the story of CFCs and stratospheric ozone

Phần 3 của môn học liên quan đến câu chuyện của CFC và quá trình mà nó tạo ra một thảm họa môi trường. Phần này còn nhấn mạnh tầm quang trọng của khí quyển đến hoạt động của con người và sự tuyệt vời của sinh quyển đến những thay đổi trong bầu không khí.

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

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.

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

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

The change in light intensity as it passes through a gas can be easily described by the Lambert 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

Beer-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

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 section (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

cross-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 O To put this into perspective, remember that I /I = cL, where

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 O2and 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 O2only? 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

This graph shows that absorption of radiation by O2accounts 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 anti-correlated, 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

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 Beer-Lambert 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

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

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

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

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

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 steady-state 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

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 O3molecule 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

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 O2and 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

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

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

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 cut-off 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 O3will 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

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)

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 molecules 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)

bio-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

Midgley and his co-workers soon followed R-12 by other useful refrigerants, and everything seemed to be OK for another 40 years

In the 1970's another inventive scientist, James Lovelock, developed a detector that was at least one million times more sensitive than previous detectors for the detection of halogenated hydrocarbons The electron-capture detector was used to detect concentrations

of trichlorotrifluoroethane Lovelock noticed that CFC-113 was always detected at about the same concentration near his home on the west coast of Ireland One of his instruments was then taken along on the research ship R.V Schackleton on a voyage from UK to Antarctic The results of these measurements showed that the concentration of CFC-113 seemed to be extremely constant throughout the world, although there was some variation in latitude A quick estimate of the total amount of CFC-113 in the atmosphere reveal that it was about equal to all the CFC-113 that had been produced to date It seemed not to be removed from the atmosphere

Destruction of hydrocarbons in the atmosphere occurs dominantly by reaction with OH, having daytime tropospheric concentrations in the low 106 radicals cm-3 range The slowest reacting hydrocarbon of all is methane on account of its high C-H bond dissociation energy compared to other hydrocarbons Actually there is a barrier to this reaction (and with most other reactions between OH and alkanes) which determines the reaction rate constant and its temperature dependence In simple terms, the barrier height can be related to a combination

of bond strength and the difference in energy between reactant and products Generally, the C-H bond is weaker than the C-F bond, but also the products of reaction of alkanes with OH have lower energy than their halogenated counterparts The result is a very slow reaction for fluorinated alkanes On the other hand C-Cl bonds are weaker than C-H bonds, but importantly the products are also of higher energy The higher-energy products more than compensates for the weaker bond thus leading to a larger barrier for Cl abstraction reactions than for H abstraction reaction and consequently a slower reaction for the former Thus, tetrachloromethane is about ten times less reactive toward OH than is methane at room temperature Tetrafluoromethane is about three thousand times less reactive at room temperature toward OH than is methane These difference become even greater with decreasing temperature

It turns out that the group of CFCs are one of the longest living groups of species emitted from the Earth’s surface Their lifetimes are generally long enough to allow nearly complete mixing in the lower atmosphere and consequently any effect they have on the atmosphere will be global rather than regional

The Chapman mechanism consistently over-predicted O3 concentrations, but it wasn’t at first obvious why this should be since there was no other species of sufficiently high concentration that could seriously affect it The answer to this puzzle was found by Paul Crutzen, a Dutch Scientist working at that time in Oxford Crutzen proposed that a high concentration of a co-reactant was not necessary if a catalytic process was operative that efficiently regenerated the co-reactant He proposed that NO molecules react with O3 to produce NO2 The so formed NO2 then reacts with O to reform NO Thus a continuous O3destruction route is established

It must be emphasised here that introduction of a catalytic destruction process for O3 does not mean necessarily that the concentration of O3 will be reduced to zero It means that O3steady-state concentrations (arising from an eventual balance of production and destruction routes) will decrease

It was just following the publication by Lovelock on the atmospheric concentration of CFC-113, that James Sherwood Rowland and his postdoctoral researcher, Mario Molina, started to think about the possible effects of CFCs on the Earth’s ozone layer In 1974 they published a paper that showed another possible catalytic cycle for the destruction of ozone (besides the NO/NO2 process), this time involving Cl atoms, and thus, for the first time, cast doubt on the safety of CFCs

For the this catalytic mechanism, the reaction of Cl with O3 results in the formation of ClO radicals, which react with O atoms to reform Cl As you will see in later pages, the effectiveness of this process depends on the presence of adequate concentrations of O atoms Since O atoms are formed

by short-wavelength UV light, they are more abundant in the upper stratosphere rather than in the mid-lower stratosphere, where the highest concentration of ozone are found As a consequence, this cycle is not the most effective Cl cycle for the removal of O3 – though it is in the upper stratosphere The more efficient Cl cycle (given later) at lower altitudes still, though, involves Cl and ClO

An efficient O3 destruction mechanism was therefore found that involved Cl atoms, but could Cl atoms be easily produced from CFCs in the stratosphere at the height of the ozone layer?

