygen radicals then combine with separate {{chem|O|2}} molecules to create two {{chem|O|3}} molecules. These ozone molecules absorb UVB light, following which ozone splits into a molecule of {{chem|O|2}} and an oxygen atom. The oxygen atom then joins up with an oxygen molecule to regenerate ozone. This is a continuing process that terminates when an oxygen atom recombines with an ozone molecule to make two {{chem|O|2}} molecules. It is worth noting that ozone is the only atmospheric gas that absorbs UVB light.

:O + {{chem|O|3}} → 2 {{chem|O|2}}

[[File:Ozone altitude UV graph.svg|thumb|upright=1.7|Levels of ozone at various altitudes ([[w:Dobson unit|DU/km]]) and absorption of different bands of ultraviolet radiation: In essence, all UVC is absorbed by diatomic oxygen (100–200&nbsp;nm) or by ozone (triatomic oxygen) (200–280&nbsp;nm) in the atmosphere. The ozone layer also absorbs most UVB. In contrast, UVA is hardly absorbed and most of it reaches the ground. Consequently UVA makes up almost all the UV light that penetrates the Earth's atmosphere.]]

The total amount of ozone in the stratosphere is determined by a balance between photochemical production and recombination.

Ozone can be destroyed by a number of [[free radical]] catalysts; the most important are the [[hydroxyl radical]] (OH·), [[nitric oxide]] radical (NO·), [[chlorine]] radical (Cl·) and [[bromine]] radical (Br·). The dot is a notation to indicate that each species has an unpaired electron and is thus extremely reactive. The effectiveness of different [[halogen]]s and [[pseudohalogen]]s as catalysts for ozone destruction varies, in part due to differing routes to regenerate the original radical after reacting with ozone or dioxygen.<ref>{{cite journal|doi=10.5194/acpd-4-5381-2004|doi-access=free |title=Atmospheric pseudohalogen chemistry |last1=Lary |first1=D. J. |journal=Atmospheric Chemistry & Physics Discussions |date=2004 |volume=4 |issue=5 |page=5381 |bibcode=2004ACPD....4.5381L }}</ref>

While all of the relevant radicals have both natural and man-made sources, human activity has impacted some more than others. As of 2020, most of the OH· and NO· in the stratosphere is naturally occurring, but human activity has drastically increased the levels of chlorine and bromine.<ref>{{Cite web |date=2009-06-01 |title=World of Change: Antarctic Ozone Hole |url=https://earthobservatory.nasa.gov/world-of-change/Ozone |access-date=2020-06-26 |website=earthobservatory.nasa.gov |language=en}}</ref> These elements are found in stable organic compounds, especially [[chlorofluorocarbon]]s, which can travel to the stratosphere without being destroyed in the troposphere due to their low reactivity. Once in the stratosphere, the Cl and Br atoms are released from the parent compounds by the action of ultraviolet light, e.g.

:{{chem|CFCl|3}} + [[electromagnetic radiation]] → Cl· + ·{{chem|CFCl|2}}

[[File:TOMS Global Ozone 65N-65S.png|thumb|upright=1.7|Global monthly average total ozone amount]]
Ozone is a highly reactive molecule that easily reduces to the more stable oxygen form with the assistance of a catalyst. Cl and Br atoms destroy ozone molecules through a variety of [[catalysis|catalytic]] cycles. In the simplest example of such a cycle,<ref>{{cite book |author= Newman, Paul A. |chapter= Chapter 5: Stratospheric Photochemistry Section 4.2.8 ClX catalytic reactions |chapter-url= http://www.ccpo.odu.edu/~lizsmith/SEES/ozone/class/Chap_5/index.htm |editor= Todaro, Richard M. |title= Stratospheric ozone: an electronic textbook |publisher= NASA Goddard Space Flight Center Atmospheric Chemistry and Dynamics Branch |url= http://www.ccpo.odu.edu/SEES/ozone/oz_class.htm |access-date= May 28, 2016 }}</ref> a chlorine atom reacts with an ozone molecule ({{chem|O|3}}), taking an oxygen atom to form chlorine monoxide (ClO) and leaving an oxygen molecule ({{chem|O|2}}). The ClO can react with a second molecule of ozone, releasing the chlorine atom and yielding two molecules of oxygen. The chemical shorthand for these gas-phase reactions is:
* Cl· + {{chem|O|3}} → ClO + {{chem|O|2}}<br /> A chlorine atom removes an oxygen atom from an ozone molecule to make a ClO molecule
* ClO + {{chem|O|3}} → Cl· + 2 {{chem|O|2}}<br /> This ClO can also remove an oxygen atom from another ozone molecule; the chlorine is free to repeat this two-step cycle

The overall effect is a decrease in the amount of ozone, though the rate of these processes can be decreased by the effects of [[null cycle]]s. More complicated mechanisms have also been discovered that lead to ozone destruction in the lower stratosphere.

