Ozone Depletion E4 and E9 Melody Mak & Vivien Tsang Table of Contents
Formation and Depletion of the Ozone…………………………………3 Ozone-­‐depleting Pollutants and their Sources………………………6 Environmental Effects of Ozone Depletion…………………………..7 Oxygen and Ozone Depletion by UV Light…………………………….8 Catalysis of Ozone Destruction by CFCs an NOx……………………9 Ozone Depletion at the Poles………………………………………………10 Sun-­‐screening Compounds………………………………………………….11 2 Formation and Depletion of the Ozone
Ozone (O3) is an allotrope of oxygen (O2). An allotrope occurs when an element can exist with more than one structure in the same physical state. This element occurs in the stratosphere between twelve to fifty kilometers of our Earth’s surface and is in a steady state meaning that its rate of production and rate of destruction are equal. This is because ozone is in a dynamic equilibrium with oxygen in the stratosphere, therefore it is continuously being formed and decomposed. 3 How is ozone formed? First of all, let us compare the structures of oxygen and ozone. Figure 1 Lewis structure of oxygen Figure 2 Lewis structure of ozone Simply by observing the Lewis structures of oxygen and the resonance structures of ozone, we can see there are some differences: • Oxygen atom has a double bond with bond order of 2 • Ozone has a bond order of 1.5 • Ozone has a longer bond length than oxygen • Ozone has delocalised electrons in its structure Now we can move on to discuss how ozone is formed. During the formation of ozone, the strong double bond present in the oxygen molecules is broken by high energy ultraviolet light that sources from the sun. 2O*(g)
O2 (g)
UV (high energy) As seen in the above equation, the high energy UV radiation breaks down the oxygen in to free oxygen radicals. Free radicals possess an unpaired electron, making them very reactive. The oxygen radical formed goes on to react with another oxygen molecule to form ozone as shown below: O* (g) + O2 (g)
O3 (g)
4 Since ozone has a longer bond length than oxygen, we can establish that it has a weaker bond strength. Thus, ultraviolet light of less energy than was needed for oxygen is able to break it. O3 (g)
O* (g) + O2 (g)
UV (lower energy) Another free oxygen radical is produced. This will react with ozone to form oxygen molecules. O3 (g) + O* (g)
2O2 (g)
As these reactions continue to take place, we can see why there is a continuous formation and destruction of ozone that naturally takes place in the Earth’s stratosphere. However, in more recent years, the ozone has been depleting at a faster rate than it is formed due to human activity. The ozone layer serves a significant role in protecting Earth from damaging radiation but our actions could be putting our own health at risk. 5 Ozone-Depleting Pollutants and their Sources
Ozone depletion is a growing problem and concern for us since a decline is the concentration of ozone in the stratosphere could lead to many health problems. Therefore it is important we investigate the pollutants, which induce ozone depletion and their source of origin. (1) Chlorofluorocarbons (CFCs) CFC’s contain a weak C—Cl bond which can easily be broken by ultraviolet light. The result of the breakdown of this bond is the production of highly reactive free radicals. These free radicals can react with 100 000 and more molecules of ozone, hence causing ozone depletion. CCl2F2 (g)
CClF2* (g) + Cl* (g)
UV radiation Cl* (g) + O3 (g)
ClO (g) + O2 (g)
Source: • Propellants in spray cans (aerosols) • Foaming agents for expanding plastic • Refrigerants • Air conditioning • Cleaning solvents (2) Nitrogen oxides (NOx) Nitrogen oxide is able to bring about ozone depletion by reacting with ozone to produce oxygen. NO (g) + O3 (g)
Source: • Combustion of hydrocarbons and fossil fuels 6 NO2 (g) + O2 (g)
Environmental Effects of Ozone Depletion
If depletion of the ozone layer in the Earth’s stratosphere, there may arise severe consequences as more damaging ultraviolet light is able to reach the Earth’s surface. This does not only have a major impact on humans but also plant-­‐life, marine ecosystems and weather patterns. In this section, the effects of ozone depletion can be looked at in closer detail. Effects on humans: • Sunburns become more frequent due to damage from ultraviolet radiation • Increase in melanoma (skin cancer) • Eye cataracts • Blindness Plants: • Inhibition of photosynthesis • Inhibition of growth • Less crop yield (this could pose as a major problem for agricultural farmers) • More susceptible to diseases Marine ecosystems: • Loss of phytoplankton • Loss of biomass • Less food for other marine organisms • Loss of carbon dioxide sink Weather patterns: • Wind and ocean circulation could possibly be affected In order to reduce the rate or amount of ozone depletion caused artificially, it is evident that production of chlorofluorocarbons and nitrogen oxides must be cut down. It is possible to substitute CFCs with a CFC alternative. However, the CFC alternatives must fall under these characteristics: • It must have similar chemical and physical properties to chlorofluorocarbons except it may not have a weak C—Cl bond • Low in reactivity • Low in toxicity • Low inflammability • It should not absorb infrared radiation as this would make it a greenhouse gas 7 Oxygen and Ozone Depletion by UV Light
