the story of CFCs and stratospheric ozone

Part 3 of the course relates the story of CFCs and how an environmental catastrophe was
avoided. This section also emphasises the great sensitivity of the atmosphere to human
activity, and in turn, the great sensitivity of the 'biosphere" to changes in our atmosphere.

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As far as the atmosphere is concerned, essentially all energy received from the Sun is in the
form of electromagnetic radiation. Other sources of energy from, for example, fossil fuel
combustion or transfer from the warm Earth's core, are so small in comparison that they can
be neglected. The average amount of electromagnetic flux reaching the Earth is quite
precisely known at (1366  3) W m-2, (this is equivalent to about 14 100% efficient, 100 W
household light bulbs per square meter) perpendicular to the direction of the photons. Over
the spherical surface of the Earth, this energy averages to 342 W per square meter of the
Earth’s surface. Solar radiation, more precisely, ultraviolet radiation (wavelengths less than
400 nm), is, of course, hugely important to the chemistry of the atmosphere. Without it, the
atmosphere would be inert and any substances released from the Earth's surface would not
be removed by either chemical reaction or photo-dissociation, leading to their increasing
atmospheric concentrations and the related environmental impact. The value 1366 Wm-2 is
known as the Solar constant, S.

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The Sun behaves very much like a black-body radiator, which is an object that is able to
absorb and emit photons of all wavelengths. Its electromagnetic spectrum follows closely
that of a black body of temperature 5800 K, with the greatest deviations occurring at very
short wavelengths. 5800 K is essentially the average surface temperature of the Sun. The
peak intensity of the emission is found in the visible region close to 500 nm, which happens
to be close to the visual response peak of the human eye at 555 nm. As can be seen in this
graph, both the intensity and the shape of the spectrum of the radiation reaching the Earth’s

surface is modified by absorption of (and scattering by) several atmospheric species.
Amongst these absorbers, both O2 and O3 are prominent in the ultraviolet (UV) and visible
regions of the spectrum. The hashed area shows the total photon flux that would reach the
Earth's surface if atmospheric species did not absorb at all. The difference between the
average incoming flux of 342 W m-2 and the flux arriving at the Earth's surface is accounted
for by reflection, mostly from water clouds and other aerosols.

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Most of the atmosphere is composed of N2, O2, and H2O vapour. The ratio of the former
two is essentially constant in the troposphere, stratosphere, and mesosphere, while the
concentration of water vapour changes spatially in three dimensions ranging from a fraction
of one percent to about four percent by number of molecules per unit volume. Although
most of the interesting chemistry of the atmosphere occurs between those minor species that
make up only a small fraction of the atmosphere (orange block, above), the macroscopic
structure of the lower and middle atmosphere is governed by the interaction of O2 with
sunlight, as will be discussed shortly.
In order to consider the fate of molecules when subjected to UV and visible radiation, one
may begin by simply looking at the dissociation energies of typical molecular bonds. In the
case of diatomic molecules, the bond dissociation energy is equal to the difference in
enthalpy of formation of XY and of X + Y. N2 has one of the strongest molecular bonds
encountered in nature. The bond dissociation energy can be related to the minimum photon
energy (and hence its associated wavelength) necessary to produce N atoms: for N2, 127 nm
or less is required to dissociate N2 to N + N, assuming that there is no barrier to the
dissociation process that would mean even shorter wavelengths were required for
dissociation. According to the Sun’s emission spectrum as observed from space, relatively
few photons of wavelength less than 127 nm reach the Earth's atmosphere. O2, on the other
hand, may dissociate a longer wavelengths of up to 240 nm. In this spectral region, many
more photons are available.

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The change in light intensity as it passes through a gas can be easily described by the BeerLambert expression, which predicts an exponential decrease in intensity with distance if the
concentration of the absorbing species remains constant. Please note the different units used
for absorption and also to the fact that the product (, k or) cL is dimensionless. In this
course we will tend to use absorption-cross section with units of cm2 (per molecule) as we
will normally use concentrations with units of (molecules) cm-3.

