Ozone Layer

The ozone layer in our atmosphere plays a vital role in the health of our planet. Ozone is the main absorber of UV-B radiation (280-320 nm) and a co-adsorber with oxygen of UV-C radiation (200-280 nm). Remember that UV-B and UV-C are the most damaging to our skin and bodies. So the ozone layer is very helpful in shielding us from what would be very harmful UV rays.

Where is the "ozone layer"? It is at the very bottom of the stratosphere covering roughly an altitude range of 15-30 km with the maximum concentration in the middle of that range. 90% of all the ozone in our atmosphere is in the ozone layer. The other 10%, unfortunately, is down here in the troposphere where it is a rather nasty air pollutant.

Let's also remember just how little gas is up there at that elevation - about 10 to 100 times LESS than down here on the surface of the planet. So that is 1 bar here on the surface and between 0.01-0.1 bar up there. It makes sense to switch to mbar (10-100 mbar). The air there is still made up of nitrogen, oxygen, and argon but there is just 10 to 100 times less of it. However, there is a LOT more ozone up there - 12 ppm or 12000 ppb vs down here where it is typically (on a good day) less than 70 ppb (which is considered "unhealthy" by the EPA). That equates to around 100 to 1000 times more up in the ozone layer than down here.

It should also be pointed out that there is a reason we give ranges. The layers of our atmosphere are somewhat different in their heights depending on the geographical locations. So a fairly broad range on the ozone layer location. The concentration of ozone also has seasonal fluctuations as well. The units used to measure the ozone layer is the Dobson unit (DU). Most of the earth has around 300 DU (averaged over a year) - but the southern hemisphere and antarctic regions have far less. This is easily seen in figure 1 below.

Figure 1
Figure 2

Figure 2 shows how the concentration of ozone (global average) over the years 1964-1993 steadily decreased. Beginning in 1994 there is a steady recovery - no doubt due to the protocols of removing CFCs from our usage patterns - that ban was in 1989. There are still issues and concerns, but if we stay on our current trajectory, some have estimated that we will get the ozone layer back to its 1980 level by around 2075. And to show this, here is yet another graphic - figure 3.

Figure 3

Let's do a deep dive on all the chemistry here...

The Chapman Cycle

If you hear a climate/atmospheric scientist mention the Chapman cycle, they are talking about the never-ending cycle of oxygen being converted into ozone and vice versa. Let's have a look at making ozone first.

Creation
O2 + UV-C photon → 2 O·
O2 + O· → O3

So a UV photon (≤ 242 nm) hits an oxygen and homolytically breaks the bond to give two identical oxygen atoms which are radicals due to unpaired electrons. Remember that there is a LOT of oxygen gas around (the O2) so it is highly likely that the newly formed oxygen atoms will collide with an oxygen molecule and make a bond to produce ozone. This is how ozone is made - and most all of it is made up in the ozone layer where the UV radiation is the most intense.

Destruction
O3 + UV-B or C photon → O· + O2
O3 + O· → 2 O2
NET RXN: 2 O3 → 3 O2

Ozone itself can take a UV photon (≤ 320 nm) hit and break back into oxygen gas and oxygen atom. Then another ozone molecule can collide with the reactive oxygen atom and produce two oxygen molecules (O2). This is how ozone can be depleted by natural means.

Notice that the two processes above are driven by UV radiation. The extent (rates) of each set of reactions is governed by the concentrations of the reactants. There is a LOT of O2 gas compared to ozone. So the oxygen creates a fairly good flux of ozone. But the ozone is easier to break back into an oxygen gas and an oxygen molecule because the wavelength is considerably longer (less energy required). The result of all this is that ozone IS able to build up to its "normal" levels up in the ozone layer. It's not much (about 12 ppm) but it matters a lot.

A Dynamic Equilibrium

All that creation/destruction stuff above is an example of a system that reaches a dynamic equilbrium state. Dynamic the opposite of static here. Static equilibrium is where there is a "match" of some sort (the equilibrium) and then everything stops. Like a balance beam and two weights, put the weights at the right distance from the fulcrum and you have a match and static equilibrium. Neither mass is changing - it is static.Contrast that with the Chapman cycle. Even though a rather constant concentration of oxygen and ozone persist over time, they are not static. One is constantly changing into the other and vice versa. Like if you were in a sinking boat with a hole in it but you also have a bucket to bail water back into the lake. As long as you match the rate of incoming water with your bucket bailing, you will not sink and the level of water in the boat will be constant. This is another example of dynamic equilibrium.

Now lets THINK about what might happen to a dynamic equilibrium system where one of the processes is interrupted or inhibited. Let's go back to our sinking boat analogy. Uh-oh, you just threw the bucket overboard - you no longer can bail water a bucket at a time. What happens? Well, the boat starts filling with even more water and it begins to sink. So you bail with your cupped hands now like crazy. It's not as good as the bucket, but you do manage to reach yet another stalemate on sinking. It's just that now, the boat is much more full of water which means the water isn't coming in as fast anymore. You've reached a new equilibrium state. The water level in the boat is much higher but the rates of water in and water out once again match. This whole analogy is exactly what happens in chemical systems. The forward and reverse processes tend to find a set of concentrations on each side such that the rates equal and all concentrations remain constant.

So let's kick a hole in the ozone "boat" and see what happens...

