ENV 102:GLOBAL CHANGE

SPRING 1999





February 1, 1999 --- Dr. Bunce

Atmospheric Chemistry (Continued)


Topics Covered Today:


Chapter 2 Protecting The Ozone Layer (continued)


Chemical Formula For Ozone is O3
Chemical Formula For Oxygen is O2

What is the difference in the bonds of Ozone and Oxygen? Diatomic oxygen has a double bond between two oxygen atoms. Ozone, which is molecule of three oxygen atoms, has a single bond and a double bond among the atoms. As stated in the previous class a single bond involves a sharing of a pair of electrons and double bond is a sharing of two paired electrons between two atoms. In the molecule of ozone there is a single bond and a double bond but one may ask which oxygen atoms are double-bonded and which are single-bonded to each other. There is no right or wrong answer. Here is the definition of this special relationship found in the glossary of your book (See page 490 of Chemistry in Context 2/ Text and Laboratory Manual): resonance – a representation of electron transfer between two (or more) electron distributions within a molecule; the actual electronic distribution is intermediate between the extremes of the resonance forms.

Resonance of Ozone: O=O-O <---> O-O=O

It takes more energy to break a single bond than a double bond. Without ozone in the atmosphere there would be very little protection from ultraviolet (UV) radiation. Some UV photons have enough energy to break apart the oxygen-oxygen double bond in diatomic oxygen or in ozone molecules. It takes a photon with a wavelength of approximately 242 nanometers (nm) to break a double bond between oxygen atoms. Wavelengths greater than 242 nm do not have sufficient energy to break these double bonds. Remember that the energy of a photon is inversely proportional to its wavelength. It would take a photon of less than 242 nm to break the single bond in ozone because it takes more energy to break a single bond than a double bond. Photons in the range of 200-242 nm will break oxygen-oxygen single bonds. Photons in the range of 242-300 nm will oxygen-oxygen double bonds. It is important to recognize that photons of high enough frequencies will do damage to biological molecules by breaking bonds and causing molecules to come apart (i.e. DNA). Oxygen and ozone act as a "screen" that inhibits high frequency, high energy photons from bombarding the surface of the earth.

Diatomic oxygen makes up approximately 21% of the atmosphere and shields the surface of the earth from high energy electromagnetic radiation by absorbing the energy in the double bond. The end results are the breaking of the double bond and separating the molecule into two oxygen atoms. A reaction we would much rather see in the atmosphere than in biological systems.

Reactions of Oxygen and Ozone in the Atmosphere:

O=O + photon (242-320nm) ---> O + O
O-O=O + photon (200-242nm) ---> O=O + O

91% of Ozone is found in the stratosphere, which is a band of the atmosphere that is 10-50 kilometers above the earth. If the stratosphere was somehow brought to the surface of the earth this band would only be 1/8th of an inch thick. This should mentally illustrate to you how spread apart molecules are in the stratosphere, yet the distance between the molecules is sufficient enough to block out most of the sun’s harmful radiation (provided we do not interfere with the natural formation and destruction of ozone in the stratosphere).


How is ozone made and destroyed?

Approximately 300 million tons of ozone is created each day. So why could there possibly any problem? Holes in ozone, how could this be possible? In nature the amount of ozone created each day is equal to the amount of ozone destroyed each day by ultraviolet radiation. When competing reactions equal out this condition is referred to as a "steady state" or "equilibrium". In your book on page 490 the definition is as follows: steady state – a condition in which a dynamic system is in balance so that there is no net change in the concentration of at least some of the participants in the reactions.

The Chapman cycle is the set of four related reactions that represents the natural steady-state formation and destruction of ozone in the atmosphere (page 487).

Chapman Cycle//Stratosphere

Step #1
O=O + high energy photon ---> 2 O (atoms)

Step#2 Ozone Creation
O=O + O ---> O=O-O

Step#3 Ozone Destroyed
O=O-O ---> O=O + O

Step#4 Slow reaction to create the reactant in Step#1
O=O-O + O ---> 2O=O

The combination of oxygen and ozone in the atmosphere helps block various photons of different energies.