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Absorption of light in the atmosphere is interesting for three reasons. (1) It shields animals
and plants from harmful UV wavelengths (2) it heats the atmosphere (3) it produces highly
reactive species.
In order to quantify the absorption process, one needs to know the absorption crosssection (or equivalently, the molar extinction coefficient, or molar absorption coefficient) of
each molecule and what the resulting dissociation products are, if any. Shown here is the
absorption cross-section for O2 and H2O. That of N2 is not shown, but it becomes significant
compared to the other two only below 150 nm, though it does not dissociate until 127 nm or
less, as already noted. Note, the values (and units) of the absorption cross section. The
highest is of the order of 10-17 cm2. This is considered to be a very high value for a molecule
in the atmosphere, although atoms can have much greater peak absorption cross-sections
than this. The essential physical interpretation of absorption cross-section of 10-17 cm2 is
that, according to a photon corresponding to a particular wavelength, the molecule appears
to have a surface area in the direction of the photons approach, of 10-17 cm2. Such a surface
area corresponds to a diameter of (4 x 10-17/)0.5 = 35 Ǻ, which is several times greater than
the collision diameter of O2. To put this into perspective, remember that Iabs/I0 = cL, where
 is the molecular absorption cross-section (cm-2 per molecule), c is the concentration
(molecules per cm-3), and L is the path length over which absorption occurs (cm). At ground

level, O2 has a concentration of about 5 x 1018 (molecules) cm-3. Thus for 20% absorption, L
= 0.2/(10-17x 5 x 1018) = 0.004 cm. Therefore, a layer of air of only 0.1 cm thick would
appear entirely "black" if the absorption cross section of air in the visible was 10-17 cm2
(only a fraction light of exp[-(10-17 cm2 x 5 x 1018 cm-3 x 0.1 cm)] = 0.007 would penetrate
such a layer. At the other extreme, in order to view an object through air that is 1 km away,
we can say the light absorption needs to be minimal (no more than 10 %, for example).
What should be the absorption cross section in this case? Here  = 0.9/(5 x 1018 cm2 x 1000
m x 100 cm/m) = 2 x 10-24 cm2. These figures are good to bear in mind. Naturally, one has
to consider that the concentrations of most absorbing species in the atmosphere are orders of
magnitude less than that of O2, and one must also take into account of the logarithmic

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change in pressure with altitude, which means that any given absorption path length in the
atmosphere in the vertical direction will not usually have a constant concentration.
Also shown on this figure, is the actinic flux: the spectral irradiance of the Sun directly
above the atmosphere (say at an altitude of 200 km) in linear units. Since most of the light
reaching the atmosphere is lies at wavelengths longer than the main absorption bands of O2
and H2O, a better representation of the influence of O2 and H2O on the solar spectrum is
that of a log/linear plot. This is given on the next page.


On the previous graph, it was difficult to see the very small absorption cross-sections
associated with O2 and H2O at longer wavelengths where the Sun's radiation flux begins to
increase rapidly. A logarithmic scale for the y-axis shows these more clearly. Shown also
are two examples of the range of the effectiveness of O2 in reducing the Sun's light
intensity. It can be clearly seen that at 150 nm the absorption cross-section is so large that a
path of O2 (at ground level) of only 1 mm is necessary to reduce the incident intensity by a
factor e (that is, a factor of 2.72). At 240 nm, 9 km of O2 is required to achieve the same

reduction factor.
It is quite clear then that due to absorption of O2 alone, light at 150 nm cannot reach the
Earth’s surface. Would the atmosphere be transparent at 240 nm due to absorption of O2
only? In order to answer this question one would need to take into account the exponential
changing concentration of O2 with altitude, as already mentioned. An example calculation is
given later on the page relating air pressure to altitude.

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This graph shows that absorption of radiation by O2 accounts for the removal of the Sun’s
radiation only in a limited region up to about 260 nm (the 240 nm cut-off for the
dissociation of O2, already mentioned, refers to zero Kelvin, at other temperatures there is
always a so-called “Boltzmann tail”, which is due to transitions from small populations of
vibrationally-excited levels and has the effect of extending this abrupt, zero Kelvin, cut-off
point - common the many physical phenomena). You will notice that part of the relative
DNA damage spectrum is also plotted. This spectrum is a subjective estimate of the degree
of damage induced in DNA when exposed to various wavelengths. In order to calculate the
total relative DNA damage, one must multiply the photon flux spectrum (orange) by the
relative DNA damage spectrum. Since these two quantities are approximately anticorrelated, the total DNA damage spectrum (sometimes referred to as the “DNA action
spectrum”) will have a relatively sharp peak. Under the conditions given above, the total
DNA damage, if only considering the wavelengths between 310 nm and 305 nm, would be
much beyond that which can be safely repaired by most living organisms. Clearly another
absorber is required to protect the Earth’s biosphere. This protection is provided by a
relatively minor atmospheric constituent, ozone.

