OZONE

OZONE
ozone is an allotropic form of oxygen having three atoms in each molecule - formula = O3.
It is a pale blue, highly poisonous gas with a strong odor and it is an irritating, corrosive, colorless gas with a smell something like burning electrical wiring;Has a detectable odor near electrical machinery;Easily produced by any high-voltage electrical arc (spark plugs, generators, arc welders, as well as photo-copiers, laser printers, CRT-tubes as used in TV and PC-sets, etc)
Ozone is chemically much more active than ordinary oxygen and is a better oxidizing agent
Used in purifying water, sterilizing air, and bleaching certain foods.In the lower troposphere – Ozone occur as a pollutant.Formed from nitrogen oxides and organic gases emitted by automobiles and industrial sources
The definition of Ozone layer
The ozone layer is a region of the atmosphere that is “Earth’s natural sunscreen” since it filters out harmful ultraviolet (UV) rays from sunlight before they can reach us on the surface of our planet and cause damage to humans and other life forms.Any substantial reduction in the amount of overhead ozone, O3, would threaten life as we know it. Thus the appearance of a large “hole” in the Antarctica represent a major environmental crisis.Ozone layer absorbs ultra violet radiation in the range of 220- 320 nm range (UV-B).Exposure of this radiation to human skin leads to sunburn and sun turn).

