Environmental Chemistry

Environmental Chemistry
Chapter 1: Stratospheric Chemistry: The Ozone Layer Part 1 - Structure of the Atmosphere Part 2 - Composition of the Atmosphere Part 3 - Pollution of the Stratosphere Part 4 - The Ozone Hole Copyright © 2007 by DBS

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Part 1 - Structure of the Atmosphere
The ozone hole provides a classic case of the workings of science at its best: the unexpected discovery of an important effect, proposals of theories to explain it, quick mounting of a logistically difficult experimental field program to test the theories, and a blending of laboratory data, field observations, and computer models to achieve understanding, in this case within only two years Graedel and Crutzen, 1993

Our Atmosphere - Review
Image from Temperature rises at tropopause but temperature is not really meaningful since pressure has dropped (fewer molecules). The atmosphere is dividided vertically into layers based on temperature Source:

Troposphere [1000 - 200 mb (29.92 in to 5.92 in)]
From 0 to 12 km (7 mi) [ mb (29.92 in to 5.92 in)] Temperature decreases 2. Winds increase with height to the jet stream Moisture decreases Sun’s heat warms the surface and is transported up by convection (mixed by updrafts and downdrafts) Image from

Stratosphere O3 is central species UV induced photochemistry dominates
Inversion Meteorology is influenced by heat generated Stratosphere is stable Earth’s ‘sunscreen’ Turco, 2002

Chemistry of Sunlight…the good
Chlorophyll transforms CO2 and H2O into O2 and carbohydrates xCO2 + yH2O + light → xO2 + Cx(H2O)y Light energy is converted into chemical energy which drives the biochemical reactions in plants

Chemistry of Sunlight…the bad
Any tan is a sign of skin damage Overexposure to UV-A and UV-B Tanning - skin produces a pigment to protect itself Damage may be intermediate or long-term: sunburn, rashes, cell and tissue damage, premature wrinkling and skin cancer Source:

Chemistry of Sunlight…the ugly
UV photons are energetic and can induce chemical reactions in biological systems Excess UV can influence plant growth and alter the ecological balance Where would the US vs. UK fall on this chart? Incidence of nonmelanoma skin cancer per 100,000 males vs. light intensity

Allotropes of Oxygen Turco, 2002 Allotropes of Oxygen

Light Absorption by Molecules
IMPORTANT SW absorbers: O2 absorbs 70 – 250 nm O3 absorbs 200 – 300 nm Source: Graedel and Crutzen, 1993 Write the chemical reactions Source: Graedel and Crutzen, 1993

SW Absorption O2 and O3 are SW filters Absorption weakens at 300 nm

Solar Flux Phytoplankton Melanoma Source: Jacob, 1999
Chemical Kinetics and Photochemical Data for Use in Stratospheric Modeling - JPL Publication97-4 Melanoma

Energy and Wavelength Why do molecules only absorb at certain wavelengths?

E = hν or E = hc / λ (since c = ν λ)
Energy and Wavelength In 1905 Einstein demonstrated the photoelectric effect which shows that photons of light have particle characteristics e- were knocked free from a metal by photons with energy E E is related to the frequency, ν, and wavelength, λ by E = hν or E = hc / λ (since c = ν λ) Where h is plank’s constant (6.626 x J s) and c = speed of light ( x 108 m/s) Only photons of a high enough frequency (above a threshold value) could knock an e- free

E = hc/ λ = 1.2 x 105 / λ kJ nm/mol (λ is in nm)
Since hc = (6.626 x J s) x ( x 108 m/s) = x kJ nm Or 1.2 x105 kJ nm/mol (N = x 1023 photons) Using this expression we can obtain the energy of a photon at a given wavelength E = hc/ λ = 1.2 x 105 / λ kJ nm/mol (λ is in nm) Examples What is the energy of a 300 nm photon? E = hc/λ =1.2 x105 / 300 = 400 kJ/mol or 4.15 eV ( kJ/mol = 1eV)

What λ is 1 eV? E = hc / λ or λ = hc / E
 λ = 1.2 x105 (kJ nm mol-1)/ (kJ mol-1) λ = 1240 nm 1 eV of photon energy corresponds to 1240 nm Determine the energy of photons of any given wavelength E = 1240 / λ 620 nm → 2eV nm → 3 eV 310 → 4 eV

Sunscreens TiO2, ZnO Active ingredients inorganic and organic
Protect against UV-B in nm range Absorb UV light and dissipate as heat Some offer UV-A nm protection (more penetrating) TiO2, ZnO