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Earlier it was noted that different wavelengths will be absorbed at different rates in the

atmosphere depending on the corresponding absorption cross-section of the species
involved and their concentration. When light passes through an absorbing medium, it is also
expected that the light intensity increases exponentially with distance according to the BeerLambert Law. However, this is valid if the concentration of the absorbing species is not a
function of distance. For the atmosphere, concentration of O2 increases exponentially with
distance (from space). This has interesting consequences for the position of maximum
absorption. In the slide above three situation are considered for which light passes through a
series of slabs of increasing concentrations having fixed absorption cross-sections and
thickness. As the light passes through, the amount of light absorbed in each slab is
calculated. It can be seen that for some situations the maximum absorption can take place
somewhere in the middle of the stack of slabs. This often occurs in the atmosphere.

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In order to work out the amount of light (or number of photons in this case) absorbed by the
atmosphere, three item of information are required. (1) the absorption cross-section as a
function of wavelength (2) the initial light intensity as a function of wavelength (3) the
concentration of the absorption species as a function of distance (for this we also assume a
constant T of 250 K). It is also more convenient to simplify the absorption cross section and
the photon flux data as indicated by the red lines. In practice this means having a table of
absorption cross section values for say each nm and the same for the photon flux.

10


The rate of absorption of photons at any given altitude can be calculated by the integration
given above. Note that there is also a factor that takes into account that not all absorbed
photons lead to photo-dissociation, this is called the quantum yield. To perform an
integration one needs to known the various functions and then be able to integrate them.
This is normally not possible and numerical integration is performed as described in the

lower box. Here F(), (), and () are taken from tables. The units of photolysis rate is
(molecules) s-1. For the numerical integration one only has to consider a relevant
wavelength range that correspond to the range over which photo-dissociation may occur
and the range over which there is a reasonably high actinic flux.

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If we perform the same calculations for each wavelength, the following absorption profiles
are obtained. Here, as expected, the longer wavelengths, corresponding to lower absorption
cross-section of O2, penetrate the atmosphere much further than do the shorter wavelengths.
Also note that the peak absorptions are much greater too. This is because the initial actinic
flux increases with wavelength.
Since one photon produces two O atoms, the rate of production of O atoms by this process
is double the photolysis rate of O2.

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It is known that O-atoms react very rapidly with O2 throughout the atmosphere. A collisions
between two mutually reactive species do not always lead to chemical reaction. This can
occur for several reasons that will be explained later. At 20 km, about one in ten thousand
collisions result in reaction. Since O2 has a concentration that is more than ten thousand
times greater than any other species that reacts rapidly with O atoms it is immediately
apparent that this must be the major O-atom loss route in the atmosphere.
The depiction of the reaction above appears at first sight to be rather complicated. It is my
own version of representing the chemical reactions for this course, but it contains much
information that should make each process easier to appreciate. The simplified version
found in nearly all texts is given in the orange box. You should use this latter version when
writing out formulas.

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For photolysis and for most chemical reactions occurring in the atmosphere, the rate of loss
of a species is directly proportional to its concentration. This is similar to the rate of leak of
water from a hole in the bottom of the bucket for which the leak rate is proportional to the
height of the water in the bucket. Provided the water does not overflow the top then
eventually the leak rate will equal the input rate via the tap and the water level will remain
constant. Should the input rate decrease, then so will the water level, again until a steadystate level is reached. Here ‘rate’ is analogous to the flow of water from the bucket, but rate
constant is analogous to the size of the hole at the bottom of the bucket. It is the rate
constant for the loss process that determines how quickly a new quasi-steady-state
concentration is established if a rate of input is changed. Many important radicals in the
atmosphere are present in extremely low concentrations. One should not, however, be
misled by this as it might simply imply that the species in question is removed very quickly
by some reactive or photolysis process: it is not the water level in the bucket that is of
importance but the rate of water is flowing through the bucket that is of importance.

