Encyclopedia of Geochemistry

Living Edition
| Editors: William M. White

Ozone and Stratospheric Chemistry

  • Slimane BekkiEmail author
  • Joël SavarinoEmail author
Living reference work entry

Later version available View entry history

DOI: https://doi.org/10.1007/978-3-319-39193-9_207-1


Ozone Concentration Ozone Layer Stratospheric Ozone Polar Vortex Lower Stratosphere 
These keywords were added by machine and not by the authors. This process is experimental and the keywords may be updated as the learning algorithm improves.


The stratosphere is the atmospheric layer comprised between 8–18 km for its lower altitudes and 40–60 km for its upper altitudes, which correspond more or less to the stratospheric ozone layer thickness. In this part of the atmosphere, intense UV radiations from the sun fuel an active and energetic photochemical reactor leading to the formation of two important by-products: the ozone molecule (O3), an oxygen allotrope, and heat that mostly originates from absorption of UV radiation by ozone. While heat causes a strong vertical stability of the air mass in the stratosphere (stratified atmosphere), large concentrations of ozone form a UV-protective shield allowing the development of life on the Earth’s surface. The intense UV radiations combined with the ozone concentration generate a distinct stratospheric chemistry where nitrogen, oxygen, and halogen compounds are highly coupled in cycles that maintain the chemical stability of this UV-protective atmospheric layer (UVs break down atmospheric molecular oxygen, leading to the formation of ozone that better absorb UV radiations than the parent molecules). In the industrialized era, emissions of halogenated compounds have destabilized this fragile chemical stability by creating “ozone-killer” molecules, reducing the thickness of the ozone layer and allowing harmful UV radiations to reach the ground. The stratosphere is also a key atmospheric compartment for the stability of the radiative budget of the Earth and thus its climate. The presence of tiny droplets of sulfuric acid in the stratosphere acts as an aerosol layer that reflects incoming solar radiation, thus cooling the Earth’s ground temperature; it also plays a significant role in the ozone budget by providing surfaces where heterogeneous chemistry can generate ozone-destroying radical species. Such acid cloud layer is exacerbated during powerful volcanic eruptions where megatons of sulfur are directly injected into the stratosphere.

Ozone Stratospheric Chemistry


This chapter introduces the main concepts about the stratospheric ozone chemistry and budget. We start by recalling briefly the history of ozone research. Then, we describe the main features of ozone distribution. The next section reviews the basic chemistry of stratospheric ozone. An analysis of the phenomenon of the so-called “ozone hole ” is presented in the fourth section. The last section is devoted to the overall perturbation of stratospheric ozone by the emissions of ozone-destroying substances and the interactions with climate change.

Historical Perspectives and Background

Ozone (trioxygen O3) was first identified at Basel in 1840 by German chemist C.F. Schöenbein from its pronounced smell following electrical discharges in the air (Schöenbein, 1840). The same particular smell, characteristic of ozone, can be found near photocopy machines when it has been used intensively. Schöenbein named this gas “ozone” from the Greek word “ozein” meaning “to smell.” The chemical nature of ozone was determined from extensive laboratory studies initiated by M. De la Rive and C. Marignac in Geneva in 1845. A. Houzeau then detected ozone in the atmosphere in 1858 after performing the first atmospheric measurements at Rouen, France. French chemist Albert Levy conducted systematic measurements of atmospheric ozone in Paris during 30 years in the late nineteenth century. A range of scientists investigated the absorption radiative properties of ozone. The important finding was that the quasi-absence of solar UV radiation (below about 300 nm) at the surface must have been due to absorption of UV by atmospheric ozone, notably the Hartley band (Cornu, 1879; Hartley, 1881). Using ozone absorption properties, French physicists H. Buisson and C. Fabry then designed a spectrograph and made the first measurements of the thickness of the ozone layer (i.e., ozone column corresponding to the vertically integrated ozone concentration) in 1913. Based on Buisson and Fabry spectrograph, the British scientist G.M.B. Dobson developed in Oxford, UK, in 1920 a spectrophotometer , which is still used nowadays to measure the ozone column (Dobson and Harrison, 1926). The first successful attempt to measure the vertical profile of ozone was led by Swiss scientist F.W.P. Götz at the start of the 1930s (Götz et al., 1934). Later, German physicists V. H. Regener started making measurements from balloons (Regener, 1964). These height-resolved ozone observations showed that the ozone concentration was peaking in the lower stratosphere, at 20–25 km. Since the late 1970s, the global distribution of atmospheric ozone has been extensively monitored with instruments on board of satellite .

Although in relatively low concentration of a few molecules per million of air molecules, atmospheric ozone is essential to sustaining life on the Earth’s surface. Indeed, by absorbing solar radiation between 240 and 320 nm, it shields living organisms including humans from the very harmful ultraviolet radiation UV-B. About 90 % of the ozone resides in the stratosphere, the region that extends from the tropopause , whose altitude ranges from 8 km at the poles to 18 km in the tropics, to the stratopause located at about 50 km altitude (Brasseur and Solomon, 2005). Stratospheric ozone is communally referred as the “ozone layer.” The remaining 10 % of the ozone is found in the troposphere , the region that extends from the surface to the tropopause. Tropospheric ozone is a major pollutant near the surface. Ozone also reacts with a number of other gases and is the main source of the hydroxyl radical, OH, that is responsible for the oxidation and removal of most trace gases including pollutants in the atmosphere. Finally, ozone absorbs not only solar radiation but also outgoing infrared radiation and so plays a significant and complex role in the Earth’s energy budget. For all these reasons, ozone is considered as one of the key species for the chemistry and radiative balance of the atmosphere (Jacob, 1999; Brasseur and Solomon, 2005).

Unlike the major atmospheric constituents, such as nitrogen, molecular oxygen, carbon dioxide, or methane, the ozone vertical profile does not exhibit a maximum at the Earth’s surface and then a uniform decrease with height. The ozone concentration is low in the troposphere except near the surface where it is formed in air polluted by human activities . Then its concentration increases sharply from the tropopause to reach a maximum in the lower stratosphere, between 15 and 20 km at high latitude winter and between 25 and 30 km in the tropics (Figure 1). Above 30 km, the ozone concentration decreases rapidly with height. Its particular vertical distribution results from the existence of a strong source of ozone in the stratosphere and not at ground level.
Figure 1

Typical vertical profile of ozone concentration (in ozone partial pressure in mPa). About 90 % of ozone resides in the stratosphere and is commonly referred as the stratospheric ozone layer. The vertical extent or thickness of the layer varies from region to region and with season. Note that in the troposphere, notably near the surface, ozone is also produced by the oxidation of hydrocarbons in the presence of high levels of nitrogen oxides; this so-called “smog” photochemistry is typical of polluted air and explained the enhanced levels of ozone found near the ground in some areas (Hegglin et al., 2014).