Ozone History 1785, Martinus Van Marum noted “the odor of electrical matter” in the description of the discharge of air Officially named as a chemical in 1840 by Christian Schönbein, after he noted that it had a smell that was similar to that of phosphorus when exposed to air It was soon realized that O3 was a good disinfectant. Otto was first to market a water purifier based on ozone Christian Schönbein

O2 + O3 as Natural UV Filter
absorbs transmits Absorption of UV-C and UV-B by O3

Production of O3 Above the stratosphere most O2 absorbs UV-C and photodissociates O2 + hν (λ < 240 nm) → O2* → 2O Due to high concentration of O2 any O atoms get converted into O3 O2 + hν → 2O 2 [O2 + O + M → O3 + M] Net: 3O2 + hν → 2O3 Ozone is constantly being formed in the earth's atmosphere by the action of the sun's ultraviolet radiation on oxygen molecules. Ultraviolet light splits the molecules apart by breaking the bonds between the atoms. A highly reactive free oxygen atom then collides with another oxygen molecule to form an ozone molecule. Because ozone is unstable, ultraviolet light quickly breaks it up, and the process begins again. + heat

Question For the dissociation of O2 ΔH0 = 495 kJ/mol
What is the longest wavelength of photons capable of dissociating O2? O2 + hν → 2O λ = hc/E = 1.2 x 105 kJ nm mol-1 / 495 kJ mol-1 = 241 nm UV light with wavelength < 241 nm can dissociate O2

Measurement In 1923, Gordon Dobson developed the first spectrometer to measure ozone in the atmosphere, and he characterized its latitudinal seasonal variability Source:

Ozone in the Atmoshpere
About 90% of the Earth’s ozone is in the stratosphere ‘Good’ O3 ptotects life from harmful UV In the troposphere O3 is a pollutant and is termed ‘bad’ <1 ppm

Question [O3] = 2 x 1012 molecules cm-3 = 3.3 x 10-12 mols cm-3
Express [O3] = 2.0 x 1012 molecules cm-3 as a volume mixing ratio (ppmv) [O3] = 2 x 1012 molecules cm-3 = 3.3 x mols cm-3 = 3.3 x mols cm-3 x 48 g/mol = 1.6 x g cm-3 = 1.6 x 10-7 mg cm-3 x (1 x 106 cm3 / m3) = 0.16 mg m-3 = 0.16 mg m-3 x 24.5 / 48 g mol-1 = ppmv From Graedel and Crutzen (1993), Chapter 1

Destruction of O3 Because most oxygen is present as O2, it is clear that there must be a process to reconvert O3 to O2… O3 + hν (λ < 325 nm) → O + O2 O + O3 → 2O2 Net: 2O3 + hν → 3O2 Ozone is constantly being formed in the earth's atmosphere by the action of the sun's ultraviolet radiation on oxygen molecules. Ultraviolet light splits the molecules apart by breaking the bonds between the atoms. A highly reactive free oxygen atom then collides with another oxygen molecule to form an ozone molecule. Because ozone is unstable, ultraviolet light quickly breaks it up, and the process begins again.

Chapman Theory Describes how the various forms of oxygen are converted fromone to another and explains why the highest concentration of O3 is found between km Constantly formed, decomposed and reformed Rates differ depending on altitude Why is O3 not formed below the stratosphere?

Part 2 – Composition of the Atmosphere

Review - Our Atmosphere
The atmosphere is dividided vertically into layers based on temperature Image from Source:

Review - Intensity of Sunlight
Absorption by Oxygen and Ozone in the atmosphere filters the UV component reaching the earth.

Review - O2 and O3 as Filters
O3 absorbs the most energetic UV light (UV-B and UV-C) which cause biological damage UV-B absorption UV-C absorption

Review - Creation of O3 in Stratosphere
Atmospheric O3 constantly formed by action of UV radiation on oxygen molecules UV light breaks the bonds between the atoms Highly reactive O atom collides with another O2 molecule to form an O3 molecule Because ozone is unstable, UV light quickly breaks it up and the process begins again

Review - Ozone in the Atmoshpere
About 90% of the Earth’s ozone is in the stratosphere ‘Good’ O3 ptotects life from harmful UV In the troposphere O3 is a pollutant and is termed ‘bad’ O3

