Oxygen Molecules

How Else Does Ozone Purify?

Ozone is also biocidal - which means it kills harmful biological and bacterial contaminants. This biocidal action results from its reaction with the double bonds of fatty acids in bacterial cell walls, membranes and the protein capsid of viruses. In bacteria, the oxidation results in a change in cell permeability and leakage of cell contents into solution. Ozone attacks these cell walls, breaking down membranes and ultrastructural components of the organism. In more simple terms, the unstable electrons of ozone blast holes through the membranes. This occurs by cell lysing or rupturing the cell wall of viruses, bacteria, yeast, and abnormal tissue cells, thereby destroying them by inactivation of the microorganism's enzymes. In viruses, alteration of the protein capsid prevents the virus from being taken up by susceptible cells.

Ozone displays an "all or nothing" effort in terms of destroying bacteria. It is such a strong germicide that only a few micrograms per liter are required to demonstrate germicidal action. Factors like humidity, temperature, pH, ozone concentration levels, type of organism and time, determine the kill rate for pathogens. The action of ozone gas in water is instantaneous. After oxidation, ozone returns to its original form of oxygen, with out leaving any toxic by-products or residues.

Atmospheric ozone is found in varying concentrations from the Earth’s surface to a height of about 60 km. The maximum amount of ozone occurs between 20 and 30 km in the region of the stratosphere characterized by increasing temperatures-the ozone layer. The ozone layer is like "Earth’s natural sunscreen" because it filters out harmful ultraviolet (UV) rays from sunlight before they reach the planet’s surface. Ozone is seen in the troposphere, however, where it is actually a harmful component of smog. Ozone is an air pollutant is found to cause respiratory problems to those exposed to it during smog events.

Above the stratosphere most oxygen exists in atomic form having been dissociated from O₂ molecules by UV-c photons from sunlight. Chemists represent this form of oxygen with an asterisk (O* ) to emphasize the fact that the atom is in a particularly reactive, excited state when it really wants to react. In the stratosphere itself the intensity of UV-C light is less and the air is denser therefore the molecular oxygen concentration is much higher.

Figure1: The types of photons being released by the sun are related to the amount of ozone found in the stratosphere. The ozone is most concentrated in the region of the stratosphere penetrated by UV-b and UV-a light. This is the type of UV-b light which is particularly harmful to humans is the type of UV light that the ozone layer protects us from. ( http://see.gsfc.nasa.gov//edu/SEES/strat/class/S_class.htm)

When an oxygen molecule absorbs a UV-C photon the process of photodissociation takes place:

Equation 1:

O₂ + UVphoton ( <240nm) 2 O^*

^* Represents electronically "excited" state of the atom with excess energy

The most likely fate of oxygen atoms in the stratosphere created by the photochemical decomposition of O₂ is subsequent collision with intact diatomic molecules of oxygen (O₃) resulting in ozone (O₃) production:

O^* + O₂ O₃^*

The excess energy represented by the asterisk must be transferred away from the O₃ molecule or the reverse decomposition will occur. The excess energy can be released by colliding with another atom or molecule represented by M (usually N₂ or O₂) and releasing some of the excess energy:

Equation 2:

O^* + O₂  O₃*

O₃* + M O₃ + M*

_____________________

O + O₂ + M O₃ + M*

Figure 2: This is a summarizes the process described in equation 1 and equation 2 above. The photochemical decomposition of O₂ is diagrammed on the left side of the figure and the formation of O₃ is shown on the equations on the right side of the figure. (http://see.gsfc.nasa.gov//edu/SEES/strat/class/S_class.htm)

At this point the ozone starts to do its job by absorbing UV light with wavelengths shorter than 320 nm providing a shield against high-energy radiation penetrating the Earth’s surface. The photodecomposition of ozone reverses the reaction leading to its formation:

Equation 3:

O₃ + hv ( <320nm) O₂ + O^* (photochemical decomposition)

O^* + O₃ 2O₂

O^*+ O^* + M O₂ + M* (spontaneous reaction)

 Processes of Ozone Depletion

Fluorinated chlorocarbons were developed in 1930 by the General Motors Research Laboratories in a search for a non-toxic non-flammable refrigerant to replace sulphur dioxide and ammonia then in use. CFC-12, pure CF₂Cl₂, is produced by reacting carbon tetrachloride with gaseous hydrogen flouride. It was widely used as a liquid coolant and was later discovered to be useful as good thermal insulator.

 CFC-11, CFCl₃, was also used for foam products used to insulate refrigerators, freezers, and buildings. Both CFC-11 and CFC-12 were used extensively in aerosol spray cans.

The other CFC of environmental concern is CF₂Cl-CFCl₂, CFC-113. It was used as a cleaning compound for electronic circuit boards in the 1970s. The uses of CFCs all ultimately lead to atmospheric release, since even ‘hermetically sealed’ refrigerators and closed-cell foams finally leak to the air. About 90 percent of all the CFC-11 and CFC-12 produced is thought to have been released.

CFCs are virtually inert near the earth’s surface because they lack hydrogen-carbon bonds that tend to break apart and release toxic chlorine radicals into the lower atmosphere. CFCs have no tropospheric sink (they are not consumed in the troposphere like many released chemicals) and therefore all molecules eventually rise to the stratosphere.