The absorption efficiency of solar ultraviolet radiation (240–320 nm) during its passage through the atmosphere depends on the total number of ozone molecules along the path of the sunrays. This represents the total amount of ozone in a column of air extending from the surface to the top of the atmosphere. It is calculated by integrating vertically the altitude-dependent local concentration and is commonly referred to as the thickness of the ozone layer or ozone column, often expressed in Dobson unit (Schwartz and Warneck, 1995). The bulk of the ozone column is found in the lower stratosphere, between 15 and 30 km (see Figure 1). The ozone column varies from region to region and season over the globe. On average, the total column ozone is about 300 DU, which is equivalent to a 3 mm ozone-pure thick layer if all the ozone molecules were brought down to Earth’s surface and uniformly distributed over the globe (Brasseur and Solomon, 2005).

Chapman Cycle: Ozone Chemistry in a Pure Oxygen Atmosphere

In 1930, the physicist Sydney Chapman proposed a theory that considers the balance between the formation and destruction of ozone from oxygen radicals only (Chapman, 1930). It lays the foundation of stratospheric ozone photochemistry . The sequence of selected reactions constitutes a cycle that allows developing a simple theory of stratospheric ozone in a pure oxygen atmosphere (Bekki and Lefevre, 2009).

The formation of stratospheric ozone occurs naturally by chemical reactions involving solar ultraviolet radiation (sunlight) and oxygen molecules, which make up about 21 % of the atmosphere. The production of ozone is a two-step mechanism:
$$ {\mathrm{O}}_2+ h\nu \to \mathrm{O}+\mathrm{O};\ \lambda <242\ \mathrm{nm} {J}_1 $$
$$ \mathrm{O}+{\mathrm{O}}_2\to {\mathrm{O}}_3^{*}+ M\to {\mathrm{O}}_3+{M}^{*} {k}_2 $$
The photolysis of molecular oxygen (R1) occurs only for photons at wavelengths below 242 nm, which explains why this source is mostly present in the middle and upper atmosphere. Ozone is then produced by the recombination of oxygen atoms with molecular oxygen following the three-body reaction (R2) where M, a third body (nitrogen (N2) or oxygen (O2)), acts as a stabilizing partner, absorbing the excess of energy produced during the formation of the meta-stable ozone, O3 * (Lin and Leu, 1982). This reaction is very fast and requires only a fraction of second in the stratosphere. It also causes a release of 24 kcal of heat per mole of ozone formed. This exothermic reaction explains the positive gradient temperature observed in the stratosphere.
Like molecular oxygen, ozone is a strong absorber of solar radiation and can dissociate:
$$ {\mathrm{O}}_3+ h\nu \to {\mathrm{O}}_2+\mathrm{O} {J}_3 $$
At wavelengths above 310 nm, the photolysis of ozone leads to the direct production of O(3P), referred as O to simplify the notations. At wavelengths below about 310 nm, the photolysis of ozone produces oxygen atom in the excited state O(1D), a very reactive form of oxygen atom that decays to O(3P), the ground state of oxygen atom. The oxygen atoms excited O(1D) produced by R3 is almost instantly deactivated back to its fundamental level O(3P) by collision. Although this channel is not treated explicitly here, J3 corresponds to the total photolysis rate of ozone including the excited state channel.
Finally, ozone and atomic oxygen are destroyed simultaneously by the reaction:
$$ \mathrm{O}+{\mathrm{O}}_3\to 2\ {\mathrm{O}}_2 {k}_4 $$
It is possible to establish the balance of ozone and atomic oxygen in the stratosphere from the four reactions above, considering only the oxygen species. Ignoring transport processes, the continuity equations (i.e., the time-dependent concentration change due to the sum of production and destruction terms for the species considered) describing the evolution of the O and O3 concentration are expressed by:
$$ \frac{d\left[\mathrm{O}\right]}{ d t}=2{J}_1\left[{\mathrm{O}}_2\right]+{J}_3\left[{\mathrm{O}}_3\right]-{k}_2\left[{\mathrm{O}}_2\right]\left[\mathrm{O}\right]\left[ M\right]-{k}_4\left[{\mathrm{O}}_3\right]\left[\mathrm{O}\right] $$
$$ \frac{d\left[{\mathrm{O}}_3\right]}{ d t}={k}_2\left[\mathrm{O}\right]\left[{\mathrm{O}}_2\right]\left[ M\right]-{k}_4\left[\mathrm{O}\right]\left[{\mathrm{O}}_3\right]-{J}_3\left[{\mathrm{O}}_3\right] $$
It is worth pointing out that the photolysis of ozone (R3) does not lead to a net loss of it. Indeed, the produced oxygen atom reacts immediately with O2 to reform ozone via reaction R2.
The rapid exchange between O3 and O operates via R2 and R3 throughout the troposphere and the stratosphere. Because of its short lifetime , the oxygen atom O is always in photochemical equilibrium with ozone and molecular oxygen. The assumption of steady-state for O (i.e., d[O]/dt = 0 meaning production rate = destruction rate) allows us to derive the abundance ratio between O and O3 using Equation 1,
$$ \frac{\left[\mathrm{O}\right]}{\left[{\mathrm{O}}_3\right]}=\frac{\left(2{J}_1\frac{\left[{\mathrm{O}}_2\right]}{\left[{\mathrm{O}}_3\right]}+{J}_3\right)}{k_2\left[{\mathrm{O}}_2\right]\left[ M\right]+{k}_4\left[{\mathrm{O}}_3\right]}\simeq \frac{J_3}{k_2\left[{\mathrm{O}}_2\right]\left[ M\right]} $$
The abundance ratio between O and O3 can be simplified because the photolysis of O3 (J 3) is so much faster than photolysis of O2 (i.e., J 3 >> 2J 1[O2]/[O3]) and the rate of (R4) is several orders of magnitude slower than the rate of (R2).

The relative abundance of O compared to O3 increases rapidly with altitude. It is, however, only above about 60 km that the ratio [O]/[O3] becomes greater than 1. Below 60 km, the concentration of O3 is greater than that of O by several orders of magnitude. In these circumstances, it is useful to combine O and O3 to define a hypothetical species Ox, called the odd oxygen family, with Ox = O + O3.