O2 + hν → O + O ΔH = 495 kJ/mol (<241 nm) (1)
Chapman Theory Above stratosphere oxygen absorbs UV-C and exists as O atoms O2 + hν → O + O ΔH = 495 kJ/mol (<241 nm) (1) Oxygen atom could react with oxygen molecule to form O3 O + O2 + M → O3 ΔH = -100 kJ/mol (2) O3 formed could react with O atoms or absorb solar radiation O3 + hν → O2 + O (<320 nm) (3) O + O3 → 2O2 ΔH = -390 kJ/mol (4) - A third molecule ‘M’ (N2 or H2O) facilitates as a heat energy carrier (is not required when there is more than one molecule produced) - Enthalpies show a great deal of heat is generated - Because of the temperature inversion vertical mixing of air is slow in the stratosphere + O3 - O3

Chapman Theory

Chapman Animation

Chapman Theory Photostationary state is established O3 + hν ↔ O2 + O
Net effect: one form of ‘odd oxygen’ is converted into another (O3 to O) At night back reaction dominates At day forward reaction dominates (hν) JO3

For about 40 years, it was generally accepted that this sequence explained the full cycle of stratospheric ozone…

K1[A]=k2[C] or [A]/[C] = k1/k2
Steady-State Rate of change = formation – destruction A → B (1) B → C (2) Rate of change of B = k1[A] – k2[C] At steady-state rate of change = 0  Rate of formation = rate of destruction K1[A]=k2[C] or [A]/[C] = k1/k2

Equilibrium vs Steady-State

SS Analysis for O3: Chapman Mechanism
O2 + hν → 2O rate 1 = JO2[O2] slow O + O2 + M → O3 + M rate 2 = k2[O][O2][M] fast O3 + hν → O2 + O rate 3 = JO3[O3] fast O3 + O → 2O2 rate 4 = k4[O3][O] slow Since O is produced and consumed in all reactions we can assume it is in the SS d[O] = 2(JO2[O2]) - k2[O][O2][M] + JO3[O3] - k4[O3][O] = 0 (A) dt Similarly, d[O3] = k2[O][O2][M] – JO3[O3] – k4[O3][O] = 0 (B) d[O + O3] = d[Ox] = 2(JO2[O2]) - 2 k4[O3][O] ) = 0 (C) dt dt Control [O] This is a simplified version of this: However, by the mid 1960’s, especially following a study by Benson and Axworthy (3), it became clear that reaction R4 is much too slow to balance the production of “odd oxygen” by reaction R1 (see Figure 1). (production-loss)

SS Analysis for O3: Chapman Mechanism
[O][O3]ss = JO2[O2]ss (C) k4 (A) – (B): 2(JO2[O2]) – 2(k2[O][O2][M]) + 2(JO3[O3]) = 0 Or JO3[O3] = JO2[O2] – k2[O][O2][M] Since rate 1 << rate 2 and rate 3, Rate 3 = rate 2 JO3[O3] = k2[O][O2][M] Or [O3]/[O] = k2[O2][M]/JO3 (D) (C) x (D): [O3]2 = JO2k2 [O2]2[M] JO3 k4 Or [O3]ss/[O2]ss = [M]0.5(JO2k2 / JO3 k4)0.5 (E)

[O3]ss/[O2]ss = [M]0.5(JO2k2 / JO3k4)0.5 (E)
The steady-state ratio of [O3] to [O2] depends on: Square root of air density Proportional to square root of JO2 x k2 corresponding to O and O3 production Inversely proportional to the product JO3 x k4 Experimental values suggest O3 maximum around 10-4 M M (and [O2] decreases with increasing altitude Solar intensity  JO2 decreases with decreasing altitude If [O2] = 0.21[M] equation reduces to [O3] = 0.21(k1/k4)^0.5 x (JO2/JO3)^0.5 x [M]^3/2 The maximum reflects largely the vertical dependence of Ox production by (R1), JO2[O2], which we have seen is the effective source for O3

Concentration of Atomic Oxygen
From (E) and (C) [O]ss = (JO2JO3/k2k4)0.5/[M]0.5 [O]ss increases with increasing altitude as [M] declines JO2 and JO3 increase with altitude  O atoms dominate over O3 at higher altitudes (> 50 km) Wayne Fig. 4.3 Concentration of atomic oxygen and ozone vs. altitude

Concentration of Atomic Oxygen
[O3]ss/[O2]ss = [M]0.5(JO2k2 / JO3k4)0.5 At 30 km JO2 = 6 x s-1 JO3 = 1 x 10-3 s-1 k2 = 4.5 x cm6 s-1 k4 = 1 x cm3 s-1 [O3]SS = [(6 x s-1) x (4.5 x cm6 s-1)] 0.5 x [M]0.5 x [O2]ss [(1 x 10-3 s-1) x (1 x cm3 s-1)] = 6 x 1012 molecules cm-3 [O3]SS ~ 0.2 ppm This is almost a factor of 4 above the true concentration! What is wrong? There must be ozone sinks missing. [O2] = 0.21[M] [M] = 3.3 x 1017 molecules/cm3

