About The Ozone Depletion Process
The ozone layer is one layer of the stratosphere and the stratosphere is the second layer of the Earth’s atmosphere. The stratosphere is a mass of protective gases clinging to our planet. The ozone layer is the thin part of the Earth’s atmosphere that absorbs almost all of the sun’s harmful ultra-violet light. This is of significance to the HVAC&R industry as chemicals that have historically been used in the refrigeration process have now been shown to have a detrimental effect on the ozone layer.
When the earliest engineers first designed mechanical refrigeration, they could only use natural elements for refrigerants because synthetic chemicals were not available at that time. The reason back then for deciding to use CFCs (chlorofluorocarbons) as refrigerants is that CFCs do not affect human health at all because they are completely inert. CFC’s are fully halogenated carbons however, and are volatile derivatives of methane and ethane, a typical example being CF2Cl (Freon).
Whilst the selection of CFC’s may not have had any direct negative implications for human health, a significant adverse effect of these chemical’s properties is their extremely long lifetime in the atmosphere, which is where they accumulate and interact with stratospheric gases. This we now know may cause the depletion of ozone.
Beyond this, CFC gases are photolyzed by ultraviolet light. Strong UV-B light is blocked in the stratosphere by the ozone layer and any UV-light reaching the troposphere is too weak to degrade CFCs significantly. So the inert CFCs penetrate into the stratosphere where they are decomposed into chlorine and fluorine radicals. It is the chlorine radicals which are responsible for the ozone depletion.
There are two distinct but related observations:
- the slow and steady decline of ozone of about 4% per decade in the total volume of ozone since the late 1970s,
- a much larger but seasonal decrease in stratospheric ozone over the polar regions – referred to in this article as the ozone hole.
The detailed mechanisms of (a) forming the ozone hole over polar regions and (b) the thinning over mid-latitudes are different. The most important process in both trends is the catalytic destruction of ozone by atomic chlorine and bromides; the main source of these halogens is the photo dissociation of CFCs (freons) and BFCs (halons) and both ozone depletion mechanisms (a) and (b) strengthened as emissions of CFCs and BFCs increased.
The ozone layer is located in the lower portion of the stratosphere (10 – 50 km above the earth’s surface); ozone concentration is greatest between 20 and 40 km above sea level, where it ranges between 2 – 8 ppm. This very much diluted concentration can be best illustrated when imaging ozone at the pressure at sea level: the layer would be only as thick as a few millimeters.
The Ozone Cycle – Overview
Three forms of oxygen participate in the ozone-oxygen-cycle:
- Oxygen atoms O
- Oxygen gas O2
- Ozone gas O3
Ozone is formed in the stratosphere as a result of photons of wavelengths <240 nm being absorbed by oxygen gas before splitting into two oxygen atoms which then re-combine with other oxygen gas molecules to form ozone. Gamma rays acting on ozone result in the splitting of ozone molecules into oxygen gas molecules and oxygen atoms. This process terminates when an ozone molecule and atomic oxygen form two oxygen gas molecules i.e. O3 + O2 → 2 O2.
The ozone layer is stable if there is a balance between photochemical production and re-combination. Ozone is in a state of equilibrium in the ozone layer i.e. is it formed and decomposed due to the interaction with UV light of proper wavelengths.
The wavelengths of UV light from the sun range between 100 and 400 nm. Very dangerous short wavelength UV light is completely absorbed in the upper ozone layer at about 35 km. UV light in the range up to 315 nm is still harmful and may cause cancer or genetic defects. Ozone is very effective at absorbing this UV wavelength light i.e. the radiation strength of the UV light before entering the ozone layer is about 109 times stronger than on the earth’s surface.
Other gases in the ozone layer also interact with the ozone, with O2 and simple oxygen atoms. Several destroy modes of ozone are known and these always need free radical catalysts such as OH–, NO–, and Cl– (Br–). All of these have natural and man-made sources. At present OH– and NO– are of pure natural origin, however, the concentration of Cl– and Br– increased significantly since man-made fabrication of CFCs in the 1920s. No significant natural resource has ever been identified for CFCs. These molecules may take up to 15 years to reach the stratosphere and they can stay there for a century.