It turns out that most CFCs have large absorption coefficients (absorption cross-sections) only in the far UV spectral region where the light flux is essentially zero at ground level But at high enough altitudes (those where ozone is found) a narrow atmospheric window caused by a minimum in the O3 absorption spectrum (and just before O2 begins to absorb strongly) allows a significant amount of light at 200 nm to reach the mid stratosphere, close

to the peak of O3 concentration This results in the photo-dissociation of CFCs and Cl atoms F atoms are not produced in this wavelength region because C-F only ruptures by interaction with shorter wavelengths In any case, as will be shown later, F atoms are not effective in removing ozone

When comparing the effectiveness of O3 and O2 in absorbing light at various wavelengths it

is important to remember that the O2 concentration at 30 km is about 1 x 1017 cm-3, whereas that for O3 is only about 2-3 x 1012 cm-3 Considering this large ratio, and the extent of O2 at higher altitudes, it can be recognised that the decrease in intensity of light at wavelengths shorter than 200 nm is due mainly to the absorption of O2 and not O3, even though O3 has

The large change in the number of photons in the “atmospheric window” around 205 nm is illustrated in this figure Notice the log scale for the number of photons per second per cm2per nm At 20 km altitude the 205-nm peak value is 3 x 107 photons cm-2 s-1 nm-1 For a first approximation lets us suppose that this value extends from 204 nm to 206 nm (a width of 3 nm) and that beyond this the solar flux is too small Now refer to the absorption cross-section values of CCl4 two pages back CCl4 has an absorption cross-section over the wavelength region 204 nm to 206 nm that is relatively constant at 1 x 10-18 cm2 per molecule We assume also that every absorbed photon leads to photo-dissociation, i.e., the quantum yield for Cl production is unity In this case then, the photo-dissociation rate (or

photolysis rate), J, of CCl4 at 20 km is about 3 nm x 3 x 107 photons cm-2 s-1 nm-1 x 1 x 10

-18 cm2 per molecule This is equal to 9 x 10-11 s-1 photons per molecule Since neither photons nor molecule are SI units, they can be omitted; giving simply 9 x 10-11 s-1 In differential form, the rate of removal of CCl4 by photo-dissociation at 20 km is d[CCl4]/dt =

-J Solution of this equation gives [CCl4]t = [CCl4]t=0exp(-Jt) The time it takes for the

initial concentration of CCl4 to reach 1/e (1/2.72) of its initial value is 1/J This is referred to

as the photolysis lifetime At 20 km the photolysis lifetime is 1/(9 x 10-11) s = 352 years If a similar calculation is done at 30 km altitude (refer to the box in the figure) one finds a lifetime of only one month Please be aware therefore that the lifetime does not refer to the

All catalytic destruction cycles that occur in the atmosphere have a limit In the case of the Cl/ClO cycle, reaction of Cl atoms with methane and ClO radicals with NO2 form relatively unreactive molecules that can only very slowly regenerate Cl atoms Methane is the only hydrocarbon with a sufficiently long reactive lifetime to make it to the stratosphere in appreciable concentrations, NO2, on the other hand, is produced in situ, from photochemical

processes origination mainly from the chemically-inert N2O molecule

The dominant process that regenerates Cl atoms following ClO formation actually depends on altitude, as mentioned previously The HO2 radical is closely linked to the OH radical as OH is easily oxidized to HO2 and HO2 is relatively easily reduced to OH In the atmosphere, they are collectively known as HOx As you will see later, other species are closely coupled too and are generally referred to by their collective name (e.g., NOx, NOy, ClOx, SOx, Ox)

HO2can react with ClO It’s rate constant isn’t as large as that between ClO and O but - as seen in the following page – its concentration becomes greater than that of O below about 30 km The reaction leads to the formation of HOCl, which has a broad absorption spectrum extending from the

UV into the visible spectral region The calculations on the following pages demonstrate that this cycle involving HOx is more efficient than that involving O atoms below about 25 km An analogous cycle exists for the catalytic destruction of ozone by Br atoms

Here is an example of how to calculate the relative effectiveness of the ClO/O and the ClO/HOx cycles Here we choose an altitude of 25 km, which has a temperature of 205 K

The rate constants for these two reactions can be expressed in Arrhenius form of k(T) =

Aexp(-E/RT) The parameters A and E/R are obtained, in this case, from a publication by

NASA and the Jet Propulsion Laboratory (JPL) on stratospheric reactions The negative value in the exponent normally means that the reaction rate constant decreases with decreasing temperature This is the situation when the reaction mechanism has a barrier (as

is the case for reactions of OH with alkanes) In the case where the two reactants are radicals, the reactions generally proceed without a barrier but, nevertheless, a weak temperature dependence is sometimes observed, which has its origin in the quantum nature

of the reaction process For the two reactions given here, the rate constants apparently increase slightly with decreasing temperature If the assumption is made that one of the co-reactant concentrations remain constant (in this case HO2 and O, since they are regenerated), then the average rate of removal of ClO is given simply by the product of the

On the previous page an assumption was made that the photo-dissociation of HOCl was not the rate-limiting step (the slowest rate) in the ClO/Cl catalytic cycle involving HO2 If this were the case then significant concentrations of Cl would always remain bound as HOCl meaning that the effective free Cl atom concentration would be small and the catalytic cycle would be less efficient

The enthalpy change, H, for the process HOCl  OH + Cl is 232 kJ mol-1 This corresponds to a wavelength of 516 nm Thus dissociation may occur at wavelengths shorter than this Note that the enthalpy change for the process HOCl  HCl + O is nearly the same as for that above, but in this case, it is unlikely to be a valid photo-dissociation pathway as it requires some re-arrangement of the molecule before dissociation; experimental measurements bear this out

We now go back to the consideration of CFCs By 1974 then a clear mechanism for ozone destruction was understood It might be asked; why it has taken such a long time to implement policies for their phasing out, or at least their reduced production There are three major reasons

Actually if you plot the global production rate of CFCs over the years you will notice that in

1974, the same year as the Sherwood-Molina paper appeared, a sudden drop occurred This was not, though, in response to the scientific findings, but apparently was due to the global oil crisis who's effects were felt on the CFC industry until the early 1980s

This is the BAS base at Halley Bay that houses the Dobson Spectrometer used to record the total column density of ozone In the mid 1980s J Shanklin was responsible for these measurements The building is apparently rebuilt each year to prevent the build-up of ice