Problem – Chapman Mechanism Predicts More O3 than Observed!
Model curve is displaced by factor of 2 The actual destruction of O3 is greater than predicted…why? What’s missing? Faster O3 + O → O2 + O2

Catalytic Loss Cycles X can be either NO, OH or Cl X is recycled
These cycles compete with production by sunlight to produce the O3 distribution NB: Both NOx and HOx cycles are natural cycles…pollution may add

HOx Cycle Important in the lower stratosphere (low [O])
OH• is produced when CH4 and H2O escapes into stratosphere and reacts catalytically with O* (also from photolysis of H2O) H2O + O* → 2•OH CH4 + O* → •OH + •CH3 •OH + O3 → HO2• + O2 HO2• + O3 → •OH + 2O2 Net: 2O3 → 3O2 Speeds up reaction 4 of Chapman’s mechanism HOx accounts for 15% of net O3 removal rate There must be other reactions that remove O3

NOx Cycle 1970 Crutzen showed that NOx (NO and NO2) react catalytically with O3, reducing overall odd O NO is produced when stable N2O rises from toposphere to stratosphere and reacts with excited oxygen atom N2O + O* → 2NO NO + O3 → NO2 + O2 NO2 + O → NO + O2 Net: O + O3 → 2O2 Rate >>> Chapman mechanism (4) (NO requires O atom for regeneration) Referred to as Mechanism I NO N2O MO’s N2O

NOx Cycle Converts two odd oxygen species into 2O2 and reduces overall odd oxygen NO acts as a catalyst and speeds up the O + O3 reaction NOx controls the overall level of odd oxygen in the stratosphere (~60% of net O3 removal is due to NOx)

Rate of reaction between NO and O3
NO + O3 → NO2 + O2 Rate = k [NO][O3] At -30 °C, k = 6 x cm3 molecules-1s-1 If [NO] = 1 x 109 molecules cm-3 and [O3] = 5 x 1012 molecules cm-3 Rate = (6 x cm3 molecules-1s-1) x (1 x 109 molecules cm-3) x (5 x 1012 molecules cm-3) = (3x107 molecules cm-3s-1)

NOx Cycle (simplified)
Removes O3 SINK HNO3 (inert) i.e. inactive until transported SOURCE N2O + O → 2NO RAIN OUT RESERVOIR Removal negligable (few reactions except O) N2O, N2 Agriculture NO3-, NO2- Mankind can alter stratospheric O3 without leaving the ground

Interaction of HOx and NOx Cycles
(a) •OH + NO2 → HNO3 Removes two catalytically active species, Protects O3 in lower stratosphere (b) NO + •HO2 → NO2 + •OH Returns odd oxyen via NO2+ hν → NO + O

Problem 1-10 = 1.57 x 10-13 mols yr-1 cm-3
At an altitude of 35 km [O*] = 100 & [CH4] = 1 x 1011 molecules cm-3 k = 3 x cm3 molecules-1 s-1 (i) What is the rate of destruction of methane? (ii) How many grams per year per cm3? Rate=k[CH4][O*] Rate = (3 x cm3 molecules-1 s-1) x (1 x 1011 molecules cm-3) x (100 molecules cm-3) = 3 x 103 molecules cm-3 s-1 Rate of destruction /yr = (3 x 103 molecules cm-3 s-1) x (3.15 x 107 s yr-1) = 9.47 x 1010 molecules cm-3 yr-1 = 9.47 x 1010 molecules yr-1 cm-3/6.023x1023 molecules mol-1 = 1.57 x mols yr-1 cm-3 = 1.57 x10-13 mols yr-1 cm-3 x 16 g/mol = 2.52 x g CH4 yr-1 cm-3 (Mol. Wt. of CH4 is 16)

CH3-Cl + •OH → oxidation products + •Cl
ClOx cycle Formation of Cl atoms from methyl chloride (naturally occurring) CH3-Cl + hν → •CH3 + •Cl CH3-Cl + •OH → oxidation products + •Cl Cl atom initiated O3 destruction (very fast) •Cl + O3 → O2 + •ClO •ClO + O → •Cl + O2 Net: O3 + O → 2O2 Note that each Cl atom can destroy many thousands of O3 molecules More efficient than NOx since chain length (~5000) is longer, main sink is HCl Stolarski & Cicerone (1974); Wofsy & McElroy (1974) “ClOx”