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The production of ozone in the stratosphere follows a relatively simple mechanism and is
unavoidable in an O2-rich atmosphere illuminated by wavelengths shorter than 240 nm. As
you have seen from the previous pages, molecular oxygen absorbs UV radiation and
radiation below 240 nm is sufficiently energetic to dissociate it to its atomic parts. By far,
most O atoms produced by this process react with O2. In fact, O atoms are reactive with
some other species too (including with O3), but the massive concentration of O2 compared
to other potential reactants ensures that O atoms collide (and undergo reaction) with O2 at a
frequency orders of magnitude greater than with other minor atmospheric constituents. The
product of this reaction is O3. You will notice that the O + O2 reaction also requires a third

collision partner, M. This represents any atmospheric constituent that can take away the
excess vibrational energy via collision with the initially-formed, vibrationally-excited O3
molecule. If this extra collisions do not occur, then the vibrationally-excited O3 molecule
will simply re-dissociate to O + O2. Re-dissociation is equivalent to no reaction.
Once ozone is formed, it also dissociates due to interaction with UV radiation, and the O
atom formed in this process again reacts with O2 to yield once again ozone (a so-called
"null" process). Finally, a minor reaction between O atoms and O3 can occur that forms
2O2.
This simple mechanism was first described by a British mathematician, Sydney Chapman,
and has since been known as the Chapman mechanism. It was used to predict O3
concentrations as a function of altitude at various places around the world. When, years
later, these predictions were compared with measurements, a consistent and clear
discrepancy was found both in peak concentrations and in the profile of ozone as a function
of altitude. This model however still remains the core of O3 chemistry in the stratosphere.
Due to the strong absorption of O2 and O3 to UV radiation the Chapman mechanism gives
rise to ozone only at altitudes several km above the Earth's surface - see later.

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It is interesting to look at the characteristic times associated with the various processes of
the Chapman mechanism. The combination of the values of absorption cross-section for O2
and the relatively low photon flux below 240 nm means that the photo-dissociation rate of
O2 is rather slow. If there were no mechanism to regenerate O2, then it would decrease over
a time scale of a few thousand years (using 27 km values for photon flux). At the opposite
end of the scale is the average lifetime of O atoms, which is a fraction of a second. This
means that at dusk, the O-atom concentration falls immediately to zero, which, in turn,
means that ozone can no longer be produced. According to the Chapman mechanism, the
loss rate of ozone is determined only by the presence of O-atoms and by UV light. Thus,
during the evening, there should be no removal of O3. In fact, as you will see later, this is

not the case, and O3 does indeed display some variation in concentration between day and
night.


The long average lifetime for O2 means that its global distribution is governed both by the
flux of incoming UV radiation and by the global motions of the air masses. The result is that
ozone has larger concentrations, or more strictly, column densities, around northern and
southern mid-latitudes than above the tropics, where most O3 is generated.

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How can we determine the concentration of O3 as a function of altitude? Let first see if we
can estimate the rate of the various processes.
From process 1, the rate of production of O atoms is 3.6 x 107 cm-3 s-1. This should be
considered a maximum as the presence of O3 at higher altitudes will absorb the light and
reduce the photolysis rate of O2.

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Here is an example calculation of the lifetime of O3 due to the removal by O atoms. Bear in
mind though that average ozone levels remain reasonably constant, so the production rate of
O3 via O2 photolysis should equal the removal rate of ozone by reaction with O atoms
(assuming only the Chapman mechanism is operative). Note also that the removal rate of O2
by photo-dissociation must be equal to the formation rate of O3 by the same process,
however, this does not mean that the characteristic times for O2 and O3 are the same
(compare 5000 years for O2 to 9 years for O3) as their concentrations are different.

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As with molecular oxygen, ozone absorbs very strongly in the UV spectral region.
Importantly for the biosphere, the absorption spectrum of O3 differs from that of O2 in that it
extends far into the near UV (near the visible region, that is), some 100 nm beyond the cutoff point (240 nm) of O2. Whether or not this has any impact on the atmosphere depends on
the concentration of O3.
The Chapman mechanism predicts that significant concentrations of O3 will be formed
only between about 10 and 40 km above the Earth’s surface. A typical profile of ozone is
given in the figure above, but plotted in two different fashions. The lowest profile gives the
concentration of ozone whilst the upper profile gives the volume mixing ratio in ppm (parts
per million). That is, the fraction of total molecules in the air to that of O3 multiplied by one
million. This latter profile peaks higher than the former due to the rapidly decreasing air
pressure with increasing altitude.