ANTARCTIC OZONE HOLE
The Antarctic ozone hole was discovered by Dr. Joe C. Farman and his colleagues in the British Antarctic Survey. They have been recording ozone levels over this region since 1957.
Their data indicated the total amounts of ozone each October had been gradually falling, with precipitous declines beginning in the mid – 1970s. This period from September to October corresponds to the spring season at the south pole, and follows a period of very cold polar winter.
By the mid 1980s, the springtime loss in ozone at some altitudes over Antarctica was complete, and amounted overall to a loss of more than 50% of the total overhead amount. It is therefore appropriate to speak of a “hole” in the ozone layer which now appears each spring over the Antarctic and which lasts for several months.
In 1993, the ozone concentration dropped to a record low of 90 DU in early October. It was not clear for several years after it discovery whether the hole was due to a “natural” phenomenon involving meteorological forces, or was due to a chemical mechanism involving air pollutants.
In the later possibility , the suspect chemical was chlorine, produced mainly from gases that are released into the air in large quantities as a result of their use, for example, in aerosol spray cans and in air conditioners.
To discover why the hole formed each spring, an American emergency ozone research expedition headed by Dr Susan Solomon of the National Oceanic and Atmospheric Administration in Boulder, Colorado went to the Antarctic in late winter in August 1986.
Using moonlight as a light source, Solomon and her co-workers were able to identify from the specific wavelength of light absorbed by atmospheric gases that certain, molecules were present in the atmosphere far above their heads.
As a results of that research and subsequent investigations, it is known that the hole indeed does occur as a result of chlorine pollution. Furthermore, it has been predicted that the hole will continue to reappear each spring for the next several decades, and that a corresponding hole may one day appear above the Arctic region.
The worldwide loss of ozone has become a major environmental concern, since it results in less protection to life at the Earth’s surface from the harmful ultraviolet component of sunlight
In this lecture, we shall investigate the chemical processes involved in the production of the ozone layer, as well as those underlying the phenomenon of ozone depletion.
The Chemistry of Formation and Depletion of Ozone layer
The chemistry of ozone depletion and other process in the stratosphere, is driven by energy associated with light from the sun.
Therefore, Investigation was made based on the relationship between light absorption by molecules and the resulting activation, or energizing, of molecules that enables them to react chemically.
Note: The environmentally relevant portion of the electromagnetic spectrum
Substance differ enormously in their propensity to absorb light of a given wavelength because of difference in the energy levels of their electrons . For example an object we perceive as black in colour absorbs light at all wave lengths of visible spectrum - that is from about 400nm(violet light) to about 750nm (red light).
Notice that the UV region begins at the violet edge of the visible region, hence the same ultraviolet. The division of the UV region into components will be discussed later. Beyond the ultraviolet range, i.e., of even shorter wavelength, are X- rays. At the other end of the spectrum, beyond the visible region, lies infrared light, which will become important to us when we discuss the green house effect in the following lectures.E.g. diatomic molecular oxygen, O2, does not absorb visible light significantly, but does absorb some types of ultraviolet (UV) light which is electromagnetic radiation with wavelengths between about 50 to 400 nm.
The O2 gas that lies above the stratosphere is responsible for filtering from sunlight most of the UV light from 120 to 220 nm, the rest is filtered by the O2 in the stratosphere. Ultraviolet having wavelengths shorter than 120 nm is filtered in and above the stratosphere by O2 and other constituents of air such as N2. Thus no UV light having wave lengths shorter than 220 nm reaches the Earth’s surface , thereby protecting our skins and eyes from damage by this part of sun’s output. O2 also filters some UV light in the 220-240 nm range, called UV-C subrange.
Ultraviolet light in the 220 - 320 nm (UV-B) range is filtered from sunlight mainly by ozone molecules, O3, that are spread between the middle and lower regions of the stratosphere. Since the molecular constitution and thus its set of energy levels of ozone is different from that of oxygen, its light absorption characteristics also are quite different.Ozone aided to some extent by O2 in the shorter wavelength range, filters out all of the sun’s ultraviolet light in the 220 – 290 nm range, which overlaps the 200 – 280 nm region known as UV-C.However, O3 can only absorb a fraction of the sun’s UV light in the 290 – 320 nm range, its ability to absorb light of such wavelengths is quite limited. The remaining amount, 10-30% (depending upon latitude), penetrates to the earth surface..
Thus O3 is not completely effective in shielding us from light in the UV-B region, defined as that which lies from 280 – 320 nm.Because neither O3 nor any other constituent of the clean air absorbs significantly in the UV-A range (320-400nm), most of this, the least biologically harmful type of ultraviolet, does penetrate to the Earth surface.
The increase in UV-B is the principal environmental concern about ozone depletion, since it leads to detrimental consequences to some life forms, including human beings.
Exposure to UV-B causes human skin to sun burn and sun tan; overexposure can lead to skin cancer.
Increasing amounts of UV-B may also adversely affect the human immune system and the growth of some plants and animals.Most biological effects arise because UV-B can be absorbed by DNA molecule, which then can undergo damaging reactions.Almost all cancers are due to overexposure to UV-B in sunlight, and so any decrease in ozone is expected to yield eventually an increase in the incidence of skin cancer.Fortunately, most skin cancer is not the often-fatal malignant melanoma but rather a slowly spreading type that can be treated.The incidence of the malignant form of skin cancer is related to short periods of very high UV exposure, particularly early in life, and particularly for fair skinned, fair- haired, freckled people who burn easily.There is some evidence also that increased UV-B levels give to more eye cataracts, particularly among the non-elderly. A 10% increase in UV-B is predicted to result in 6% more cataracts among 50- year- olds.Increased UV-B exposure also leads to a suppression of the human immune system with resulting increase in the incidence of infectious diseases.It is speculated that increasing amounts of UV-B interfere with the efficiency of photosynthesis, and thus that plants will respond by producing less leaf, seed, and fruit.
All organisms that live in the first five meters or so below the surface in bodies of water will also experience increased UV-B exposure and may be at risk.
In Particular, it is feared that production of the microscopic plants called phytoplankton near the surface of sea water may be at significant risk from increased UV-B; this would affect the marine food chain for which it forms the base.The recent worldwide drop in the population of frogs and other amphibians has now been linked to increasing levels of UV.
Principles of photochemistry
As Albert Einstein was the first to realize, light can be considered not only a wave phenomena but also to have particle-like properties in that it is absorbed (or emitted) by matter in finite packets now called photons.The energy E of each photon is related to the frequency ν and the wave length λ of the light by the formulas.
E = h ν or E = hc/ λ since λν = C
Where h is Planck’s constant (6.626218 × 10-34J s) and C is the speed of light ( 2.997925 ×108 ms-1).
From the equation, it follows that the shorter the wavelength of the light, the greater the energy it transfers to matter when absorbed

UV light is high in energy content, visible light is of intermediate energy, and Infrared is lower in energy. Furthermore, UV – C is higher than UV – B which in turn is more energetic than is UV – A.

For convenience, the product hc in the equation E = hc/ λ can be evaluated on a molar basis to yield a simple formula relating the energy absorbed by 1 mole of matter when each molecule in it absorbs one photon of a particular wavelength of light.
Expressing wavelength in nanometers, the value of hc is 119,627 kJ mol-1, E = 119,627/ λ. Where E is in kJ mol-1.