CH3-Br + •OH → oxidation products + •Br
BrOx cycle Formation of Br atoms from methyl bromide CH3-Br + hν → •CH3 + •Br CH3-Br + •OH → oxidation products + •Br Br atom initiates O3 destruction •Br + O3 → O2 + •BrO •BrO + O → •Br + O2 Net: O3 + O → 2O2

Interaction with Other Cycles
Free radicals are short-lived and are readily converted into stable forms – so called reservoir species that are catalytically inactive •ClO + NO2 ⇌ ClONO (chlorine nitrate) •Cl + CH4 ⇌ HCl + •CH3 HCl and ClONO2 are inactive since they do not react directly with O3…chlorine reservoirs…transported out of stratosphere? When it was realized in 1980s that the chlorine in the atmosphere exists in the inactive form , the predicted loss of ozone in the stratosphere was lowered sunlight

Interaction with Other Cycles
Why no F cycle? OH + HF endothermic Unlike Cl, inactive HBr and bromine nitrate are decomposed by sunlight…active free radical Br and BrO remains making Br a more efficient ozone destroyer However, inactive form of chlorine (reservoir or chlorine nitrate and hydrogen chloride) can become active and destroy massive amounts of ozone! – a discovery that was not made until 1980’s

Summary: Formation vs Destruction
The rate of O3 production from oxygen depends on light intensity and oxygen concentration The rate of O3 destruction is more complex – depends on sunlight intensity, O3 concentration and catalyst concentration The rate of formation and destruction keeps the O3 balance in the stratosphere (–steady state condition) If the rate of destruction is increased by the introduction of additional molecules of a catalyst, the steady state concentration of O3 will decrease However, this decrease will alter the balance but will not completely destroy ozone from the atmosphere Penetration of UV-C at lower altitude will continue to produce O3 – a self healing process

Part 3 – Pollution of the Stratosphere
Stratosphere is vulnerable due to its low density and stability against mixing Any reduction in O3 increases UV-B undesirable biological effects Concern of pollution by SST’s in 1970’s Effect of water vapor (HOx cycle) and NOx is negligable More concern over growing use of fertilizers (N2O)

CFnCl4-n + hν → CFnCl3-n + Cl
Chlorofluorocarbons Chlorofluorocarbons (CFC’s, Freons,…) non-toxic, non-flammable, non-carcinogenic. e.g. CFC-11(trichlorofluoromethane - CFCl3) CFC-12 (dichloro-difluoromethane - CF2Cl2) Inert in troposphere (no sinks!), not soluble in water Present conc. ~10-10 by volume ~ integrated world production Sink for CFC’s is the stratosphere CFnCl4-n + hν → CFnCl3-n + Cl Photolysed at λ < 220nm Enter ClOx cycle, may become inactive

Link Between Freon and O3 Depletion
Two important contributors: Lovelock demonstrated that man-made, chemically inert, CFC gases had spread throughout the atmosphere Stolarski and Cicerone showed that free chlorine atoms can decompose ozone catalytically (similarly to NOx) “…are unusually stable chemically and only slightly soluble in water and might therefore persist and accumulate in the atmosphere …The presence of these compounds constitutes no conceivable hazard.” Source: Distribution of CCl3F in and over the N. and S. Atlantic Ocean, Nature, Vol. 241, 1973 Concentrations (pptv) of CCl3F

Spray Cans + Refrigerators Damage the O3 Layer
1974, Molina and Rowland (Nature) proposed that CFC’s would react with UV light to produce free Cl atoms Predicted tens of percents of O3 loss CFCl3 + hν → CFCl2 + •Cl 1614 citations even with typo!

Cl + O3 → ClO + O2 ClO + O → Cl + O2 ClO + NO → Cl + NO2 ClO + NO2 + M → ClONO2 ClONO2 + hν → Cl + NO3 → ClO + NO2 Cl + HO2 → HCl + O2 Cl + CH4 → HCl + CH3 ClO + OH → HCl + CH3 → Cl + HO2 HCl + OH → H2O + Cl ClO + HO2 → HOCl + O2 HOCl + hν → Cl + OH What is really going on? Couples ClOx to NOx Couples ClOx to HOx