Since [O] << [O3] in the troposphere and stratosphere, [Ox] ~ [O3] and the evolution equation of O3 can simply be derived by summing Equations 1 and 2 in conjunction with Equation 3:
$$ \frac{d\left[{\mathrm{O}}_3\right]}{ d t}+\frac{d\left[\mathrm{O}\right]}{ d t}\approx \frac{d\left[{\mathrm{O}}_3\right]}{ d t} = 2{J}_1\left[{\mathrm{O}}_2\right] - 2{k}_4\left[\mathrm{O}\right]\left[{\mathrm{O}}_3\right]=2\left({J}_1\left[{\mathrm{O}}_2\right]-\frac{k_4{J}_3{\left[{\mathrm{O}}_3\right]}^2}{k_2\left[{\mathrm{O}}_2\right]\left[ M\right]}\right) $$
The steady-state ozone concentration (\( \frac{d\left[{\mathrm{O}}_3\right]}{ d t}=0 \)) can then be derived from Equation 4,
$$ \left[{\mathrm{O}}_3\right]=\left[{\mathrm{O}}_2\right]{\left(\frac{k_2\left[ M\right]{J}_1}{k_4{J}_3}\right)}^{\frac{1}{2}}={x}_{{\mathrm{O}}_2}\left[ M\right]{\left(\frac{k_2{J}_1\left[ M\right]}{k_4{J}_3}\right)}^{\frac{1}{2}}\ \mathrm{with}\ {x}_{{\mathrm{O}}_2}=\frac{\left[{\mathrm{O}}_2\right]}{\left[ M\right]} $$
The rate of molecular oxygen photolysis (i.e., \( {J}_1\Big) \) increases with altitude, while [M] decreases at the same time. As a result, the ozone concentration profile calculated from Equation 5 shows a maximum in the stratosphere, which is consistent with observations.

Global Stratospheric Ozone Chemistry

Although, the Chapman’s gives approximately the correct altitude distribution, Equation 5 overestimates observed ozone concentrations by a factor of 2. Note that the transport only redistribute the ozone spatially, it is neither a sink nor a source of ozone. Indeed, global mass balance budgets of ozone have shown that loss processes does not balance the production rate, so there should be missing loss processes not taken into account in the chapman model (Johnston, 1975). This discrepancy is the most compelling argument to expend the chapman reaction scheme, involving other trace constituents that act as ozone destroyer, notably hydrogen radicals (Bates and Nicolet, 1950; Hampson, 1964), nitrogen oxides (Crutzen, 1970; Johnston, 1971), and chlorine radicals (Stolarski and Cicerone, 1974; Molina and Rowland, 1987). Note that P.J. Crutzen, M. Molina, and F.S. Rowland were awarded the Nobel Prize in 1995 “for their work in atmospheric chemistry, particularly concerning the formation and decomposition of ozone.

These radicals destroy ozone via chemical catalytic cycles . The simplest catalytic cycles are of the form:

Cycle X
$$ X + {\mathrm{O}}_3\to X\mathrm{O}+{\mathrm{O}}_2 {k}_5 $$
$$ X\mathrm{O}+\mathrm{O}\to X+{\mathrm{O}}_2 {k}_6 $$
$$ \mathrm{n}\mathrm{e}\mathrm{t}:\mathrm{O}+{\mathrm{O}}_3\to 2\ {\mathrm{O}}_2 $$
where X is the minor constituent that is the catalyst of the cycle of reactions.
The net overall reaction, reflecting the simultaneous disappearance of odd oxygen impairs, is equivalent to the reaction R4. However, the reactions involved in the cycle above are faster and are initiated by a catalyst, which, by definition, is not consumed in the process. One can then understand how a radical X with a concentration much lower than that of ozone (in practice, up to a million times less) can destroy a large number of O3 molecules until the radical is converted into less reactive molecules. The rate of destruction of O3 by the cycle X is determined by the rate of reaction XO + O. Indeed, it is much slower that the reaction X + O3, which reform XO very quickly. The destruction rate of ozone by the cycle above can be written as:
$$ \frac{d\left[{\mathrm{O}}_x\right]}{ d t} = - 2{k}_6\left[ X\mathrm{O}\right]\left[\mathrm{O}\right] $$
where k 6 is the rate of the reaction between XO and O.

The factor 2 stems from the fact that two molecules of O3 are destroyed at each pass through the cycle. In the stratosphere, X corresponds to H, OH, NO, Cl, or Br. The relative importance of these species depends on their concentrations and the reaction rates of their specific catalytic cycles. Whatever the catalyst X involved, it can be noted that the general cycle above requires the presence of atomic oxygen O. Therefore, this type of process is very effective above 30 km altitude, when the concentration of O becomes sufficient. At lower altitudes, the destruction of odd oxygen also occurs through other types of catalytic cycles, which do not require atomic oxygen. For example, one of the most effective catalytic cycles in the lower stratosphere involves a coupling bromine and chlorine radicals.

The ozone-destroying radicals mainly originate from the stratospheric oxidation of source gases emitted at the Earth’s surface. Most trace gases are quickly destroyed in the troposphere by oxidation and/or dissolution in clouds . However, certain gases have sufficiently long lifetimes in the troposphere (i.e., removed very slowly in the troposphere) for them to be transported in significant quantities into the stratosphere through air motions. The sources of stratospheric ozone-destroying hydrogen radicals (OH, HO2) are water vapor and methane emitted at the surface. The source of stratospheric nitrogen oxides radicals (NO, NO2) is the oxidation of nitrous oxide (N2O) emitted at ground level by biological processes or linked to human activities. The only significant natural source of stratospheric chlorine is the oxidation of a chlorinated organic compound, methyl chloride CH3Cl, which is emitted by the oceans. On the top of this natural source, there have been industrial sources of chlorinated organic compounds, typically chlorofluorocarbons (CFCs), for over 50 years. Almost all of these industrial compounds emitted at the surface enter the stratosphere because their lack of chemical reactivity results in very long lifetimes in the troposphere, up to 100 years (Brasseur and Solomon, 2005).

Polar Ozone

The total column ozone has been regularly measured for over half a century in the polar regions . In Antarctica , the observations until the late 1970s show an austral spring maximum, with an average over the month of October of about 350 Dobson Units (DU). From the early 1980s, a rapid and continuous decline in total column ozone is observed every spring above Antarctica. The Antarctic monthly mean thickness of the ozone layer in October is currently close to 200 DU. Local minima below even 100 DU have sometimes been seen since the late 1990s, which represents a decrease of nearly 70 % of the total ozone column compared with measurements taken twenty years earlier. The analysis of individual vertical ozone profiles (Figure 1) shows a destruction even more pronounced with often an almost complete disappearance of ozone in the lower polar stratosphere, between around 15 and 20 km (Brasseur and Solomon, 2005).