Another Representation of the Chapman Cycle Using Previous Equations:

See Figure 2.6 of Chemistry in Context / Text and Laboratory Manual: Applying Chemistry to Society

Students should be able to write the Chapman Cycle out in words, replacing the chemical formulas, to demonstrate their knowledge of the steady state of ozone.




February 3, 1999 --- Dr. Bunce

Atmospheric Chemistry (Continued)


Ozone in the Stratosphere

1. Ozone is not evenly distributed in the stratosphere. It is thinnest at the equator, and thickest at the North and South Poles.

2. Actual measurements of ozone indicate that there is less ozone than there should be as calculated by the Chapman cycle. Therefore, we must look for other ways for ozone to be destroyed. Some of these ways include:

3. Water vapor also affects ozone. Water vapor rarely gets as high as the stratosphere but when it does, the following reaction occurs:

The right hand side of the equation shows free radicals of H and OH. A free radical is an unstable chemical species with an unpaired electron. This free radical is very reactive chemically and can bombard O3 and change it to O2.

4. NO (nitric oxide) can cause O3 to break down. (NO can be formed two ways: the natural way and the unnatural way. It is formed naturally by plants and bacteria growing in the soil, which release N2 O. When this reacts with O2 in the air, it forms NO

In the unnatural process, NO is formed from the gases found in the air when they are subjected to the high heat of engines. Most of these engines are automotive, but airplanes flying at high altitudes, i.e., the Concord, can form it also, thus destroying ozone in the stratosphere.

***One NO molecule can knock out thousands of O3 molecules!

However, even taking into consideration all of these reasons, there was still too much ozone being destroyed. Another source had to be identified.


Ozone and Chlorofluorcarbons

Enter a new chemical culprit --- a class of chemicals known as the chlorofluorocarbons. These are compounds containing chlorine, fluorine and carbon, differing in their chemical makeups by the number of F and Cl atoms associated with each. Examples are:

    CFC # 12 --- Dichlorodifluoromethane (2 CL, 2 F): CCl2F2

    CFI # 11 --- Trichlorofluoromethane (3 Cl, 1F): CCl3F

    Chloroflurocarbons do not occur naturally; they are a artificially produced. CFCs were products developed by duPont in the 1930s for use in the first electric refrigerators. It is an almost inert compound, and replaced the highly toxic ammonia and sulfur dioxide that had been used in earlier refrigerators.

    Following use in refrigeration, CFCs were found to be effective in aerosol dispensing cans and air conditioning systems. The molecules used in these systems leaked readily as a liquid, entering the air around us. The molecules are long-lived, also: a freon molecule can last 120 years!


    Interaction of CFCs with Ozone

    Photon ~~~~~> CCl2 F2 ----> C-ClF2 + Cl- (cleaves off a Cl)

    Next the Cl- free radical pulls an O away from O3:

      Cl- + O3 ---> Cl-O + O2 (this leaves another Cl free radical)

    The Cl-O then reacts with one of the O molecules:

      Cl-O + O ---> Cl- + O2

    If the above two equations are added together, and we cancel out the Cl-s and Cl-Os, what we have is the Chapman cycle:

      O + O3 ---> 2O2

    In this way, Cl can knock out ozone, and then be regenerated, going on and on and on.


    The Antarctica Ozone Hole

    1. South Pole is the coldest place on earth, due to winds that circle the pole, keeping warm air out

    2. During southern hemisphere winter (June - September), the circular polar winds prevent warmer air from entering region.

    3. Water vapor normally found in the atmosphere freezes into thin clouds in the stratosphere, forming ice crystals.

    4. These ice crystals form a "table", or a surface upon which chemical reactions can occur. Reactions on this "surface" can form HOCl and Cl2. When summer comes in October, radiation melts these ice crystals in 1-2 days, and all these chemicals are "dumped" into the stratosphere, forming free radicals of Cl-.

    5. Because of these reactions, a lot of ozone is destroyed by the Cl- in a very short amount of time.

    6. The Antarctica ozone hole is the largest in October but filled again by November because warm winds now can come in, and bring some O3 with them. Also, after the "big dump", very little in the way of chemical reaction takes place in the stratosphere (no more table!) However, the hole is never completely filled in.

    Now, the Arctic is showing some signs of the creation of an ozone hole phenomena.


    (Preparation for Friday’s lab...)



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