The photon energies associated with some wavelengths of critical importance to stratospheric chemistry have been calculated from this formula. These photon energies for UV and visible light are of the same order of magnitude as the enthalpy (heat) changes, ∆H0, of chemical reactions, including which dissociate atoms from molecules.
For example, dissociation of molecular oxygen into its monatomic form requires an enthalpy change of 495 kJ-1
O2 → 2 O ∆H0 = 495 kJ mol-1
Recall that ∆H0 stands for the enthalpy change determined under standard conditions; to a good approximation, for a dissociation reaction ∆H0 is equal to the energy required to drive the reaction under stratospheric pressure and temperature conditions.
Since the energy is all to be supplied by one photon per molecule, the corresponding wavelength for the light is
λ = 119,627 kJ mol-1 nm / 495 kJ mol-1= 241 nm.
Thus any O2 molecule that absorbs a photon from light of wavelength 241nm or shorter has sufficient excess energy to dissociate;The photon energies associated with some wavelengths of critical importance to stratospheric chemistry have been calculated from this formula and listed in the Table below.
If energy in the form of light initiates a reaction, it is called a photo-chemical reaction.
O2 + UV photon → 2 O
The oxygen molecule in the above reaction is variously said to be photochemically dissociated or photochemically decomposed or to have undergone photolysis.
Molecules that absorb light generally do not retain the excess energy provided by the photon for very long. Within a tiny fraction of second, they must either use the energy to react photochemically or else it is dissipated usually as heat energy that becomes shared among several neighbouring molecules as a result of collisions ( i.e., molecules must “ use it or lose it”).Thus molecules normally cannot accumulate energy from several photons until they receive sufficient energy to react; all the excess energy required to drive a reaction usually must come from a single photon.
Therefore light of 241 nm or less in wavelength can result in the dissociation of O2 molecules, but light of higher wavelength does not contain enough energy to promote the reaction at all.
The energy from a photon of wavelength greater than 241 nm absorbed by an O2 molecule is rapidly converted to an increase in the energy of motion of it and of the molecules that surround it.


