1-9: Ozone Reactions: Physics and Chemistry of the Stratosphere

Unit Objectives

• To review atmospheric measurements of manmade chemicals and concurrent measurements of stratospheric ozone.
• To understand the atmospheric conditions, chemicals, and complex reactions leading to ozone depletion in certain regions of the atmosphere.

Preparation for Online and Class Discussion

1. Read the summary information.

   Class images

2. Print for offline review.

3. Take the summary information quiz. Quizzes are available from your portfolio.

4. Additional Related Reading. Although you are not ultimately responsible for content listed in this section, we encourage you to fully explore the links in order to provide a better scope on the upcoming class discussion.

   Houghton, J. T., G. J. Jenkins, and J. J. Ephraums, 1990: Climate Change, The IPCC Scientific Assessment. Cambridge University Press. p. 27-28.

Online and Class Discussion

• Social Comments
• Knowledge-building Comments
• Ethical Comments

• Unit Problem to Ponder
• Unit exam questions from previous years
• Unit discussions from previous years

Unit Wrap-up Materials

• Group Discussion Summary:

  Summary prepared by: Francis Otieno and Heather Bickert

  Ozone is an atmospheric compound, that is concentrated in the stratosphere.The name ozone comes from Greek ozein, which means "to smell". In its standard state, ozone is pale blue, highly poisonous gas with a strong odor. It is much more active chemically than ordinary oxygen and is a better oxidizing agent. It is used in purifying water, sterilizing air, and bleaching certain foods. In the lower atmosphere atmosphere, ozone is formed from nitrogen oxides and organic gases emitted by automobiles and industrial sources. In the stratosphere, ozone is formed when a molecule of oxygen absorbs a photon of light with wavelengths shorter than 200nm, is split into atomic oxygen and the atomic oxygen reacts withan oxygen molecule to form ozone.

  O2 + UV light -> 2 O
  O + O2 + M -> O3 + M (where M indicates conservation of energy and momentum)

  These reactions absorb up to 98% of the UV-light therby shielding the earth¢¥s surface from biologically damaging ultraviolet radiation. The exchange between O3 and O2 is estimated at 300 million tons a day. Ozone molecules exposed to ultraviolet may break down into O2 and O releasing heat and warming the upper atmosphere.

  O3 + UV, visible light -> O + O2

  The free oxygen atom may then combine with an oxygen molecule, creating another ozone molecule, or it may take an oxygen atom from an existing ozone molecule to create two ordinary oxygen molecules.

  O + O2 -> O3 or O3 + O -> O2 + O2

  These light catalyzed reactions, in which ozone is produced and destroyed, are known as "Chapman Reactions." Ozone, levels change periodically as part of regular natural cycles. Sunspots, emitting high levels of electromagnetic radiation with a period of about 11-years, contribute to ozone production and account for 2 to 4 % of the total variation in ozone concentrations. The quasi-biennial oscillation in which tropical winds switch from easterly to westerly every 26 months also accounts for approximately 3 % of the natural variation in ozone concentration. Concentrations are also affected by isolated events such as volcanic eruptions.

  Because, ozone is a highly unstable molecule it readily donates its extra oxygen molecule to free radical species such as nitrogen, hydrogen, bromine, and chlorine.

  O3 + X -> XO + O2 ( where X may be O, NO, OH, Br or Cl)

  Chlorine, released from CFCs, and bromine (Br), released from halons, are two of the most important chemicals associated with ozone depletion. Halons are mainly used in fire extinguishers while CFCs which may remain in the atmosphere for at least 40 years are used largely in aerosols, air conditioners, refrigerators, and cleaning solvents. In the stratosphere, high-energy ultraviolet radiation causes the CFC molecules to break down through photodissociation releasing atomic chlorine. The free chlorine atom initially reacts with an unstable oxygen containing compound, such as ozone, to form chlorine monoxide (ClO). The chlorine monoxide then reacts with atomic oxygen to produce molecular oxygen and atomic chlorine. The regenerated chlorine atom is then free to initiate a new cycle.

  CCL2F2 + photon->CCLF2+CL
  CL+O3->ClO +O2
  ClO+O->CL+ O2
  CL +O3 +ClO +O ->ClO+Cl+O2+O2

  The other reactions involving Nitrous oxide and water vapor are summarized below.

  H + O3 -> HO +O2
  HO + O -> H + )2
  H+O3+HO+O->HO +H+2O2 H AND OH SERVE AS CATALYSTS.
  N + O3 ->NO +O2
  NO + O-> N +O2
  N+O3+NO+O->NO+N+2O2 NITROUS OXIDE IS A CATALYST FOR OZONE DESTRUCTION.

  In each case the general reaction is O+O3->2O2 and the H, OH, nitrous oxide, atomic chlorine only act as catalysts.

  Problem to ponder

  These reactions are responsible for the ozone hole. The topography of Antarctica creates a stagnant whirlpool of extremely cold stratospheric air over the region during the long polar night. The whirlpool prevents warm air from reaching the south pole while infrared radiation freezes the air within this polar vortex all winter, leading to formation of thin polar stratospheric ice clouds. These ice crystals provide surfaces on which the destructive reactions which convert molecules like ClONO2 and HCl to more reactive ones like HOCl and Cl2. These two compounds release Cl atoms.

  One student responded to the question to ponder citing the difficulties of filtering and collecting all the ozone, the economics and increase in solar radiation, which reaches the surface and is ultimately trapped by the CO2. A further clarification was posted on the role of the stratospheric ice clouds and the ozone hole.

  Some useful Websites

  High Level Ozone
  Ozone Production and Destruction