Oxygen: Ozone: Previously, the structures of oxygen and ozone were compared and it was evident that the bonds in the two molecules were different; oxygen molecules contain a double bond while ozone molecules have a bond that is between that of a single and double bond. Apart from this difference in structure, understanding how UV light of different wavelengths can break these bonds also suggests that they are different in strength. The double bond of an oxygen molecule has an average bond enthalpy of 496 kJmol-­‐1. The wavelength of light needed to break up this double bond is 241nm, which is very short, meaning that it is high in energy. On the other hand, an ozone molecule has an average bond enthalpy of 362 kJmol-­‐1. The wavelength of light needed to break up the bond in this molecule is 330nm. This is longer than that wavelength of UV light needed to break up the double bond in an oxygen molecule, which clearly indicates that the bond in ozone is weaker. 8 Catalysis of Ozone Destruction by CFCs and NOx
In the presence of high energy ultraviolet radiation, homolytic fission of the C—Cl bond in chlorofluorocarbons occurs. This results in the production of chlorine radicals (radical initiation) that go on to deplete ozone molecules and generate further more free radicals (propagation). The process of generating free radicals continues until the radicals escape or terminate. Thus, just one molecule of CFC can catalyse the breakdown of 100 000 ozone molecules, highlighting the damaging impact CFCs have on the environment. This is an example of catalytic ozone destruction. Radical initiation: CCl2F2 (g)
Propagation: Cl* (g) +
UV radiation O3 (g)
CClF2* (g) + Cl* (g)
ClO* (g) + O2 (g)
ClO* (g) + O* (g)
Cl* (g) + O2 (g)
NO2 (g)
NO (g) + O* (g)
Aside from chlorofluorocarbons, nitrogen dioxide also catalytically decomposes ozone by a radical mechanism. Nitrogen dioxide is broken down in the presence of ultraviolet light and oxygen radicals are produced. O* (g) + O3
(g)
UV radiation 2O2 (g)
Nitrogen oxide generated from the breakdown of nitrogen dioxide can also react with ozone to regenerate the catalyst so that the process continues. NO (g) + O3 (g)
NO2 (g) + O2 (g)
9 Ozone Depletion at the Poles
Ozone depletion is greatest in the North and South Poles during the winter and early spring. During the winter season, the low temperature allows more ice crystals to form. These ice crystals contain small amounts of hydrochloric acid and chlorine compounds (e.g. ClONO2). Catalytic reactions can occur on the surface of the ice crystals and as a result, HClO and Cl2 are produced. In the early springtime, the appearance of the sun causes the HClO and Cl2 molecules to break down to produce free chlorine radicals. As seen previously from the effects of CFCs, free chlorine radicals are able to catalyse the destruction of ozone. Therefore, during this time, the ozone layer will have the largest hole. As the weather warms though, the ice will begin to melt and the reaction is no longer possible as there are no more ice crystals. Thus the ozone concentration in the stratosphere gradually increases again. 10 Sun-screening Compounds
We know that the ozone layer is important because it prevents damaging UV radiation from reaching Earth’s surface and causing harm to our skin. UV light is able to break bonds in DNA molecules within the genes of our skin, this could lead to the growth of cancerous cells and result in skin cancer (e.g. melanoma). Our body has a natural sunscreen of its own. This is in the form of the dark brown pigment named melanin. The concentration of melanin increases by the action of sunlight on the surface of our skin, which explains why we get darker when we stay in the sun for too long. Yet this natural sunscreen is not protective enough, especially with a decline in the ozone concentration. Commercially available sunscreens are able to absorb light of particular frequency as they contain delocalized pi electrons. The UV light is absorbed when the pi electrons become excited to higher energy levels without breaking their bonds. An example of such a substance is 4-­‐aminobenzoic acid. Sun blocks can also protect our skin from damaging radiation. They contain white pigments that either reflect or scatter ultraviolet radiation from the sun. Zinc oxide or titanium (IV) oxide is usually contained within these sun blocks. 11