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5…Photochemistry of the Stratosphere

5…Photochemistry of the Stratosphere

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5 Atmospheric Photochemistry


Northern Hemisphere, but October in the Southern Hemisphere, when the equatorial photodissociation rate will have decreased. Whilst ozone will be produced

over the whole of the stratosphere, the spring-time accumulation towards the poles

is presumed to be due to a net transfer, over a period of months from the ‘hot’ low

latitude, but high altitude region of the stratosphere, to the ‘cold’ high latitude, but

low altitude regions at the poles.

The steady-state ozone concentration at any point is determined by the balance

of O2 photodissociation (5.1) and the removal mechanisms represented by reactions (5.22) and (5.23), and also the rates of transport in and out of the gas volume

of interest, all of which vary depending on the conditions. The chemical lifetime

(see Sect. 5.7.1 for a formal definition) of O3 is very short at the top of the

stratosphere, because of the high rate of its photodecomposition (5.22), and steadystate levels are relatively low, but it increases towards lower altitudes and higher

latitudes, where it can be several months.

The vertical concentration profile increases significantly when passing down

from the mesosphere into the stratosphere, reaching a maximum between 30 and

20 km (see Fig. 5.5), tending towards the higher z when close to the equator.

Fig. 5.5 Example of the vertical profile of ozone concentration; these data (averaged and

smoothed) are extracted from Ref. [7] and represent balloon based measurements at 57o N and

107o W, during August in the years 1998–2004


R. S. Mason

It rapidly decreases across the tropopause, when the necessary UV levels have

become effectively quenched. Spring-time levels in the stratosphere reach up to

1013 molecules cm-3 at a mixing ratio of *5 ppm (e.g. if z = 20 km, where

p = 55.3 mbar and T = 220 K). This compares with background ozone, at sea

level in the mid-latitudes, which can be in the range of 20–45 ppb [6] (*5 9 1011

cm-3, pair = 1 atm, T = 293 K). The source of most ozone at ground level,

however, is different to that in the stratosphere.

The mechanism for net removal of ozone in the stratosphere is that of the

catalytic chain reaction:

X ỵ O3 ! XO ỵ O2


XO ỵ O ! X ỵ O2


where X is thought to be mainly any one of the various trace gases present: H, OH,

NOx, Cl, and Br. Cl and Br are particularly effective, one atom possibly removing

up to 104 O3 molecules overall. The direct reaction (5.23) may contribute, but it is

slow by comparison.

As has been discovered over the past 40 years, a principal contributor is

X = Cl, and the principal source of Cl is chlorofluorocarbon pollutants. These

entirely man-made materials, developed as highly efficient refrigerants, aerosol

propellants and solvents, are inert in the troposphere. They are either gases or have

a high vapour pressure and therefore all of them, unless deliberately destroyed,

find their way into the atmosphere, and being highly stable, they ultimately end up

in the stratosphere, where they do become vulnerable to photodecomposition, for


CF2 Cl2 ỵ hmk\460 nmị ! CF2 Cl ỵ Cl


Levels of Cl from these anthropogenic sources now heavily outweigh natural

sources (the most abundant of which is biogenic CH3Cl, with a global mixing ratio

at 550 ppt [8]). Their overall effect therefore is to remove ozone; the chemistry

involved is, however, highly complex, since all these species react with many

other species present, which may themselves react with ozone or even promote its

formation. Of particular note is the formation of ‘reservoir’ species. For the

X = Cl reactions, an important example is HCl. It forms via the sequence:

Cl ỵ CH4 ! CH3 þ HCl


and is eventually regenerated later as Cl by the reaction

OH ỵ HCl ! H2 O ỵ Cl


Whilst in the HCl form it is inactive towards ozone, and therefore this sequence

of reactions acts as a ‘holding’ mechanism, delaying the release of Cl, rather than

necessarily actually permanently removing it from the mix. It is thought that as

much as 70 % of stratospheric Cl is held in this form. Other ‘reservoir’ processes

trap ClO, which temporarily forms HOCl and ClONO2, in the reactions:

5 Atmospheric Photochemistry


ClO ỵ HO2 ! HOCl ỵ O2


ClO ỵ NO2 þ M ! ClONO2 þ M


HO2 is formed by the reaction:

H ỵ O2 ỵ M ! HO2 ỵ M


OH ỵ O ! H ỵ O2


and H by the reaction:

OH is formed from O(1D), mainly by its reactions with water and methane:

O 1 D ỵ H2 O ! OH ỵ OH



O 1 D ỵ CH4 ! OH ỵ CH3

X = NOx is another important catalyst. The main source is the reaction sequence:

O 1 D ỵ N2 O ! 2NO


HO2 ỵ NO ! OH ỵ NO2


NO2 ỵ hm ! NO ỵ O


N2O is an otherwise relatively unreactive trace gas, whose main source is

biogenic emission. Each of the reagent species (OH, HO2, H, NO2) themselves

undergo many other reactions besides those shown, so that there is an intricate web

involved in the final balance of O3 in the stratosphere as a function of height,

season, time of day and location. Model predictions of ozone depletion, not surprisingly, are very sensitive to the accuracy of laboratory kinetic data and the

assumptions and approximations applied to these systems.

All the gases are held in a more or less ‘steady’ state, in which the net sink

processes are thought to be removal via their longer lived reservoir species (HCl,

HOCl, and ClONO2 from above, but, for example, HNO3, HBr, BrONO2, and

HO2NO2 may also be involved). These are all water soluble and it is thought that

they eventually cross the tropopause where they dissolve in clouds and other

aerosols and get ‘rained’ out of the atmosphere.

5.5.1 Heterogeneous Chemistry

Aerosols are also important in stratospheric photochemistry. They are thought to

be formed as a result of the oxidation of dimethyl sulphide in the troposphere. This

is formed from decaying organic matter (e.g. oceanic algae) and is emitted into the

air where it breaks down to form carbonyl sulfide: COS. This is chemically stable


R. S. Mason

in the troposphere (s = *7 years) and therefore is the most abundant atmospheric

sulfur compound to be found. It therefore penetrates up into the stratosphere,

where it becomes photolysed and oxidised to SO2, eventually forming H2SO4.

COS ỵ hv ! CO ỵ S


S ỵ O2 ! SO ỵ O


SO ỵ O2 ! SO2 ỵ O


OH ỵ SO2 ỵMị ! HSO3 ỵMị


HSO3 is converted to H2SO4, possibly via reactions:

HSO3 ỵ O2 ! HO2 ỵ SO3


SO3 ỵ H2 O ! H2 SO4


This highly hygroscopic molecule readily combines with water molecules to

form an acid aerosol droplet. Other aerosols are formed by nucleation around

mineral particles injected as a result of volcanic activity. Under very cold conditions, such as at the poles in winter, these aerosols freeze to form polar stratospheric ice clouds (PSCs), the surfaces of which provide a substrate for important

heterogeneous catalytic processes. An example of this is the well-known ‘ozone

hole’ effect. This arises because the steady state concentration of O3 is sustained

by the series of reactions (5.1) and (5.21)–(5.25). As already discussed, the sink

mechanism (5.24) and (5.25) requires the presence of catalyst X, of which Cl

atoms are nowadays the most important, and which are provided, such as reaction

(5.26), mainly by the photolysis of CFCs present at trace levels in the upper

atmosphere; and much of the Cl is temporarily locked up into the reservoir species

such as HCl and ClOx.