Hydroxide is a diatomic anion with the chemical formula OH− . It consists of an oxygen and hydrogen atom held together by a covalent bond, and carries a negative electric charge (anion)
NO− is nitroxide anion and Cl− and Br−are chloride and bromide anions
CFCs find their way in the stratosphere without being destroyed in the troposphere due to their low chemical reactivity. Once in the stratosphere, Cl– (and Br–) ions are created by absorption of UV light:
C F Cl3 + hν → C F Cl2+ + Cl–
Now there exists a variety of catalytic cycles. Cl– ions are efficient in interacting with ozone since the ions are not consumed during chemical reactions but only recycled:
Cl– + O3 → Cl O– + O2
Cl– + O3 → Cl O– + O2
Cl O– + Cl O– + m → Cl2 O2 + m
Cl2 O2 + sunlight → 2 Cl– + O2
In total: 2 O3 → 3 O2
In any chemical reaction we see the exchange of particles between the substances, either full atoms or molecules or parts thereof. So, exchange of electrons and the generation of ions are a good example of that. It also turns out that some reactions need external energy to take place (i.e. the quantity “hν” with h as the Planck quantum and ν (Greek nu) as the frequency of a ϒ quant). Once the reaction is over the total energy balance must be reassured and in order to collect all contributions a quantity m is introduced which represents a fictive and inert molecule that absorbs excess molecular energies which might have occurred during the reaction. Due to Einstein’s famous formula E = m c² (with E as the energy, m as a mass and c as the velocity of light) any energy corresponds to a mass.
More complicated mechanisms have been discovered leading to ozone destruction also in the lower atmosphere.
This means that chlorine radicals catalyze the decomposition from ozone to “normal” oxygen molecules or in other words: the UV light triggered decomposition of CFCs leads to Cl O– radicals.
To give an idea about the effects of chlorine onto the ozone gas: a single chlorine atom can react with up to 106 ozone molecules and keeps on destroying ozone for up to two years; this period corresponds to the duration of the way back to the troposphere if no other reactions take place leading to substances like HCl or Cl O NO2:
Cl O + NO2 → Cl O NO2 (chlorine nitrate)
or Cl O + NO + CH4 → HCl + … (hydrogen chloride)
Cl O NO2 + HCL →HNO3
Important to know that both HCl and Cl O NO2 do not react with ozone and are chemically stable products.
Ozone is not measured in concentrations (in ppm) at certain levels in the atmosphere but by the total ozone column above a point on the Earth’s surface; the unit to measure ozone in columns is the Dobson Unit (DU). Measurements have been performed by use of the Total Ozone Mapping Spectrometer (TOMS) which is a satellite instrument to measure ozone. Five of these instruments have been built, four of these entered in orbits:
- Nimbus-7 and Meteor-3: daily measurements between November 1978 and December 1994
- ADEOS TOMS: measurements between August 1996 and June 1997
- Earth Probe TOMS: replaced failed ADEOS TOMS and made measurements until December 2006
- Quik TOMS: launched in September 2001 and failed to reach orbit
Since January 2006 the Ozone Monitoring Instrument (OMI) has replaced the Earth Probe TOMS. OMI resides on the Aura satellite (EOS CH-1, Earth Observing System) and uses ultraviolet and visible radiation to produce daily high-resolution maps. OMI was developed by the Finnish Meteorological Institute and the Netherlands Agency for Aerospace Programmes.
The image below shows the minimum ozone values in DUs:
In austral spring ozone is reduced by about 70% in the ozone column over Antarctica; this was first observed 1985. In the 90s, September and October values were 40-50% lower than in the pre-ozone-hole era. The reduction of ozone is about 30% over the Arctic region.
Reactions that take place on polar stratospheric clouds (PSCs) likely play an important role in ozone deletion. PSCs form in the extreme cold of Antarctic stratosphere. This is why ozone holes first formed, and are deeper, over Antarctica.
In middle latitudes it is preferable to speak of ozone depletion rather than holes. Declines are about 3% below pre-1980 values for 35–60°N and about 6% for 35–60°S. In the tropics, there are no trends.