Satellite measurements have revealed the geographical extension of this phenomenon, known as the “ozone hole,” with the reduction in ozone taking place over a wide circular area centered on the South Pole (Figure 2). This area indicates the polar vortex, an enormous ring of circling fast-moving air. Although of large magnitude, it is important to note that this now recurring feature of the southern Antarctic spring remains a seasonal phenomenon. The ozone destruction indeed starts at the end of August, when sunlight comes back on the edge of the polar vortex after the winter polar night. The rate of ozone decrease peaks in September and minimum values for total column are usually reached in October. Antarctic ozone columns remain low in November. Then, they quickly recover values close to 1970s climatology at the end of December.
Figure 2

Total ozone values are for high southern latitudes on 14 September 2013 as measured by a satellite instrument. The dark blue and purple regions over the Antarctic continent show the severe ozone depletion or “ozone hole” now found during every austral spring. The ozone hole area is usually defined as the geographical area within the 220-DU contour (see white line) on total ozone maps (Hegglin et al., 2014).

When the ozone hole was discovered by a team from the British Antarctic Survey in 1987 (Farman et al., 1985), a rise in atmospheric chlorine content was proposed in an attempt to explain a phenomenon that, at that time, was totally unexpected. Indeed, as a result of human activities, the atmospheric concentration of chlorine has indeed grown almost exponentially since the mid-twentieth century, resulting in 5-fold increase compared to preindustrial values. However, the large spring polar ozone depletion cannot be accounted for by the catalytic cycles of destruction discussed previously because they are not effective enough in the lower stratosphere due to the lack of atomic oxygen. An unusual chemistry operates in the polar regions, during winter and spring, that makes ozone very sensitive to the atmospheric chlorine content. This particular chemistry is linked to the existence of chemical reactions occurring on the surface of clouds that form during winter and beginning of spring in the Antarctic lower stratosphere. This heterogeneous chemistry on polar stratospheric clouds converts chlorine compounds, communally called reservoir species (HCl, ClONO2), into ozone-destroying chlorine radicals such as the chlorine monoxide (ClO) (McElroy et al., 1986; Solomon et al., 1986). The resulting high levels of chlorine radicals greatly increase the probability that ClO reacts with itself to form a complex, the dimer Cl2O2 (Molina and Molina, 1987). This three body reaction initiates a powerful ozone-destroying catalytic cycle in the winter polar regions that is responsible for 50–70 % of the observed ozone destruction:

Cycle dimer ClO
$$ \mathrm{C}\mathrm{l}\mathrm{O}+\mathrm{C}\mathrm{l}\mathrm{O}+ M\to {\mathrm{Cl}}_2{\mathrm{O}}_2+ M {k}_7 $$
$$ {\mathrm{Cl}}_2{\mathrm{O}}_2+ h\nu \to \mathrm{C}\mathrm{l}+{\mathrm{Cl}\mathrm{O}}_2 {J}_8 $$
$$ {\mathrm{ClO}}_2+ M\to \mathrm{C}\mathrm{l}+{\mathrm{O}}_2+ M {k}_9 $$
$$ 2\left(\mathrm{Cl}+{\mathrm{O}}_3\to \mathrm{ClO}+{\mathrm{O}}_2\right) {k}_{10} $$
$$ \mathrm{n}\mathrm{e}\mathrm{t}:\ 2{\mathrm{O}}_3\to 3{\mathrm{O}}_2 $$
No atomic oxygen is involved in this cycle, and it needs little sunlight to operate effectively because the photolysis of the chlorine dimer (Cl2O2), its rate-limiting step, is fast. The chlorine dimer may also be thermally decomposed to give back ClO. However, at the low temperatures encountered in the polar vortex, the thermal decomposition of Cl2O2 can usually be neglected against its photolysis. Assuming quasi-steady state for Cl2O2 (destruction = production), the rate of ozone destruction can be expressed:
$$ \frac{d\left[{\mathrm{O}}_3\right]}{ d t}=-2{J}_8\left[{\mathrm{Cl}}_2{\mathrm{O}}_2\right]=-2{k}_7\left[\mathrm{ClO}\right]\left[\mathrm{ClO}\right]\left[ M\right] $$
where the factor 2 is explained by the number of ozone molecules destroyed for each molecule of Cl2O2 photolyzed.

This expression shows that the rate of ozone destruction is proportional to the square of the concentration of chlorine. It is estimated that the levels of chlorine in the stratosphere have tripled since the 1960s, resulting in a factor 9 increase in polar ozone destruction rate by the ClO dimer cycle. This quadratic dependence largely explains the relatively sudden appearance of the phenomenon of the ozone hole. Due to the large increase in ClO levels, the cycle that couples ClO and BrO becomes important in the polar regions. This cycle is responsible for 30–50 % of the overall seasonal polar ozone destruction (Bekki and Lefevre, 2009). Along with the increase in atmospheric chlorine content, its role has also risen rapidly in recent decades as a result of increasing emissions of long-lived anthropogenic bromine-carrying compounds into the atmosphere. The destruction of polar ozone described above continues until the warming of the polar vortex during spring which leads to the evaporation of polar stratospheric clouds and hence disappearance of heterogeneous chemistry. Without it, chlorine radicals are converted back into reservoirs species during spring.

Past and Future Evolution of Ozone at Global Scale

The balance of stratospheric ozone is controlled by hydrogen, nitrogen, chlorine, and bromine radicals, whose levels in the stratosphere are themselves determined by the levels of gas sources emitted at ground level. The main gas sources are water vapor and methane (CH4) for hydrogen compounds, nitrous oxide (N2O) for nitrogen compounds, chlorofluorocarbons (CFCs) for chlorinated compounds, and methyl bromide (CH3Br) and halons for brominated compounds. The most significant factor for the evolution of the ozone layer during the last century has been the considerable increase in industrial emissions of long-lived chlorine and bromine compounds. These gases are decomposed in the stratosphere under the influence of intense ultraviolet radiation , releasing the atoms of chlorine and bromine that destroys ozone through catalytic cycles. Because of the widespread use of chlorofluorocarbons in the industry, the stratospheric chlorine content has risen almost exponentially since the 1960s. As a result of growing emissions of halons , the stratospheric bromine content has also increased during the recent decades, although to a lesser extent than the chlorine content (Brasseur and Solomon, 2005; WMO, 2014).