THE CREATION AND NONCATALYTIC DESTRUCTION OF OZONE
In this section, the formation of ozone in the stratosphere and its destruction by noncatalytic processes are analysed. The formation reaction of ozone generates sufficient heat to determine the temperature in this region of the atmosphere.
Above the stratosphere, the air is very thin and the concentrations of molecules is so low that most oxygen exists in atomic form, having been dissociated from O2 molecules by UV-C photons from sunlight.The eventual collision of oxygen atoms with each other leads to the reformation of O2 molecules, which subsequently dissociate again when more sunlight is absorbed.
In the stratosphere itself, the intensity of the UV-C light is much less since much of it id filtered by the oxygen that lies above, and since the air is denser, the molecular oxygen concentration is much higher. For this combination of reasons, most stratospheric oxygen exist as O2 rather than as atomic oxygen,Because the concentrations of O2 molecules is relatively large and the concentration of atomic oxygen is small, the most likely fate of the stratospheric oxygen atoms created by the photochemical decomposition of O2 is their subsequent collision with undissociated oxygen molecules, thereby resulting in the production of ozone:
O + O2 → O3 + heat.
atomic oxygen diatomic oxygen Ozone.Indeed, this is the source of all the ozone in the stratosphere. During the daylight, ozone is constantly being formed by this process, the rate of which depends upon the amount of UV light and the concentration of oxygen molecules at a given altitude.
At the bottom of the stratosphere, the abundance of O2 is much greater than at top because air density increases progressively as one approaches the surface. However, relatively little of the oxygen at this level is dissociated and thus little ozone is formed because almost all the high energy UV has been filtered from sunlight before it descend to this altitude. For this reason the ozone layer does not extend much below the stratosphere.
In contrast, at the top of the stratosphere, the UV-C intensity is greater but the air is thinner and therefore relatively little ozone, O3, is produced since the oxygen atoms collide and react with each other rather than with small number of intact O2 molecules.
Consequently, the density of ozone reaches a maximum where the product of UV-C intensity and O2 concentration is maximum.
This maximum density of ozone occurs at about 25 km over tropical areas. Most of the ozone is located between 15 and 35 km, i.e., the lower and middle stratosphere, known as the ozone layer.
The release of heat by the reaction results in the temperature of the stratosphere being greater than the air below or above it. Indeed, the stratosphere is defines as the region of the atmosphere that lies between these temperature boundaries.
Notice that, in the stratosphere the air at a given altitude is cooler than that which lies above it. The general name for this phenomenon is a temperature inversion.
Because cool air is denser than hot air, it does not rise spontaneously due to the force of gravity; consequently, vertical mixing of air in the stratosphere is very slow process compared to that in the troposphere
The air in this region therefore is stratified – hence the name stratosphere.
The absorption of a UV-C and UV-B photon by an ozone molecule in the stratosphere results in the destruction of that molecule:
O3 UV Photon (λ < 320nm) → O2 + O*
The oxygen atom produced in the reaction of ozone with UV light have an electron configuration that differs from that which is that of lowest energy. Atoms or molecules that exist temporarily in such situations are said to be in an excited state and their symbols are marked with a superscript asterisk
They posses additional positive energy compared to the lowest energy arrangement of electrons, which is referred to as the ground state. Unless an excited atom or molecule quickly reacts with some other atom or molecules, its excess energy is lost.
Most oxygen atoms produced in the stratosphere by photochemical decomposition of ozone or of O2 subsequently react with intact O2 molecules to reform ozone.
Some oxygen atoms react with intact ozone molecules to destroy them by conversion to O2:
O3 + O → 2 O2
This reaction is inherently slow, since its activation energy of 18kJmol-1 is a sizable one, the result being that few collisions occur with sufficient energy to result in reaction.
To summarize, ozone in the stratosphere is constantly being formed, decomposed, and reformed during daylight hours by a series of reactions that proceed simultaneously.
The ozone gas filters UV-B and UV-C from sunlight, but is destroyed temporarily in this process or by reaction with oxygen atoms.
Ozone is not formed below the stratosphere due to a lack of the UV-C required to produce the O atoms necessary to form O3, because this type of sunlight has been absorbed by O2 and O3 in the stratosphere.
Above the stratosphere, oxygen atoms predominate and usually collide with other O atoms to reform O2 molecules.
CATALYTIC PROCESSESS OF OZONE DESTRUCTION.
In the early 1960 it was realized that there are mechanisms for ozone destruction in the stratosphere in addition to the normal processes described earlier.
In particular, there exist a number of atomic and molecular species, designated in general as X, that react efficiently with ozone by abstracting (removing) an oxygen atom from it:
X + O3 → XO + O2
In those region of the atmosphere where the atomic concentration is appreciable, the XO molecules react subsequently with oxygen atoms to produce O2 and to reform X:
XO + O → X + O2
The net sum of the two reactions above is
O3 + O → 2 O2
Thus the X are catalysts for ozone destruction in the stratosphere, since they speed up a reaction (here, between O3 and O) but are eventually reformed intact and are able to begin the cycle again – with, in this case, the destruction of further ozone molecules.
The X catalysts greatly increase the efficiency of this reaction. All the environmental concerns about ozone depletion arise from the fact that we are inadvertently increasing the stratospheric concentrations of several X catalysts by release at ground levels of certain gases, especially those containing chlorine.Such an increase in the catalyst concentrations will lead to a reduction in the concentration of ozone in the stratosphere.
A factor that minimizes the catalyzed gas- phase destruction of ozone is the requirement for atomic oxygen to complete the cycle by reacting with XO in order to permit the regeneration of the X catalyst is a usable form.

Natural variations of the ozone layer
The thickness of the ozone layer varies with season and location on Earth. Under normal conditions the thickness of the ozone layer is larger at high latitudes than at low latitudes.For example, larger ozone concentrations are usually found above Northern Europe, Canada and Siberia than above the equator.Another important fact is the annual cycle. On average, the thickness of the ozone layer is greater in March and April and lowest in October November. This cycle is related to the general air circulation in the stratosphere at different seasons.
Lewis Structures of Free Radicals
Most of the free radicals that are important in atmospheric chemistry have their unpaired electron located on a carbon, oxygen, hydrogen or halogen atom.The specific atomic location can be denoted by plain a “dot” above the relevant atom symbol to represent the unpaired electron e.g. F., Cl. (their unpaired electrons are not in actual use as bonding electrons).Thus a carbon atom on which unpaired electron is located forms three rather than four bonds H-O-C=O , Oxygen forms one rather than two bonds, Halogen or hydrogen forms one bond.The unpaired electron exists as a nonbonding electron localized on one atom, not as a single bonding electron shared between atoms. For many polyatomic free radicals, the choice of atom to which the unpaired electrons is to be assigned in deducing the Lewis structure is from atom-atom connections.
In hydroperoxy radical HOO, the hydrogen atom can not be the radical site since it must form a bond to the adjacent oxygen; and neither can the central oxygen since it must form two bonds one to each neighbor. The terminal oxygen is the appropriate radical site




REFERENCES
David Urbinato(Summer 1994).Environmental Protection Agency.London,United States
Spengler,John D;Sexton,K.A.(1983)."Indoor Air Pollution.A Health Perpectives