Ozone depletion also explains much of the observed reduction in stratospheric and upper tropospheric temperatures. The source of the warmth of the stratosphere is the absorption of UV radiation by ozone, hence reduced ozone leads to cooling. The World Meteorological Organization Global Ozone Research and Monitoring Project (Report No. 44) comes out strongly in favor for the Montreal Protocol, but notes that a UNEP 1994 Assessment overestimated ozone loss for the 1994–1997 periods. So predictions of ozone levels remain difficult.
The Antarctic ozone hole is an area of the Antarctic stratosphere in which recent ozone levels have dropped to as low as 33% of their pre-1975 values. The ozone hole occurs during the Antarctic spring, from September to early December, as strong westerly winds start to circulate around the continent and create an “atmospheric container”. Within this polar vortex, over 50% of the lower stratospheric ozone is destroyed during the Antarctic spring.
As explained above, the primary cause of ozone depletion is the presence of chlorine-containing source gases. In the presence of UV light these gases dissociate and release chlorine atoms, which then go on to catalyze ozone destruction. The Cl-catalyzed ozone depletion can take place in the gas phase, but it is enhanced in the presence of polar stratospheric clouds (PSC).
These polar stratospheric clouds (PSC) form during winter in the extreme cold. Polar winters are dark, consisting of 3 months without solar radiation. The lack of sunlight contributes to a decrease in temperature and the polar vortex traps and chills air. Temperatures hover around or below -80 °C. These low temperatures form cloud particles. There are three types of PSC clouds that provide surfaces for chemical reactions that lead to ozone destruction:
- nitric acid trihydrate clouds,
- slowly cooling water-ice clouds,
- rapid cooling water-ice clouds.
Most of the chlorine in the stratosphere resides in stable “reservoir” compounds, i.e. hydrochloric acid (HCl) and chlorine nitrate (ClONO2). During the Antarctic winter and spring, however, reactions on the surface of the polar stratospheric cloud particles convert these “reservoir” compounds into reactive free radicals (Cl and ClO). The clouds can also remove NO2 from the atmosphere by converting it to nitric acid, which prevents the newly formed ClO from being converted back into ClONO2.
The role of sunlight in ozone depletion is the reason why the Antarctic ozone depletion is greatest during spring.
During winter, even though PSCs are at their most abundant, there is no light over the pole to drive the chemical reactions.
During the spring, however, the sun comes out, providing energy to drive photochemical reactions, and melt the polar stratospheric clouds, releasing the trapped compounds.
Near the end of the spring break warming temperatures break up the vortex around mid-December.
As warm, ozone-rich air flows in from lower latitudes, the PSCs are destroyed, the ozone depletion process shuts down, and the ozone hole closes.
Most of the ozone that is destroyed is in the lower stratosphere, in contrast to the much smaller ozone depletion through homogeneous gas phase reactions, which occurs primarily in the upper stratosphere.
In some detail:
Extremely low stratospheric temperatures of – 80°C occur during polar nights: HNO3 and water form icy clouds which are only stable at those low temperatures and unstable at higher temperatures. On the surface of such clouds HCl and Cl O NO2 react as follow
HCl + Cl O NO2 → HNO3 + Cl2 (pure chlorine)
Cl2 is a stable molecule and does not react with ozone, but
Cl2 + sunlight → 2 Cl–
The dissociation of Cl2 into chlorine ions eventually leads to the ozone depletion which does not happen before the arctic spring and the process goes as long as the relevant reactants in the cloud are thawed and Cl– radicals come free.
The most concentrated part of the ozone layer (between 14 and 22 km above the earth) should not be affected, but downwinds transport Cl– to lower regions. Hence polar vortices can build up.
The very special conditions for large scale ozone depletion are: deep temperatures in polar nights, ice cloud formation, build-up of polar vortices, and polar sun rise.
CFCs are banned on a global scale since the Montreal Protocol in 1987. Due to the long lifetime of CFCs in the stratosphere it will take in average about 50 years until the molecules have been removed and a new ozone state of equilibrium has been established. The maximum for CFC 11 with a lifetime of 45 years has been in 1994, the one for CFC 12 with a lifetime of 100 years is “now”. CFCs have a twofold importance for the atmosphere: as a greenhouse gas with very high GWP and due to the very long lifetime.
Global Warming Potential (GWP) is defined as the relative radiative effect of a given substance compared to CO2 integrated over a time horizon.