In the winter time, air from the equator, carrying the reacting ozone mixture,

becomes trapped into a freezing cold mass (vortex) circling the pole, almost

isolated from air in the lower latitudes, until spring time sunshine arrives. In the

presence of PSCs HCl gets adsorbed onto the surface hence removing Cl from the

gas phase and thus halting the O3 removal process, which therefore remains at its

steady-state level. It appears that a number of surface reactions involving adsorbed

reservoir species occur, examples of which are:





the products of which are retained until the onset of spring (i.e. October in the

southern hemisphere, where the effect is most prominent). With the appearance of

sunlight, evaporation occurs, releasing ‘large’ quantities of Cl2. This is very easily

photolysed by visible light, the intensity of which is very much greater than the

UV normally required for photodecomposition.

5 Atmospheric Photochemistry


Cl2 ỵ hvk

410 nmị ! 2Cl


Cl atoms react with ozone to form ClO:

Cl ỵ O3 ! ClO ỵ O2


as part of the chain reactions (5.24) and (5.25); but the important reaction in the

cold conditions of the Antarctic region is thought to be the formation of the (ClO)2

dimer. Aircraft based UV measurements show a direct correlation between O3

depletion and ClO formation during ‘ozone hole’ formation. The low temperatures

favour three body reactions such as (5.48) and (5.50), and (ClO)2 is easily photolysed to yield two Cl atoms leading to an increased removal rate for ozone.

ClO ỵ ClO ỵ M ! ClOị2 þ M


ðClOÞ2 þhm ! Cl þ ClOO


ClOO þ M ! Cl ỵ O2 ỵ M


Cl ỵ O3 ! ClO ỵ O2 Â 2


The net reaction is equivalent to: 2O3 ? hm ? 3O2.

There are also lesser contributions from other catalytic species, such as BrO and

NOx, which are involved in similar reaction cycles.

So, whereas for the winter and early spring the ozone cycle has lain dormant,

the sink mechanism is triggered with the production of ‘large’ quantities of Cl

atoms over a short period of time in the later spring, and O3 concentrations rapidly

decline by as much as one half over a few weeks (the so-called ozone hole). Of

course the status quo is restored when the polar and lower latitude air masses start

to mix effectively again. This effect did not appear to become significant (and was

not observed until the early 1980s), until Cl catalysis became the dominant

removal mechanism due to the build up of man-made CFCs in the atmosphere.

5.5.2 Ozone Levels and Life on Earth

There are two critical outcomes to the patterns of upper atmosphere ozone levels,

which are vital to life on Earth. The first is the vertical layering of the atmosphere

as already described. The second is the fact that, after O2, the O3 filters out

virtually all the remaining UV-B (280–315 nm) from the solar spectrum at the

Earth’s surface. The DNA damage inflicted by extraterrestrial levels of UV would

severely restrict the survival of life forms at the surface: it is O2 combined with O3

which provides the essential UV filter.


R. S. Mason

5.6 Photochemistry of the Mesosphere

Like the stratosphere, the mesosphere photochemistry is dominated by O2, O3, and

O atom reactions. Being at a much lower pressure, the absolute concentrations are

much lower, but the ratio of [O]/[O3] is higher particularly because of the

increased rate of reaction (5.22). A distinguishing feature of this reaction is the

formation of the singlet D O atom and the singlet delta O2 molecule, (5.52).

O3 ỵ hmk\310 nmị ! O 1 D þ O2 ð1 Dg Þ


Both products are metastable, with long radiative lifetimes (110 s and 44 min

respectively). The O(1D) atom is however very rapidly converted, reaction (5.20),

by collision with air molecules into the ground state O(3P), whereas O2(1Dg) (often

called singlet oxygen) is collisionally much more stable. As a result, during the

daytime, there is an accumulation of this excited molecule in the lower reaches of

the mesosphere and the upper stratosphere, reaching values as high as 1010 molecules cm-3 (mixing ratio is at the ppm level). Despite its ‘forbidden’ nature, the

O2(1Dg ? 3Rg ) transition does occur, at 1270 nm, and forms the strongest

emission in the daytime airglow (i.e. the dayglow), and is known as the Infrared

Atmospheric Band. It is observable spectroscopically because of the very deep

optical path lengths (tens of km) of the emission volumes which can be sampled

during measurement. (Emission from the O2(1Dg ? 3Rg ) transition can also be

experimentally measured in the laboratory — see Chaps. 14 and 15).