And NOx is coupled to HOx!
OH + O3 → HO2 + O2 OH + O → O2 + H OH + CO → CO2 + H H + O2 + M → HO2 + M HO2 + O → OH + O2 HO2 + O3 → OH + 2O2 OH + HO2 → H2O+ O2 HO2 + HO2 → H2O2 + O2 HO2 + NO → NO2 + OH OH + NO2 + M HNO3 + M HNO3 + OH → H2O + NO3 HNO3 + hν → OH + NO2 HO2 + NO2 → HO2NO2 + M HO2NO2 + hν → OH + NO3 + halogen reactions… And NOx is coupled to HOx! Couples HOx to NOx

The World Breathes a Sigh of Relief… in the Form of CFC’s
1978 ban on certain CFC use

The Big Surprise of 1985 Isolated local concern or global problem?
Farman et al. revealed a dramatic and unpredicted decline in stratospheric O3 in a surprising location Antarctica Shocked the world Showed dramatic decline in springtime O3 starting in 1970’s 30% by 1985 70% by 2000 Farman, J. C.;Gardiner, B. G.;Shanklin, J. D., Large Losses of Total Ozone in Antarctica Reveal Seasonal Clox/Nox Interaction. Nature, 1985, 315, Isolated local concern or global problem? Chemical explanation? Physical explanation? Min O3 at Antarctic in Spring (Sep-Nov)

Convention for the protection of the ozone layer

Montreal Protocol 1985: scientific concerns about damage to the ozone layer prompted governments to adopt the Vienna Convention on the Protection of the O3 Layer 1987: international negotiators met again to adopt legally binding commitments in the Montreal Protocol on Substances that Deplete the O3 Layer Developing countries have agreed to freeze most CFC consumption as of 1 July 1999 based on averages, to reduce this consumption by 50% by 1 January 2005 and to fully eliminate these CFCs by 1 January 2010

How Effective is the MP?

How Effective is the MP? Illegal CFCs imperil the ozone layer
17 December 2005 Duncan Graham-Rowe Magazine issue EFFORTS to speed up the phasing out of CFCs in China may be unwittingly encouraging illegal trade in the ozone-depleting chemicals. If this trade continues unabated, repair of the Antarctic ozone hole will be set back years. An undercover investigation has revealed how lax trade controls are fostering a booming black market in Chinese CFCs. This news comes hard on the heels of a study suggesting that even without illegal CFC trading, the springtime hole in the Antarctic ozone layer is likely to close 15 years later than had been predicted.

CFC Replacements CFCs and CCl4 have no tropospheric sinks (not soluble in water/rain), not decomposed by UV-A or visible light HCFCs contain H atoms bonded to C atom. Consequently a majority of the molecules are removed in the troposphere by the hydrogen abstraction reaction of hydroxyl radicals CHF2Cl (HCFC-22) the current replacement for refrigerator coolants Because of the presence of chlorine, it has a ozone-reducing potential. It decomposes quickly and chlorine atoms released in the atmosphere have greater potential for short term ozone destruction.

2C, 3H, 1F, and 2Cl: HCFC-141b = C2H3FCl2
The prefix describes what kinds of atoms are present the next step is to calculate the number of each type of atom. Add 90 to the number; the result shows the number of C, H, and F atoms. For HCFC-141b: = #C #H #F Next decipher the number of Cl atoms. All of these chemicals are saturated; that is, they contain only single bonds. The number of bonds available in a carbon-based molecule is 2n(C) + 2. Thus, for HCFC-141b, which has n=2 carbon atoms, there are 6 bonds. Cl atoms occupy bonds remaining after F and H atoms 2C, 3H, 1F, and 2Cl: HCFC-141b = C2H3FCl2 Notice that the HCFC designation (hydro chloro fluoro carbon) is a good double-check on the decoding; this molecule does, indeed, contain H, Cl, F, and C. The "b" at the end describes how these atoms are arranged; different "isomers" contain the same atoms, but they are arranged differently

Other O3 Depleting Substances
CCl4 (carbon tetrachloride) Used as dry cleaning solvent CH3-CCl3 (methyl chloroform) Cleaning agent for metals Halons Bromine containing hydrogen free substances such as CF3Br and CF2BrCl Methyl bromide is used as a pesticide, fumigate crops (40% of world use is in the US) They rise to stratosphere and decompose to produce Br. – an efficient catalyst for O3 depletion Developing countries have agreed to reduce consumption, but not complete elimination. Meantime China & Korea have increased their production!