The lack of reactivity that makes CFCs so useful commercially is also what allows them to survive in the atmosphere and diffuse into the stratosphere where they are exposed to high-energy radiation causing photodissociation. Because vertical motion in the stratosphere is slow, the atmospheric lifetime of CFC-11 molecules is around 60 years and that of CFC-12 is around 105 years. The C-Cl bonds in CFCs is weak enough that free chlorine radicals are readily formed in the presence of light with wavelengths in range 190 to 225 nm. For example the reactions of C-12 are:

Equation 4:

CF₂Cl₂ + hv CF₂Cl + Cl

Cl + O₃ ClO + O₂

ClO + O Cl + O₂

_________________

O + O₃ O₂ + O₂

Another important class of ozone-depleting chemicals are halons: bromine-containing, hydrogen free substances such as CF₃Br and CF₂BrCl. Halons were developed by the U.S. Army Corps of Engineers after WWII for extinguishing fires in tanks and aircraft. Total global production of the three halons (Halon-1211,-1300, and –2400) was about 17M lbs in 1985. Like CFCs, halons have no tropospheric sink. They are of special concern because their ozone depleting potential per molecule is estimated to ten times that of CFCs.

Like chlorine, bromine radicals released by photochemical decomposition can destroy ozone:

Equation 5:

Br + O₃ BrO + O₂

BrO + O Br + O₂

_________________

O₃ + O 2O₂

Figure 4: This figure diagrams in the left hand corner the formation of an O₃ molecule which can then be destroyed by the free chlorine radical released by CFCs in the stratosphere like in equation 3. Bromine containing compounds (BrO_x)destroy ozone in a similar reaction as in equation 4.

The Antarctic Ozone Hole

In 1985, Josesph Farman, Brian Gardiner, and Jonathan Shanklin in the British Antartic Survey discovered that ozone was disappearing over Antarctica during this southern hemisphere spring period in amounts much less than even the naturally occurring low amounts over Antarctica during this season.

The ozone hole occurs because of special polar winter weather conditions in the lower stratosphere that temporarily convert all the Cl that is stored in the inactive forms HCl and ClONO₂ into the active Cl and ClO . As the stratosphere cools to very cold temperatures over the Antarctic during the southern winter, polar stratospheric clouds (PSC’s) form. When temperatures drop a few degrees below –80^oC a special type of crystal forms. These inactive forms of Cl can react on the thin aqueous layers on the surfaces of these PSC crystals and produce Cl products that can catalytically destroy ozone.

Overall, there is an ozone destruction rate of about 2% per day occurring each September due to the combined effects of the various catalytic reaction sequences. By early October almost all the ozone is wiped out between altitudes of 15 and 20 km.

The key ingredients for polar ozone losses are: high chlorine and bromine levels, cold temperatures during the late winter, and relative isolation of the polar region from the midlatitudes. Increasing levels of CFCs and bromine compounds over the past few decades has caused the Antarctic ozone hole since temperatures are always cold in the late winter.

Figure 5: The amount of ozone over Antarctica in October when the concentration is the lowest is steadily declining. The amount of ozone is represented in Dobson units -one DU is equivalent to a 0.01 mm thickness of pure ozone if it were at ground level temperature and pressure.

 Biological Effects of Ozone Depletion

Most biological effects of sunlight arise because UV-B can be absorbed by DNA molecules. Most skin cancers in humans are due to overexposure to UV-B in sunlight. The incidence of malignant skin cancer affecting one in 100 Americans is thought to be associated with high UV exposure especially at a young age.

It is predicted that there will be a 1-2% increase in malignant skin cancer for each 1% decrease in ozone. People living at latitudes around 45^o N are at an increased risk. It is predicted that there will be an eventual 22% increase in malignant cancer due to a 6.6% decrease in ozone between 1979 and 1992 over those regions.

It is also speculated that ozone depletion increasing UV-B exposure will interfere with the photosynthesis process and plants will produce less leaf, seed, and fruit. It is also predicted that production of phytoplankton near the surface of sea water will be at risk thus affecting the marine food chain.

 Policy and Ozone Depletion

The usage of aerosol spray cans containing CFCs was essentially eliminated in the 1970s by the US, Norway, Canada, and Sweden although not in other parts of the world. A national agreement wasn’t formerly made until 1987 at the conference in Montreal, Canada. The Montreal Protocol was initiated at this conference and amended in 1990, 1992, 1995, and again in 1997. It was finally decided that "all ozone-depleting chemicals" are destined for phaseout all over the world. All legal production of CFCs was ended in 1995 in developed nations. Developing nations have been given until 2010 to curb production.

Developed countries have also agreed to end production of HCFCs by 2030 and developing countries by 2040. Halon production was banned in 1994 but the use of stock piles continues. Methyl bromide will be phased out by developed countries by 2005 and developing countries by 2015.

Figure 6: This is a look at CFC use in 1991 as compared to 1974 (Figure 3). The overall amount of CFC use is much smaller in 1991 and the percentage released by propellants is much lower due, in part, to aerosol bans in the 1970s. (http://see.gsfc.nasa.gov//edu/SEES/strat/class/S_class.htm)

The EPA also enacted a market-based system to phaseout ozone depleting substance production. In 1988 the EPA assigned annual permits called "allowances" to five CFC producers, three halon producers, fourteen CFC importers, and six halon importers. Allowances permitted a company to make or import a certain quantity of CFCs and halons for the U.S. market and also let companies produce an additional 10% for export to developing nations.

Companies could trade their allowances within their company and to other firms as well as outside the U.S. In 1990 the system was amended so that each ozone depleting substance(ODS) was given a separate allowance taking into account its ozone depleting potential instead of all ODS being treated equally. The amended system also included a 1% tax to "pay" the environment on every allowance trade.

 The Future of the Ozone Layer?