At night there is a nightglow, but the emissions come from the top of the

mesosphere, the thermosphere and ionosphere. The IR Atmospheric Band features,

but so also do many chemiluminescent reactions due to, for example:

O ỵ O ! O2 ! O2 ỵ hm


NO ỵ O3 ! O2 ỵ NO2 ! NO2 ỵ hm


NO ỵ O ! NO2 ỵ hm



S ! O 1 D ỵ hmk ẳ 558 nmị


2 ỵ e !O P ỵ O



which emit in the visible part of the spectrum, and so can be observed by eye from

space, as a glow on the edge of the night time atmosphere. Reactions (5.53)–(5.56)

obviously also occur during the daytime in the meso- and stratospheres, but their

emissions are impossible to detect by eye against the high intensity background

solar radiation. The chemiluminescences can however be used spectroscopically to

detect and monitor the gases involved.

These emissions of light occur over the whole upper atmosphere and should be

distinguished from the aurora which are reactions induced by collisions with gas

molecules at higher altitude, of charged cosmic particles captured in the Earth’s

magnetic field, and frequently visible in the night sky of the northern and southern


5 Atmospheric Photochemistry


5.7 Photochemistry in the Troposphere

5.7.1 Non-Polluted Atmosphere

There are, of course, thousands of gaseous emissions into the atmosphere, other

than N2, O2, H2O, Ar and CO2. The latter are stable gases, whose fractional

composition remains constant throughout the homosphere. But for most emissions

the degree of mixing and their transport within the atmosphere depends on the

chemical lifetime (schem) of each species. This is best defined by example, such as

the hydrogen abstraction reaction (5.57) which is one of the most important types

of reaction in the unpolluted troposphere;

OH ỵ C5 H8 ! H2 O ỵ C5 H7 :


here the hydroxyl radical reacts with a hydrocarbon (in this case: isoprene emitted

from the leaves of plants), to form water and another radical. This sets off a chain

reaction, by which the hydrocarbon is eventually removed from the atmosphere.

The rate of reaction is given by

oẵC5 H8

ẳ kẵOHẵC5 H8 ẳ k0 ẵC5 H8



where t is the reaction time and k is the rate constant. If OH is assumed to be

present at an approximately constant steady-state concentration, the rate equation

is pseudo-first order, with pseudo-first order rate constant k0 . The definition of the

lifetime of a first order reaction is the time it takes for the reactant to decay by the

fraction 1/e of its starting concentration, which is given by 1/k0 . According to the

NIST Chemical Kinetics database, k58 = 1 9 10-11 molecules-1 cm3 s-1 at

ambient temperatures. Since [OH]troposphere & 106 molecules cm-3, then

k0 = k[OH] & 10-4 s-1, and schem (isoprene) & 104 s, i.e. *2.8 h. On a calm

day (wind speed \1.6 km hr-1) the molecules from such an emission source

would therefore be carried no more than 5 km.

The main ingredients of air have a comparatively very long chemical lifetime

and their residence time (s) is therefore long. Thus it takes about 107 years to

completely replace all the N2, 2500 years for O2 and 120 years for CO2 (not

counting the continuous exchange with CO2 dissolved in the oceans, the main

reservoir for this gas; if counted the replacement time goes down to between 3 and

15 years [9]). The long-lived species therefore have plenty of time to cross the

layer boundaries (e.g. the tropopause) and to mix effectively, so that the composition of the main ingredients remains the same up to the top of the homosphere.

The great majority of gases emanate as volcanic, biogenic or radiogenic

emissions from the surface. Whilst the original (prebiotic) atmosphere was dominated by volcanic emissions (e.g. N2, CO2 and H2O), it is now almost completely

dominated by biogenic exchange, with radiogenic exceptions (such as Ar and He).

Others, such as COS (see earlier), are the products of chemical or photochemical

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