HFC’s CH2F-CF3 called HFC-134a is considered replacement for CFC-12 (CF2Cl2) - atmospheric lifetime in decades Hydrofluorocarbons (HFCs) are compounds containing carbon, hydrogen, and fluorine. Because the HFCs contain no chlorine they do not directly affect stratospheric ozone Oxidation of HFCs by the hydroxyl radical is believed to be the major destruction pathway for HFCs in the atmosphere. Although it is believed HFCs will not deplete ozone within the stratosphere, this class of compounds has other adverse environmental effects e.g. formation of trifluoroacetic acid in rain

Summary of important points
Gases that are long-lived in the troposphere will eventually reach the stratosphere, where they nearly all break down (‘oxidize’) to produce highly reactive radicals that catalytically destroy ozone. It doesn’t matter where these gases originate from – the troposphere is the great homogenizer. The 1995 Nobel Prize in Chemistry was awarded to Crutzen, Molina, and Rowland for recognizing the importance of this concept The radical ‘families’ are highly coupled – changes in abundances of one family will result in changes in the others. Thus, the system is non-linear (although reasonably well behaved). However, it means that you can’t just scale ozone losses with emissions. A ‘simple’ stratospheric model has dozens of chemical species and hundreds of chemical reactions Having a good idea isn’t good enough. It takes a lot of measurements to prove your point – or a global crisis…

Question Deduce the formula for the compounds with code numbers 12, 113, 123, 134 a = 102 C=1, H=0, and F= 2, ( Cl=2) → CF2Cl2 b = 203 C=2, H=0, and F= 3, ( Cl=3) → C2F3Cl3 c = 213 C=2, H=1, and F= 3, ( Cl=2) → C2HF3Cl2 d = 224 C=2, H=2, and F= 4, ( Cl=0) → C2H2F4

Part 4 - Ozone Hole http://hq.unep.org/ozone/

TOMS http://jwocky.gsfc.nasa.gov Indirect measurements
Measures O3 by mapping UV light emitted by the Sun to that scattered from the Earth's atmosphere back to the satellite O3 is inferred from Earth’s albedo Shows strong spatial variability Low around equator, high in mid-latitude (why) Very low at Antarctic (especially in September/October)

What is a Dobson Unit? 1 DU is the number of molecules of O3 required to create a layer of O mm (0.001 cm) thick at 0 °C and 1 atm A column of air with an O3 concentration of 1 DU would contain about 2.69 x 1016 O3 molecules cm-2 at the base of the column Over the Earth’s surface, the O3 layer’s average thickness is about 300 DU (3 mm layer) if brought to sea level O3 “Hole” [O3] ~100 DU

Determination of O3 Concentration
What is the total mass of ozone corresponding to 350 DU? V(O3) = 4/3π [(r+d)3 -r3] [(r+d)3 - r3] = [r3 + d3 + 3rd2 + 3r2d - r3] ~ 3r2d …since r>>d V(O3) = 4/3 π [3r2d] = 4 π r2d = 4 x 3.14 x (6.4 x 106)2 x 3.5 x 10-3 = 1.8 x 1012 m3 or 1.8 x 1015 L 350 DU = 3.5 mm or 3.5x10-3 m r = 6400 km or 6.4x106 m n = PV/RT = 1.0 atm x 1.8 x 1015 L / L atm mol-1K-1 x 273 K = 8 x 1013 moles = 8 x 1013 mols x 48 g mol-1 = 4 x 1015 g

1 DU at 1 atm… d = thickness of O3 (1 DU=10-5 m), r = 6.4 x 106 m
From Ideal gas law, PV = nRT = NkBT N = number of molecules kB=R/NA Boltzmann’s const, kB=1.381 x JK-1 NO3 = [(Pstp x V) /(kB.Tstp)] = [(Pstp x 4 πr 2 d) /(kB.Tstp)] Dividing both sides by 4 πr 2 NO3/4πr 2 = [(Pstp x d)/(kB.Tstp)] = 2.69 x 1020 molecules m-2 From before: V = 4 πr2d If all ozone in the atmosphere were spread out around the earth in a homogeneous spherical shell at standard temperature, Tstp ( K) and standard pressure, Pstp ( Pa), one DU is equivalent to a horizontal density (number of ozone molecules per unit area) of 2.69 x1020 molecules per sqaure meter

O3 in the Stratosphere O2 → 2O (1) O + O2 + M → O3 + M (2)
From Steady-State analysis (Chapman mechanism) [O3]ss/[O2]ss = [M]0.5(JO2k2 / JO3k4)0.5 Continuous formation and destruction of Ozone in sunlight is supposed to keep its concentration steady in the stratosphere

Catalysts •X + O3 → •XO + O2 •XO + O → •X + O2 O3 + O → O2 + O2
O3 + hν → O2 + O N2O + O → NO + NO H2O + O → •OH + •OH CFxCly + hν → + O2 + hν.... → CO2 + xF + yCl• SSTs – H2O, NOx Chlorine – volcanoes, space shuttle N2O from fertilizing H2O from CH4 oxidation CFCs – refrigerants, propellants, foam blowing, etc.

