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Chapter 2: Protecting the Ozone Layer

Chapter 2: Protecting the Ozone Layer. Why do we need to do to protect the ozone layer? Isn’t ozone hazardous to human health? Why is the ozone layer getting smaller? What can we do (if anything) to help stop the depletion of our ozone layer?. Catatan: Diambil dari berbagai sumber.

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Chapter 2: Protecting the Ozone Layer

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  1. Chapter 2: Protecting the Ozone Layer Why do we need to do to protect the ozone layer? Isn’t ozone hazardous to human health? Why is the ozone layer getting smaller? What can we do (if anything) to help stop the depletion of our ozone layer? Catatan: Diambil dari berbagai sumber

  2. Ozone Formation Energy + 3 O2 2 O3 Energy must be absorbed (endothermic) for this reaction to occur. Ozone is anallotropicform of oxygen. The name ozone comes from a Greek word meaning “to smell” An allotrope is two or more forms of the same element that differ in their chemical structure and therefore their properties. ElementAllotropes oxygen O2, O3 carbon graphite, diamond, buckminister fullerenes 2.1

  3. -The sum of the protons and neutrons. Mass number (A) O 16.00 8 -The number of protons. (nuclear charge) Atomic number (Z) 2.2 2.2

  4. The electrons in the outermostenergy levels are called valence electrons. The group number (of the representative elements) on the periodic table tells you the number of valence electrons. Group 1A: 1 valence electron 1A 8A Group 3A: 3 valence electrons 3A 6A 7A 4A 5A 2A 2.2

  5. Isotopesare two or more forms of the same element (same number of protons) whose atoms differ in number of neutrons, and hence in mass. Isotopes of carbon: C-12, C-13, C-14 also written as: 12C 13C 14C • Carbon 14 ( 6C) • Used in dating old materials • Atomic number: 6 • Atomic mass: 14 • # of neutrons = Atomic mass-Atomic # • 6 protons, 6 electrons, 8 neutrons *The average mass is very close to 12.000 b/c 6C is by far the most abundant isotope. 2.2

  6. Representing molecules withLewis structures: Consider water, H2O: • 1. Find sum of valence electrons: • 1 O atom x 6 valence electrons per atom = 6 • + 2 H atoms x 1 valence electron per atom = +2 • 8 valence electrons 2. Arrange the electrons in pairs; use whatever electron pairs needed to connect the atoms, then distribute the remaining electron pairs so that the octet rule is satisfied: 2.3

  7. Representing molecules withLewisstructures: Typical valence for selected atoms = the # of bonds an atom typically forms 2.3

  8. Representing molecules withLewisstructures: Multiple bonds Doublebond Triplebond Occasionally a single Lewis structure does not adequately represent the true structure of a molecule, so we useresonanceforms: 2.3

  9. The Nature of Light Low E High E Wavelength (l) = distance traveled between successive peaks (nm). Frequency (n) = number of waves passing a fixed point in one second (waves/s or 1/s or s-1 or Hz). 2.4

  10. The Electromagnetic Spectrum The various types of radiation seem different to our senses, yet they differ only in their respective l and n. 2.4

  11. Visible: l = 700 - 400 nm M MUDASIR Decreasingwavelength Infrared (IR) : longest of the visible spectrum, heat ray absorptions cause molecules to bend and stretch. Microwaves: cause molecules to rotate. Short l range: includes UV (ultraviolet), X-rays, and gamma rays. 2.4

  12. The wavelength and frequency of electromagnetic radiation are related by: c = ln where c = 3 x 108 m/s (the speed of light) The energy of a photon of electromagnetic radiation is calculated by: E = hn where h = 6.63 x 10-34 J.s (Plank’s constant) Energy and frequency are directly related-higher frequency means higher energy. 2.5

  13. What is the energy associated with a photon of light with a wavelength of 240 nm? E = hn C = ln E = (6.63 x 10-34 J.s) (1.3 x 1015 s-1) n = C l E = 8.6 x 10-19 J n = 3 x 108 m/s =1.3 x 1015 s-1 10-9 m 240 nm x nm UV radiation has sufficient energy to cause molecular bonds to break 2.5

  14. 2.6

  15. The Chapman Cycle l≤ l≤ A steady state condition 2.6

  16. Factors influencing steady-stateconcentration of O3 • Intensity of UV radiation • Expect O3 concentration will depend on season and sun cycle • Concentration of O2 and other reacting species • Temperature • Reaction rates (some of the steps are faster than others) • *Many of these factors depend on altitude, so O3 concentrations also depend on altitude.