The measurements unambiguously show that the change in atmospheric composition since the 1980s has led to a reduction in global stratospheric ozone. The decline in total column averaged over the year has reached about 3–4 %, a value far greater than the 1 % natural variations of ozone. The amplitude of long-term changes in the ozone layer varies with latitude and season. The greatest reduction in the total ozone column occurs at high latitudes in the southern hemisphere (−15 to −25 % on annual averages) as a result of the Antarctic ozone hole that forms each austral spring. In the northern hemisphere , the decrease is also the most pronounced at high latitudes (−4 % to −6 % on annual averages) and again mostly by some ozone depletion inside the arctic polar vortex during winter and spring. The ozone decrease observed at mid-latitudes largely results from transport of ozone-depleted air from polar regions, as well as direct destruction of ozone by global catalytic cycles that do not require heterogeneous chemistry on polar stratospheric clouds, Finally, ozone has not changed significantly in tropical regions .

Concerns arising from the discovery of the Antarctic ozone hole in 1985 and from a possible alteration of the stratospheric ozone layer by human activities have led to a process of regulating the production and emission of long-lived compounds that are sources of chlorine and bromine in the stratosphere. The Montreal Protocol (1987) was negotiated and ratified by producing countries to conduct a concerted reduction in emissions that would ultimately lead to elimination of these products in the industry . The protocol has been amended several times since 1987 by adopting increasingly stringent measures and extending its scope to a larger number of chlorine and bromine-carrying gases. Developed countries have halted production of halons in 1994 and of the most dangerous chlorinated compounds (CFCs , methyl chloroform CH3CCl3, and carbon tetrachloride CCl4) in 1996. In most applications, CFCs are now replaced by hydrochlorofluorocarbons (HCFCs) that have much shorter lifetimes in the lower atmosphere or hydrofluorocarbons (HFCs) that contain no more chlorine atoms. Implementation of the Montreal Protocol has led to a stabilization of total chlorine content since mid-1990s. The effectiveness of the Montreal Protocol is now well established with the atmospheric burden of CFC-11 (lifetime of about 45 years) declining since 1996, while that the burden of CFC-12 (lifetime of 100 years) has stabilized. Given these findings and the most probable scenarios of production, the most likely return to the atmospheric chlorine contents to levels comparable to those that preceded the appearance of the ozone hole is not expected to occur before the middle of this century (WMO, 2014).

The slow decay of stratospheric halogen levels should lead to a gradual reduction of the large polar ozone losses and an overall increase in the thickness of the ozone layer. It is worth pointing out that although the chlorine content by the end of the century is expected to be comparable to that of 1960s, before the appearance of the ozone hole, the atmospheric composition will nevertheless be very different from that of 1960s. Indeed, human activities since 1960s have led to an important increase in the concentrations of greenhouse gases such as CO2, CH4, and N2O, which will carry on during this century. Although the magnitude of this future increase is now difficult to quantify, it is already well established that the growing emissions of greenhouse gases has led to a warming of the Earth’s surface and the lower atmosphere, accompanied by a cooling of the stratosphere. In the upper stratosphere, this cooling should result in increase in ozone levels because the efficiency of the catalytic ozone destruction decreases when the temperature decreases. In the lower stratosphere, on the contrary, lower temperatures will promote more frequent formation of stratospheric clouds and therefore will tend to counteract the effects of the chlorine reduction on the destruction of polar ozone in winter and spring. Our ability to predict the long-term evolution of temperature in the stratosphere has therefore become a key factor in determining the future of the ozone layer. In a sense, the issue of stratospheric ozone will be increasingly linked in the coming years with the issue of greenhouse gases emissions and climate change (Brasseur and Solomon, 2005; Bekki and Lefevre, 2009; WMO, 2014).

Stratospheric Sulfur Chemistry

This section is devoted to the sulfur cycle and budget in the stratosphere, notably focusing on stratospheric aerosol particles because they play a very significant role in stratospheric ozone chemistry and in the radiative balance of the atmosphere, especially after large volcanic eruptions (Hamill et al., 1997; Brasseur and Solomon, 2005; SPARC, 2006). The first section introduces the life cycle of stratospheric sulfur. We then review stratospheric sulfur oxidation chemistry and key aerosol microphysical processes. The next section presents the perturbations of the stratospheric aerosol layer by large volcanic eruptions and the effects on ozone and climate. The last section is devoted to the possibility of cooling the climate intentionally through deliberate sulfur injections in the stratosphere in order to counteract the climate warming from increasing concentrations of GHGs (the so-called sulfur geo-engineering ).

Sulfur Life Cycle

Elevated concentrations of aerosol particles are found all over the globe in the lower stratosphere, between about 10 and 35 km. It was first discovered in 1961 by Junge et al. (1961) who noted that, as balloon-borne instruments rose into the stratosphere, they registered an increase in the concentration of large particles (Hamill et al., 1997). Further measurements have shown that stratospheric aerosols consisted primarily of submicron liquid (super-cooled) sulfate aerosols.

Figure 3 illustrates the processes that govern the stratospheric aerosol life cycle and global distribution. Most of the sulfur enters the stratosphere under the form of gases through the tropical tropopause (around 17 km in the tropics and 8 km in the polar regions). The source gases originate from emissions at the Earth’s surface or even direct injections by large volcanic eruptions. Once in the tropical stratosphere, they are transported throughout the stratosphere by the Brewer-Dobson circulation (ascending in the tropics and descending in the subtropics). During their stratospheric transit, they are oxidized and converted into sulfuric acid (H2SO4). Gaseous H2SO4 either combines with water vapor to nucleate and form new very small sulfate aerosols or condenses on pre-existing aerosol particles. Once formed, sulfate particles evolve under the action of several microphysical processes (condensation/evaporation, coagulation, sedimentation (Seinfeld and Pandis, 1998)). Particles grow by coagulation and by co-condensation of H2O and H2SO4, which is produced by the oxidation of sulfur gases. Sulfate aerosol particles evaporate completely above 35 km because of the increase of the temperatures in this region of the atmosphere. In the end, aerosols are removed from the stratosohere by their gravitational sedimentation . Once in the troposphere, they are then removed by wet (i.e., precipitating clouds) and dry (i.e., sticking to the earth’s surface) deposition .
Figure 3

Schematic of the stratospheric aerosol life cycle (SPARC, 2006 adapted from Hamill et al., 1997).