A single chlorine atom would continuously destroy ozone (thus a catalyst) for up to two years (the time scale for transport back down to the troposphere) except for reactions that remove it from this cycle by forming reservoir species such as [[hydrogen chloride]] (HCl) and [[chlorine nitrate]] ({{chem|ClONO|2}}). Bromine is even more efficient than chlorine at destroying ozone on a per-atom basis, but there is much less bromine in the atmosphere at present. Both chlorine and bromine contribute significantly to overall ozone depletion. Laboratory studies have also shown that fluorine and iodine atoms participate in analogous catalytic cycles. However, fluorine atoms react rapidly with water vapour, methane and hydrogen to form strongly bound [[hydrogen fluoride]] (HF) in the Earth's stratosphere,<ref>{{cite journal |first1=P. |last1=Ricaud |first2=F. |last2=Lefèvre |title=Fluorine in the Atmosphere |journal=Advances in Fluorine Science |volume=1 |pages=1–32 See 12–13 |date=2006 |doi=10.1016/S1872-0358(06)01001-3 |id=hal-00256296 |url=https://hal.archives-ouvertes.fr/hal-00256296/document   |quote=Thus, fluorine chemistry does not represent a significant sink for stratospheric ozone. All fluorine released from the source gases ends up in the form of HF, which accumulates in the stratosphere (Fig. 8). ... The high stability of HF makes it an effective tracer of fluorine input in the stratosphere arising from fluorinated anthropogenic gases}}</ref> while organic molecules containing iodine react so rapidly in the lower atmosphere that they do not reach the stratosphere in significant quantities.<ref>{{cite web |url=https://csl.noaa.gov/assessments/ozone/2010/twentyquestions/Q7.pdf |archive-url=https://web.archive.org/web/20210226175130/https://csl.noaa.gov/assessments/ozone/2010/twentyquestions/Q7.pdf |archive-date=2021-02-26 |url-status=live |pages=3–4 |title=Q7 What emissions from human activities lead to ozone depletion? |work=20 Questions: 2010 Update: Section II The Ozone Depletion Process |publisher=Chemical Sciences Laboratory, National Oceanic and Atmospheric Administration (NOAA) |access-date=22 October 2022 |quote=Iodine is a component of several gases that are naturally emitted from the oceans. Although iodine can participate in ozone destruction reactions, these iodine-containing source gases generally have very short lifetimes and, as a result, only a very small fraction reaches the stratosphere. There are large uncertainties in how these emissions vary with season and geographical region.}}</ref>

A single chlorine atom is able to react with an average of 100,000 ozone molecules before it is removed from the catalytic cycle. This fact plus the amount of chlorine released into the atmosphere yearly by chlorofluorocarbons (CFCs) and hydrochlorofluorocarbons (HCFCs) demonstrates the danger of CFCs and HCFCs to the environment.<ref>{{cite web|url=http://www.eoearth.org/article/Stratospheric_Ozone_Depletion_by_Chlorofluorocarbons_(Nobel_Lecture) |title=Stratospheric Ozone Depletion by Chlorofluorocarbons (Nobel Lecture)—Encyclopedia of Earth |publisher=Eoearth.org |url-status=dead |archive-url=https://web.archive.org/web/20110909064451/http://www.eoearth.org/article/Stratospheric_Ozone_Depletion_by_Chlorofluorocarbons_%28Nobel_Lecture%29 |archive-date=September 9, 2011 }}</ref><ref>{{Cite web |last=Laboratory (CSL) |first=NOAA Chemical Sciences |title=NOAA CSL: Scientific Assessment of Ozone Depletion: 2010 |url=https://csl.noaa.gov/assessments/ozone/2010/ |access-date=2024-04-01 |website=csl.noaa.gov |language=en}}</ref>