  17. Biological Effects of Ultraviolet Radiation • The consequences depend primarily on: • The energy associated with the radiation. • The length of time of the exposure. • The sensitivity of the organism to that radiation. The most deadly form of skin cancer, melanoma, is linked with the intensity of UV radiation and the latitude at which you live. As l decreases, damage to DNA increases. Note the logarithmic scale on the y-axis. An Australian product uses “smart bottle” technology; bottle color changes from white to blue when exposed to UV light. 2.7

  18. O3 concentrations vs. altitude

  19. O3 concentrations over time http://earthobservatory.nasa.gov/cgi-bin/texis/webinator/printall?/Library/Ozone/ozone_2.html 1 Dobson unit (DU) = 1 O3 molecule per 1 billion (109) molecules of air 320 Du is average ozone level over the northern hemisphere 250 DU is typical at the equator “hole” is the area over Antarctica with < 220 DU

  20. Natural causes of O3 destruction 1) Reactions with water vapor and its breakdown products: H2O + UV photon  H• + •OH These free radical products can participate in reactions that convert O3 to O2: O3 + H•  O2 + •OH O3 + •OH  2O2 +H• Example of a chain reaction. Free radicals are very destructive Free radicals very reactive b/c they have unpaired e- Octet Rule is NOT satisfied

  21. 2) Reactions with stratospheric nitrogen monoxide (NO) • Formation of natural NO: N2O + O  2NO • N2O produced by microorganisms • Formation of man-made NO: N2 + O2 + energy  NO • Energy is from lightning or high-temp engines of supersonic transport airplanes (like the Concorde) that are designed to fly at altitudes of 15-20 km (the region of the ozone layer) • Lewis dot structure of NO? • Turns out that NO is also a free radical… energy Source of VOC and NO in the Upper Atmosphere leading to Tropospheric Air Pollution Chemistry • However, these natural causes do not explain the amount of ozone destruction seen in recent years…

  22. CFCs seem to be the culprit • Properties and uses of CFC • Stable • Nontoxic • Nonflammable • Low boiling point (makes it a good refrigerant) • Cheap • Widely available • Originally developed as refrigerants to replace the toxic ones that had previously been used (SO2 and ammonia—NH3) • Also used as: • Solvent for computer chips • Propellants for aerosol cans • Styrofoam • Fire extinguishers

  23. Before we knew about the ozone hole… • Two scientists, Roland & Molina, set out to study fate of atmospheric CFC molecules • The reactions they hypothesized would occur suggested that CFCs could destroy significant amounts of O3 • The destruction of O3 by CFCs later shown by scientific experiments. • ~ 5 years after release in troposphere, CFCs make their way into the stratosphere. • A high-energy photon corresponding to wavelengths  220 nm has enough energy to break the chlorine-carbon bond: CCl2F2 + photon  •CClF2 + •Cl Revolutionized modern life! Throughout the 1960s and 70s, cheap air conditioning using CFCs helped spur the booming growth of Southern cities such as Atlanta, Dallas-Fort Worth and Houston.

  24. Reactions with •Cl Removal of •Cl • Carried to lower atmosphere by winds—once in lower atmosphere, there are many other compounds for •Cl to react with. • If it encounters another •Cl the following reaction may occur: •Cl + •Cl  Cl2 • May react with other species to form stable compounds such as HCl and ClONO2 •Cl +O3 O2 + •OCl •OCl + O  •Cl + O2 O3 + O  2O2 You can see that once •Cl is created that it begins a chain reaction—it can destroy many (~ 100,000!) O3 molecules. •Cl is acting as a catalyst—it is not used up in the reaction The presence of •Cl will affect the Chapman cycle balance!!!

  25. Proof of relationship b/w ozone depletion and stratospheric Cl • 75-85% is from human activity • CFCs • Rocket fuel from space shuttles (<1%) • 15-20% is from methyl chloride • Most from natural sources and burning of biomass • ~2% - Large, explosive volcanoes • Increase amount of HCl which can be converted to •Cl • Only a few volcanoes have enough explosive power to project material into the stratosphere.