Conversion of Reduced Sulfur Gases into Sulfate

During nonvolcanic (also called background) conditions, the main precursor of stratospheric aerosol particles is thought to be carbonyl sulfide (OCS), the only long-lived sulfur species (Crutzen, 1976). Apart OCS, all the reduced sulfur compounds are rapidly oxidized in the troposphere and thus, once emitted at the surface, they have little time to reach the stratosphere before being removed. In contrast, the vast majority of OCS emitted at the surface, predominantly from natural sources, reaches the stratosphere. It is worth pointing out that other nonvolcanic sulfur dioxide (SO2) sources and tropospheric sulfate particles might also be important (Brock et al., 1995; Sheng et al., 2015), especially in the tropics where the vertical transport of air from the surface to the tropical tropopause can be rapid. Once it is in the stratosphere, OCS is photolyzed to give SO2:
$$ \mathrm{O}\mathrm{C}\mathrm{S}+ h\nu \to \mathrm{S}+\mathrm{C}\mathrm{O} {J}_{11} $$
$$ \mathrm{S} + {\mathrm{O}}_2\to {\mathrm{SO}}_2 {k}_{12} $$
Then the oxidation of SO2 is initiated by the hydroxyl radical OH
$$ {\mathrm{SO}}_2+\mathrm{O}\mathrm{H}+ M\to {\mathrm{HSO}}_3+ M {k}_{13} $$
with M being an air molecule.
R15 is very rapidly followed by the following reactions:
$$ {\mathrm{HSO}}_3 + {\mathrm{O}}_2\to {\mathrm{SO}}_3+{\mathrm{HO}}_2 {k}_{14} $$
$$ {\mathrm{SO}}_3 + {\mathrm{H}}_2\mathrm{O}\to {\mathrm{H}}_2{\mathrm{SO}}_4 {k}_{15} $$
The limiting step in the SO2 oxidation sequence R13R14R15 is the reaction with OH (R13). The e-folding lifetime of SO2 in the stratosphere is about 1 or 2 months. The net effect of the sequence can be summarized by Bekki (1995),
$$ \mathrm{N}\mathrm{e}\mathrm{t}:{\mathrm{SO}}_2+\mathrm{O}\mathrm{H} + {\mathrm{O}}_2 + {\mathrm{H}}_2\mathrm{O} - \to\ {\mathrm{H}}_2{\mathrm{SO}}_4+{\mathrm{H}\mathrm{O}}_2 $$
The sequence converts OH into HO2 and consumes one molecule of H2O. When SO2 levels are high, for instance following an injection of SO2 by a volcanic eruption , SO2 conversion to sulfate particles is expected to reduce levels of OH, the main atmospheric oxidant.
As gaseous H2SO4, the end oxidation product, has a very low saturation vapor pressure at the temperatures prevailing in the lower stratosphere, it quickly combines with H2O to form new very small sulfate particles via binary homogeneous nucleation , typically at the tropical tropopause or in the polar regions. However, in most of the lower stratosphere, where there are quite a few particles available, gaseous H2SO4 tends to co-condense with H2O on pre-existing particles and make them grow. The average composition of stratospheric sulfate aerosols is about 75 % by weight of H2SO4, which is roughly equivalent to a molecular composition of one molecule of H2SO4 for two molecules of H2O. As the result, the H2SO4 conversion into sulfate can be expressed by,
$$ {\mathrm{H}}_2{\mathrm{SO}}_4 + 2{\mathrm{H}}_2\mathrm{O}\ \to\ \mathrm{Sulphate} $$
Aerosols also grow via coagulation. Coagulation is the process of two aerosols of various sizes colliding with each other and combining into a single larger particle (Kremser et al., 2016). Finally, sulfate particles are mostly removed from the stratosphere by gravitational sedimentation . The sedimentation velocity of a particle is very strongly dependent on its size. When the sedimentation of a population of particles is considered, the critical characteristic is the particle size distribution , i.e., concentration of particles as a function of size. Size distribution is also critical for a range of properties such as the amount of surface area available for heterogeneous chemical reactions or for their scattering efficiency of sunlight (a key property for their cooling effect on climate; (Seinfeld and Pandis, 1998; Jacob, 1999; SPARC, 2006)).
At the top of the stratospheric aerosol layer (~30-35 km), sulfate aerosols also evaporate and release H2SO4 in the gas-phase. There are indications that H2SO4 can then be photolyzed by visible sunlight to give SO3, which in turn is also subject to photolysis and give back SO2 (Kremser et al., 2016).
$$ {\mathrm{H}}_2{\mathrm{SO}}_4+ h\nu\ \to {\mathrm{SO}}_3+{\mathrm{H}}_2\mathrm{O} {J}_{16} $$
$$ {\mathrm{SO}}_3 + h\nu\ \to {\mathrm{SO}}_2 + \mathrm{O} {J}_{18} $$
As the result, the dominant sulfur species above the stratospheric aerosol layer is SO2.

Impacts of Volcanic Eruptions

Although rare, major volcanic eruptions have an important impact on the stratospheric aerosol levels and loading, as illustrated by satellite data for the last 30 years (Figure 4). They inject massive amount of sulfur, mostly under the form of SO2, directly into the stratosphere, resulting in very large enhancements in sulfate aerosol levels lasting up to several years. For example, the largest volcanic eruption of the last 50 years is the eruption of Mount Pinatubo , Philippines, in June 1991. It injected about 15–20 Mega-tons of sulfur into the stratosphere, which is more than an order of magnitude greater than the annual stratospheric sulfur mass flux into the stratosphere from tropospheric surface sources during a nonvolcanic period.
Figure 4

History of integrated backscatter from two tropical sites (São José dos Campos and Mauna Loa) and two mid-latitude sites (Hampton and Garmisch). Integrated backscattering indicates the aerosol load of the stratosphere. The dashed horizontal lines are meant only to aid the reader. The times of the most significant volcanic eruptions during the period are indicated in the top panel with the green triangles (El Chichon 1981, Mexico and Pinatubo 1991, Philippines). It can clearly be seen that these two major stratospheric eruptions are responsible for the large increase of the optical opacity of the stratosphere at the global scale, induced by the formation of sulfuric acid aerosols (Adapted from SPARC, 2006).

The drastic increases in stratospheric aerosol loading impact stratospheric ozone and climate . They influence ozone because large volcanic sulfate aerosol particles act as sites for heterogeneous reactions that lead to ozone destruction when the atmospheric levels of halogens (chlorine, bromine) are high (see ozone section). They also impact climate because sulfate aerosols scatter the incoming sunlight back to space and hence cool the surface. To illustrate these effects, let us consider the Mount Pinatubo eruption. The injection of SO2 and conversion into sulfate led to stratospheric sulfate aerosol loadings and surface area densities (aerosol surface available for heterogeneous chemical reactions) enhanced by up two orders of magnitude, resulting in enhanced heterogeneous chemistry that converts chlorine- and bromine-containing reservoir species such ClONO2 and BrONO2 into ozone-destroying radicals. Global stratospheric ozone dropped by several percentages after the Mt. Pinatubo eruption (WMO, 2010). Stratospheric aerosol optical depth increased by almost two orders of magnitude (Figure 4, bottom), enhancing the planetary albedo and hence cooling the Earth’s surface by about half degree (Kremser et al, 2016). The volcanic forcing of climate represents one of the two major natural external climate forcers; the other one is the solar variability .