CFCs

CFCs are chlorofluorocarbons. As the name clearly states, it's an alkane (hydrocarbon) but with all the H's swapped out for F's and Cl's. Two very popular and famous ones are/were Freon-11 and Freon-12 (aka CFC-11 and CFC-12), not to mention R-11 and R-12 which stood for Refrigerant 11 and 12. Let's not worry about the numbering system for now. They also have a systematic organic name. You can see their name and structure below in figures 4 and 5. So you got...

Figure 4
Figure 5

Great stuff for refrigerant purposes. Definitely solved a lot of unpleasantness with previous refrigerants like ammonia and sulfur dioxide. Those carbon halogen bonds are pretty strong and the stuff just didn't react with anything - basically inert. It wouldn't react and wouldn't catch on fire. Pretty much all refrigeration and AC systems switched over to it and life was good. One could argue that the big expansion and population booms of cities like Dallas and Phoenix, were now far more livable with central AC. Like I said, life was good.

Fast forward a bit. The discovery of the hole in the ozone layer was a big (scientific) concern. The somewhat controversial explanation was that CFCs were to blame - and specifically chlorine. Turns out there was a LOT of chlorine up there in the ozone layer when there should have been none. Guess what else was up there? You got it - CFCs. Their inert nature meant that when they were released down here at ground zero they just followed the winds and eventually ended up in the clouds and higher - finally making it all the way to the ozone layer (the stratosphere). It is estimated that CFCs last up to 110 years in our atmosphere before decomposing somehow (that's pretty inert!). Also, remember that thing about the stratosphere were the UV radiation is so much more intense? Yeah - well that potent UV intensity is perfect for breaking C-Cl bonds into free radicals. Once you have a chlorine atom free radical, bad things happen to ozone... over and over again. Have a look.

1. Cl· + O3 → ClO· O2
2. ClO· + O· → Cl· + O2
  NET RXN: O3 + O· → 2 O2

The CRITICAL thing to see here is that Cl· is not even IN the overall net reaction. Notice that it enters as a reactant in reaction 1, but exits as a product in reaction 2. What this means is that Cl· is a catalyst and not a reactant. It is NOT consumed during the reaction. It is used and regenerated over and over - that is what a catalyst does. Catalysts participate in reactions but are never consumed. The rates of catalytic reactions ARE heavily dependent on the catalyst concentration (the more the faster), but the catalyst is recycled thousands of times. It has been estimated that every chlorine atom in the ozone layer participates in the destruction of over 100,000 ozone molecules. That is why I said "over and over again" in the earlier paragraph. Chlorine is catalytic. Over time, the chlorine atoms do eventually drift back downward and react with gases in the troposphere - but only after thay have done their damage in the stratosphere. Yikes.

Catalysts

  1. Catalysts speed up chemical reactions - they increase the rate
  2. Catalysts are not consumed in a reaction
  3. Catalysts enter the reaction mechanism AND exit the mechanism
  4. Catalysts facilitate an alternate reaction mechanism which is faster
  5. Catalyzed reactions have a lower activation energy which leads to faster reactions

HCFCs to the Rescue?

OK - so CFCs are bad. Bad for ozone anyway and we'd like to KEEP half of our UV protectant of the earth in place. So a new refrigerant type was developed and put into use. HCFCs - hydrochlorofluorocarbons (a mouthful, right?). So now, we think that maybe a little more reactivity down here on the earth's surface wouldn't be so bad. So we stick a H back on freon-12 to get HCFC-22 which is known as R-22 and is really just chlorodifluoromethane (CHClF2). That added H helps facilitate the molecule's decomposition (less inert) - as in it hangs out in the atmosphere for like 12 years instead of 110. The H make the molecule more prone to decomposition here in the troposphere. So the upside of this is that HCFCs have only about 5% of the ozone destroying potential of CFCs. Way better, but that is still not 0%. HCFCs still have one chlorine on them and when they do reach the stratosphere, boom - same thing happens. A Cl· atom is ejected by a UV photon and the whole depletion of ozone is happening again. Although, at a much slower rate than that from the CFCs. Note there is only 1 Cl per molecule (instead of 2 for CFC-12) and many of those molecules break up before reaching the stratosphere. All good right?

Be careful what you wish for... The whole HCFCs were a stop-gap solution. A quick fix to HELP, but not completely solve the problem. We need a NO-chlorine type of refrigerant if we really want to keep Cl· out of the stratosphere. HFCs to the rescue! HFCs are hydrofluorocarbons - no chlorine at all! Our old freon-12 is now going to be HFC-32 where the chlorines are now H's and the formula is CH2F2 and fully named difluoromethane. Great - except that the thermodynamic properties are way different than CFCs and HCFCs. You can't just swap out - you have to build new machinery to handle the different condensation/evaporation cycles. This HAS happened and many modern cars, refrigerators, and home AC units are using the newer eco-friendly HFCs. So we're all good again - right?

Well not quite. There is a major drawback to ALL of these refrigerants. All of them - the CFCs, HCFCs, and the newest HFCs are all really potent greenhouse gases (see previous section). So while we solve one problem we stand to make another problem worse. So back to the drawing board. There are currently newer substances that "work" as refrigerants and have a purposeful functional group such that the lifetime in the atmosphere is greatly reduced. This might be the happy compromise. That, AND maybe trying our best to capture all the refrigerants as best we can and not just release them with reckless abandon into the atmosphere.



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