If all of the ozone in the atmosphere was compressed to atmospheric pressure into one
thin layer surrounding the Earth, then the thickness of this pure ozone layer (at atmospheric
pressure and 273 K) would be only 3 mm. This concept is, in a sense, used when referring
to ozone concentrations. For many environmental issues, and especially when considering
transmission of light, the total amount of ozone overhead is of most importance rather than
its precise vertical profile and, incidentally, the former is much easier to determine. The unit
of measure for column densities of ozone is the Dobson Unit. One Dobson unit implies an
amount of overhead ozone that is equivalent to a thickness of pure ozone at atmospheric
pressure (at 273 K) of 0.01 mm. Thus, a typical column density for ozone throughout the
world is 300 Dobson Units (DU).
As an example, one can quickly determine the effectiveness on O3 absorption on light at a
wavelength of, say, 260 nm using Beer-Lambert's law. Itr/Io = exp(-cL). Here  is the
absorption cross-section of ozone at 260 nm, which is 1 x 10-17 cm2; c is the concentration

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of ozone. For this we can use pure ozone at atmospheric pressure (and 273 K). This gives
about 2.7 x 1019 molecules cm-3 (using the ideal-gas law). L is the path length. Since we
have (artificially) compressed the gas, the path length is now only 3 mm (300 DU).
Substituting these values into the above equation leads to Itr/Io = 6.6 x 10-36. Thus light at
240 nm is substantially reduced- to effectively zero by the presence of O3.


Most people associate ozone with UV protection, but another extremely important
consequence of the presence of ozone is the formation of the stratosphere. It is due to the
strong absorption by ozone that the air temperature begins to warm above about 20 km.
Cooling occurs again above the wide ozone layer after about 50 km. Thus both at 20 km
and at 50 km a reversal in temperature trend occurs. These reversals effectively partition the
lower atmosphere into three distinct regions, the troposphere, stratosphere, and mesosphere.
Thus the macroscopic structure of the lower-mid atmosphere arises naturally from the
interaction of the UV light from the Sun and molecular oxygen, which leads to the
formation of ozone. Note, that the heating effect on the atmosphere follows closely the
mixing ratio of ozone rather than its concentration due to the rapid change in density of the
surrounding air that has to be heated (i.e., the heat capacity of the air decreases with
increasing altitude, requiring less ozone to induce a given temperature change).

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The effect of the relatively small amount of ozone on the light reaching the Earth’s surface
is dramatic. The solar spectrum reaching the Earth’s atmosphere looks like that shown in the
figure in yellow. The presence of O2 ensures that wavelengths shorter that 250 nm do not
penetrate the atmosphere (orange). But this is not sufficient for living organisms because
significant DNA damage begins at about 305 nm. The addition of ozone to the atmosphere
blocks all wavelengths below about 300 nm (brown). Importantly though, small changes in

O3 concentration (grey) have a relatively large effect on damage to DNA and other biomolecules because of the strong increase in their damage spectrum at wavelengths shorter
than 300 nm. So any process that can decrease the stratospheric ozone concentration has the
potential to seriously effect the Earth’s eco-systems. This was indeed observed due to the
quest to find less toxic refrigerants that resulted in the development of chloro-fluoro-carbons
(CFCs).

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The story of CFCs begins about one hundred and twenty years ago in Belgium with a
chemist named Frederic Swarts who discovered a catalytic technique to replace C-Cl bonds
with C-F bonds in chlorinated hydrocarbons. At that time, Swarts was in a very specialised
field, with only a hand full of researchers in the world working on this subject. His work
began to receive attention in the nineteen twenties when the extremely inventive Thomas
Midgley working as part of a collaboration between Fridigaire (owned by General Motors)
and DuPont to developed safer alternatives to the toxic and corrosive refrigerants (e.g. NH3)
existing at the time. Fluorinated hydrocarbons appeared to be candidates for this.
Substituting a hydrogen with iodine, bromine, or chlorine is not so difficult and such
modified hydrocarbons are often found in nature. However, placing a fluorine atom onto a
hydrocarbon (by removing a H or another halogen) is somewhat more challenging. So
challenging in fact that it occurs extremely rarely in nature, such that nearly all fluorinated
hydrocarbons now found in the world are manmade.
Midgley and his team used a technique similar to Swats‘s and made useful amounts of
dichlorodifluoro methane, known as R-12, CFC-12, or Freon 12, which proved to be ideal:
it was unreactive, non-toxic, non corrosive and had good thermal properties to be used as a
refrigerant.

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