Actual mitigation efforts to reduce global warming are already falling far short of what many analysts think is needed to avoid dangerous changes in climate. These failings arise from a political logic that will soon be difficult to rectify. Deep cuts may be too costly and thus politically difficult to implement, organize, and sustain; they imply radical change in energy systems that will be difficult to implement and manage effectively in many countries even if they could gain the needed political support. More importantly, it has been proven extremely difficult to design competent international institutions for coordinating and enforcing worldwide efforts to mitigate emissions (e.g., the Kyoto protocol, COP21). As a result, the possibility of modifying the climate intentionally through large-scale technical schemes (also called geo-engineering ) in order to counteract the warming effect of greenhouse gases (GHGs) has started to be considered more seriously, even by prominent scientists (Crutzen, 2006; Wigley, 2006; Lenton and Vaughan, 2009).

Climate geo-engineering is a large-scale engineering of our environment in order to counteract the effect of changes in atmospheric composition due to human activities, in particular the increase of GHGs. Most of the proposed schemes are based on the decrease of solar flux reaching the surface to balance the increase in thermal infrared radiation reaching the surface from increasing GHG concentrations. They are based on injection of gases or particles to increase the albedo (i.e., the reflectivity) of the Earth atmosphere and clouds. By analogy with the global cooling observed after volcanic emissions of sulfur in the stratosphere, it has been proposed to inject sulfur in the stratosphere in order to enhance the stratospheric aerosol loading and hence the scattering back to space of the incoming solar radiation. According to the literature review, at least mega-tons of sulfur appears to be necessary for counteracting significantly the climate warming driven by increasing GHGs with highly uncertain costs, results, and consequences, notably on stratospheric ozone (e.g., Robock et al., 2009). Geo-engineering can certainly not be considered as a long-term solution to the climate change problem and any induced perturbation in the Earth’s climate system may have unwanted consequences that should be carefully evaluated before any implementation.


While making only ≈10 % of the mass of the atmosphere, the stratosphere plays a crucial role for the stability of the chemical and radiative balance of the Earth. The presence of a rich layer of ozone protects the Earth’s ground from harmful UV radiations that otherwise would prevent the development of life on the Earth’s surface. In return, by absorbing UV radiations this ozone layer generates a temperature inversion that makes the stratosphere particularly stable dynamically in comparison of the troposphere. In some way, the ozone stratospheric layer creates its own living conditions. However, the stability of the ozone layer has been drastically challenged by emissions of human manufactured halogenated species, which can cross the tropopause barrier and result in catalytic destruction of ozone. Equally, during powerful volcanic eruptions, the formation of a highly enhanced sulfate aerosol layer in the stratosphere strongly disturbs the energy and chemical balance of the stratosphere, leading to significant decreases in stratospheric ozone concentrations. In this regard, given the importance of stratospheric ozone for climate, chemistry, and life on Earth, any geo-engineering actions that would lead to change the radiative properties of the stratosphere in order to balance global warming should be viewed with a lot of caution, at minimum with distrust and critical analysis.