  26. Why is the hole over Antarctica? • A special mechanism appears to be in effect here, related to the fact that the lower stratosphere over the South Pole is the coldest spot on Earth. • During the Antarctic winter (June to September), a strong circumpolar wind develops in the middle to lower atmosphere • keeps warmer new air (w/ new O3) from entering • Temps can get as low as -90oC! • At temps < 80oC, clouds of ice crystals containing sulfates and nitric acid can develop in the stratosphere. • In spring, when the sun comes out, chemical reactions occurring on the surface of these cloud particles convert otherwise safe (non-ozone depleting) molecules to more reactive species. • ClONO2 & HCl converted to HOCl & Cl2 • HOCl•Cl + •OH ; Cl2 2 •Cl • UV light UV light http://earthobservatory.nasa.gov/Study/Tango/

  27. Summer conditions replenish hole • In summer, sunlight warms the atmosphere • Ice crystals evaporate • No ice crystals means the conversion of ClONO2 & HCl to more reactive species halts. • Air from lower latitudes flows into the polar regions, replenishing depleted ozone levels. • Thus, “Ozone Hole” is largely replenished by the end of summer. summer spring • Same process is developing in the Arctic

  28. The Global Response • CFCs in spray cans banned in North America in 1978, and use as foaming agents for plastics discontinued in 1990. • 1987 Montreal Protocol—established a schedule to reduce production & consumption of CFCs. • The US and 140 other countries agreed to a complete ban on CFC production after Dec. 31, 1995. • Other ozone depleting compounds to be eliminated b/w 2002-2010: • CFBr compounds (fire-fighting) • CCl4 (an industrial solvent) • CH3Br (agricultural fumigant) Global production of CFCs, 1950-2000

  29. Black-market CFCs • While the protocols set dates for the halt of production, sale of existing stockpiles will remain legal. • US govt. is promoting conversion to less harmful substitute refrigerants by imposing a tax of $5.35/lb on CFCs • Price of CFCs has risen $1/lb in 1989 to $15/lb (as of 2003) • Creates a black-market for CFCs, which are 2nd only to illegal drugs as the most lucrative illegal import. • “Freon busts” in 1997 led to the confiscation of 12 million lbs of illegal CFCs

  30. Ban on CFCs is making a difference--slowly What to use instead? • Stratospheric concentration of chlorine peaked at 4.1 ppb in the mid-90s and is now declining. • However, the stratospheric chlorine is not expected to reach 2 ppb* until 2050 or 2075. • *Antarctic ozone hole first appeared when chlorine levels reached 2 ppb. • DON’T want to go back to previous technologies (NH3, SO2 as refrigerants)! • Need to balance 3 undesirable properties and come up with the best compromise • Toxicity • Flammability • Extreme Stability

  31. Fluorocarbons (FCs) • Made of only carbon and fluorine • Non-toxic • Nonflammable • Don’t decompose in the stratosphere • Why? The C-F bond is stronger—the energy requirement to break this bond is higher than that of a UV photon. • The negative: No destruction pathways—the molecules will accumulate in the atmosphere and contribute to the greenhouse effect by absorbing IR radiation.

  32. Hydrochlorofluorocarbons (HCFCs) • Replacing one or more of the halogen atoms with hydrogen makes the compound more reactive • Compounds decompose long before reaching stratosphere • However: • Too many H atoms increases flammability • Too many Cl atoms increases toxicity

  33. Suitable HCFCs Decompose in the troposphere more readily than CFCs • HCFC-22 is the most widely-used HCFC • Used in air-conditioners and in the production of foamed fast food containers • 5% ozone depleting potential of CFC-12 • HCFC-141b used to form foam insulation • 250 million lbs produced worldwide in 1996 for this purpose • Montreal Protocol calls for a production ban of HCFCs by 2030. CFC-12

  34. Hydroflurocarbons (HFCs) H • Example: HFC-134a • Has no Cl atoms to interact with ozone • The 2 H atoms facilitate decomposition in the lower atmosphere, without making it flammable under normal conditions. C C H US Emissions http://www.eia.doe.gov/oiaf/1605/vr98rpt/chapter6.html

  35. Timeline • 1928 – CFCs invented • 1950s-1970s – Consumption and use of CFCs rises rapidly • 1971 – CFCs measured in the atmosphere • 1974 – Rowland & Molina link CFCs with ozone depletion • 1985 – First scientific assessment of stratospheric ozone levels • 1987 – “Smoking gun” evidence linking decreasing ozone levels with increasing stratospheric chlorine levels – Montreal Protocol established a schedule to reduce production & consumption of CFCs. • 1991 – Multinational fund established to provide financial and technical assistance to developing countries to enable them to comply with control measures. - B/w 1991 & 2004 $1.6 billion given to more than 100 countries • 1995 – Production ban of CFCs in US and 140 other countries goes into effect – Rowland & Molina win Nobel Prize for their work • 2030 – Production ban of HCFCs goes into effect

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