  1. Bates, D. R., and Nicolet, M., 1950. The photochemistry of atmospheric water vapor. Journal of Geophysical Research, 55(3), 301–327, doi:10.1029/JZ055i003p00301.CrossRefGoogle Scholar
  2. Bekki, S., 1995. Oxidation of volcanic SO2: a sink for stratospheric OH and H2O. Geophysical Research Letters, 22(8), 913–916, doi:10.1029/95GL00534.CrossRefGoogle Scholar
  3. Bekki, S., and Lefevre, F., 2009. Stratospheric ozone: History and concepts and interactions with climate. Eur Phys J Conferences, 1, 113–136Google Scholar
  4. Brasseur, G., and Solomon, S., 2005. Aeronomy of the Middle Atmosphere – Chemistry and Physics of the Stratosphere and Mesosphere, 3rd edn. Dordrecht: Springer.Google Scholar
  5. Brock, C. A., Hamill, P., Wilson, J. C., Jonsson, H. H., and Chan, K. R., 1995. Particle formation in the upper tropical troposphere: a source of nuclei for the stratospheric aerosol. Science, 270(5242), 1650–1653, doi:10.1126/science.270.5242.1650.CrossRefGoogle Scholar
  6. Chapman, S., 1930. A theory of upper atmospheric ozone. Quarterly Journal of the Royal Meteorological Society, 3, 103–125.Google Scholar
  7. Cornu, A., 1879. Sur la limite ultra-violette du spectre solaire a diverses altitudes. Proceedings of the Royal Society of London, 29, 47–55.CrossRefGoogle Scholar
  8. Crutzen, P. J., 1970. The influence of nitrogen oxides on atmospheric ozone content. Quarterly Journal of the Royal Meteorological Society, 96(30), 320–325, doi:10.1029/JC076i030p07311.CrossRefGoogle Scholar
  9. Crutzen, P. J., 1976. The possible importance of CSO for the sulfate layer of the stratosphere. Geophysical Research Letters, 3(2), 73–76, doi:10.1029/GL003i002p00073.CrossRefGoogle Scholar
  10. Crutzen, P. J., 2006. Albedo enhancement by stratospheric sulfur injections: a contribution to resolve a policy dilemma? Climatic Change, 77, 211, doi:10.1007/s10584-006-9101-y.CrossRefGoogle Scholar
  11. Dobson, G. M. B., and Harrison, D. N., 1926. Measurements of the amount of ozone in the Earth’s atmosphere and its relation to other geophysical conditions. Proceedings of the Royal Society of London Series A, Containing Papers of a Mathematical and Physical Character, 110(756), 660–693.CrossRefGoogle Scholar
  12. Farman, J. C., Gardiner, B. G., and Shanklin, J. D., 1985. Large losses of total ozone in Antarctica reveal seasonal ClOx/NOx interaction. Nature, 315, 207–210.CrossRefGoogle Scholar
  13. Götz, F. W. P., Meetham, A. R., and Dobson, G. M. B., 1934. The vertical distribution of ozone in the atmosphere. Proceedings of the Royal Society of London, 145(855), 416–446.CrossRefGoogle Scholar
  14. Hamill, P., Jensen, E. J., Russell, P. B., and Bauman, J. J., 1997. The life cycle of stratospheric aerosol particles. Bulletin of the American Meteorological Society, 78(7), 1395–1410, doi:10.1175/1520-0477(1997)078<1395:TLCOSA>2.0.CO;2.CrossRefGoogle Scholar
  15. Hampson, J., 1964. Chemical Instability of the Stratosphere. Paper presented at the International Association of Meteorology and Atmospheric Physics (IUGG) Symposium on Atmospheric Radiation, Leningrad.Google Scholar
  16. Hartley, W. N., 1881. On the absorption of solar rays by atmospheric ozone. Journal of the Chemical Society, Transactions, 39, 111–128, doi:10.1039/CT8813900111.CrossRefGoogle Scholar
  17. Hegglin, M. I., Fahey, D. W., McFarland, M., Montzka, S. A., and Nash, E. R., 2014. Twenty Questions and Answers About the Ozone Layer: 2014 Update. Scientific Assessment of Ozone Depletion: 2014. Geneva: World Meteorological Organization.Google Scholar
  18. Jacob, D. J., 1999. Introduction to Atmospheric Chemistry. Princeton: Princeton University Press.Google Scholar
  19. Johnston, H. S., 1971. Reduction of stratospheric ozone by nitrogen oxide catalysts from supersonic transport exhaust. Science, 173, 517–522.CrossRefGoogle Scholar
  20. Johnston, H. S., 1975. Global ozone balance in the natural stratosphere. Reviews of Geophysics, 13(5), 637–649, doi:10.1029/RG013i005p00637.CrossRefGoogle Scholar
  21. Junge, C. E., Chagnon, C. W., and Manson, J. E., 1961. Stratospheric aerosols. Journal of Meteorology, 18(1), 81–108, doi:10.1175/1520-0469(1961)018<0081:SA>2.0.CO;2.CrossRefGoogle Scholar
  22. Kremser, S., Thomason, L. W., von Hobe, M., Hermann, M., Deshler, T., Timmreck, C., Toohey, M., Stenke, A., Schwarz, J. P., Weigel, R., Fueglistaler, S., Prata, F. J., Vernier, J.-P., Schlager, H., Barnes, J. E., Antuña-Marrero, J.-C., Fairlie, D., Palm, M., Mahieu, E., Notholt, J., Rex, M., Bingen, C., Vanhellemont, F., Bourassa, A., Plane, J. M. C., Klocke, D., Carn, S. A., Clarisse, L., Trickl, T., Neely, R., James, A. D., Rieger, L., Wilson, J. C., and Meland, B., 2016. Stratospheric aerosol – observations, processes, and impact on climate. Reviews of Geophysics, 54, 278–335, doi:10.1002/2015RG000511.CrossRefGoogle Scholar
  23. Lenton, T. M., and Vaughan, N. E., 2009. The radiative forcing potential of different climate geoengineering options. Atmospheric Chemistry and Physics, 9(15), 5539–5561.CrossRefGoogle Scholar
  24. Lin, C. L., and Leu, M. T., 1982. Temperature and third-body dependence of the rate constant for the reaction O + O2 + M → O3 + M. International Journal of Chemical Kinetics, 14(4), 417–434, doi:10.1002/kin.550140408.CrossRefGoogle Scholar
  25. McElroy, M. B., Salawitch, R. J., Wofsy, S. C., and Logan, J. A., 1986. Reductions of Antarctic ozone to synergistic interactions of chlorine and bromine. Nature, 321, 759–762.CrossRefGoogle Scholar
  26. Molina, L. T., and Molina, M. J., 1987. Production of Cl2O2 from the self-reaction of the CIO radical. Journal of Physical Chemistry, 91, 433–436.CrossRefGoogle Scholar
  27. Molina, L. T., and Rowland, F. S., 1987. Stratospheric sink for chlorofluoromethanes: chlorine atom catalyzed destruction of ozone. Nature, 249, 810–814.CrossRefGoogle Scholar
  28. Regener, V. H., 1964. Measurement of atmospheric ozone with the chemiluminescent method. Journal of Geophysical Research, 69(18), 3795–3800, doi:10.1029/JZ069i018p03795.CrossRefGoogle Scholar
  29. Robock, A., Marquardt, A., Kravitz, B., and Stenchikov, G., 2009. Benefits, risks, and costs of stratospheric geoengineering. Geophysical Research Letters, 36, L19703, doi:10.1029/2009gl039209.CrossRefGoogle Scholar
  30. Schöenbein, C., 1840. Recherche sur la nature de l” odeur qui se manifeste dans certaines actions chimiques. Comptes Rendus de l’Académie des Sciences, 10, 706–710.Google Scholar
  31. Schwartz, S. E., and Warneck, P., 1995. Units for use in atmospheric chemistry (IUPAC Recommendations 1995). Pure and Applied Chemistry, 67. doi:10.1351/pac199567081377Google Scholar
  32. Seinfeld, J. H., and Pandis, S. N., 1998. Atmospheric Chemistry and Physics. New York: Wiley, Vol. 1.Google Scholar
  33. Sheng, J.-X., Weisenstein, D. K., Luo, B.-P., Rozanov, E., Stenke, A., Anet, J., Bingemer, H., and Peter, T., 2015. Global atmospheric sulfur budget under volcanically quiescent conditions: Aerosol-chemistry-climate model predictions and validation. Journal of Geophysical Research, 120(1), 256–276, doi:10.1002/2014jd021985.Google Scholar
  34. Solomon, S., Garcia, R., Rowland, F. S., and Wuebbles, D. J., 1986. On the depletion of Antarctic ozone. Nature, 321, 755–758.CrossRefGoogle Scholar
  35. SPARC, 2006. SPARC Assessment of stratospheric aerosol properties (ASAP). SPARC Report No. 4. World Climate Research Programme.Google Scholar
  36. Stolarski, R. S., and Cicerone, R. J., 1974. Stratospheric chlorine: a possible sink for ozone. Canadian Journal of Chemistry, 52(8), 1610–1615, doi:10.1139/v74-233.CrossRefGoogle Scholar
  37. Wigley, T. M. L., 2006. A combined mitigation/geoengineering approach to climate stabilization. Science, 314(5798), 452–454, doi:10.1126/science.1131728.CrossRefGoogle Scholar
  38. WMO, 2010. Scientific Assessment of Ozone Depletion: 2010. Global Ozone Research and Monitoring Project – Report No. 52. Geneva: World Meteorological Organization.Google Scholar
  39. WMO, 2014. Scientific Assessment of Ozone Depletion: 2014. Global Ozone Research and Monitoring Project – Report No. 55. Geneva: World Meteorological Organization.Google Scholar

Copyright information

© Springer International Publishing AG 2016

Authors and Affiliations

  1. 1.LATMOS-IPSL (Laboratoire Atmosphères, Milieux, Observations Spatiales)Sorbonne Universités-UPMC, Paris Saclay-UVSQ, CNRSParisFrance
  2. 2.Laboratoire de Glaciologie et Géophysique de l’EnvironnementUniversité Grenoble-Alpes, UMRCNTS/UGAGrenobleFrance