What is the Ozone Hole? Occurs at the beginning of Southern Hemisphere spring (August-October) The average concentration of O3 in the atmosphere is about 300 Dobson Units Not a “hole” but a region of depleted O3 over the Antarctic Any area where O3 < 220 DU is part of the O3 hole Image from: Ozone is ‘thinning’ out

Ozone Destruction In september destruction occurs ~2 % per day
Most O3 is wiped out at km altitude Could not be explained by natural cycles since under low light conditions [O] is too low (UV-C required to produce O does not penetrate past 20 km)

Post 1980 O3 levels decreasing Around the globe
Trends in Global O3 Trends are obvious Post 1980 O3 levels decreasing Around the globe Solomon (1999) Figure 1 Source: Solomon, 1999

Arctic O3 ‘Dimple’ Winter of O3 destruction ~30% due to record cold temperatures UV levels remained near normal through the winter, however, because unusual weather conditions brought O3 from the Earth’s O3 -rich mid-latitudes to the pole to fill in the gaps left by the extreme O3 depletion Greenhouse gases have a net cooling effect because they radiate away the heat generated from the absorption of UV by O3 Artic weather is more stable than antarctic Feb, march and March 3/11/05

Mid-lattitudes Slow, steady decline, of about 3% per decade during the past twenty years Enhanced by volcanic eruptions (Mt. Pinatubo) Kerr, 2002

Antarctic Hole Size and Minimum O3
Size: This image shows the growth of the average area of the ozone hole from 1979 to 2004 using data from Version 8 of the TOMS algorithm. We define the ozone hole as the area for which ozone is less than 220 Dobson Units, a value rarely seen under normal conditions. It shows that the ozone this low hardly occured at all in 1980, but by year 2000 covered an area larger than North America, 26.5 million square kilometers. Minima : This image shows the lowest value of ozone measured by TOMS each year in the ozone hole. Global average ozone is about 300 Dobson Units. Before 1980 ozone less than 200 Dobson Units was rarely seen. In recent years ozone near 100 Dobson Units has become normal in the ozone hole. Ozone in the year 2002 ozone hole was higher than we have come to expect because of unusually high temperatures in the Antarctic stratosphere. NASA FACTS Cf. Fig. 1-2, 1-3

Summary Turco, 2002

The Smoking Gun! O3 and •ClO are anticorrelated
Zurer, 1988 (Baird graph) Anderson, J.G., Toohey, D.W., and Brune, W.H. (1991) Free Radicals within the Antarctic Vortex: The Role of CFC’s in Antarctic Ozone Loss. Science, Vol. 251, pp Concentration of Stratospheric O3 and •ClO vs Latitude

The New Catalyic Cycle Reactions Responsible for the Hole
Step 1: •Cl + O3 → ClO• + O2 Step 2: 2ClO• → Cl-O-O-Cl Step 2b: Cl-O-O-Cl → •Cl + ClOO Step 2c: ClOO → •Cl + O2 Step 2 net: 2ClO• → ClOOCl + hν → 2 •Cl + O2 Step 1 and 2 represent Mech II Occurs when [O] (needed for Mech I) is low controls season Net: 2O3 → 3O2 The key behind discovery of this catalytic cycle was the laboratory observation that photolysis of the ClO dimer (ClOOCl) takes place at the O-Cl bond rather than at the weaker O-O bond. It was previously expected that photolysis would take place at the O-O bond, regenerating ClO and leading to a null cycle. One molecule of chlorine can degrade over 100,000 molecules of ozone before it is removed from the stratosphere or becomes part of an inactive compound These inactive compounds, for example ClONO2, are collectively called 'reservoirs'. They hold chlorine in an inactive form but can release an active chlorine when struck by sunlight Nearly 75% of the ozone depletion in the antartica occurs by this mechanism (Cl. As a catalyst)

Reservoirs •Cl + CH4 → HCl + •CH3 ClO• + NO2 + M → ClONO2
ClONO2 + hv → ClO• + NO2 Turco, 2002 If it wasn’t for these reservoir species entire O3 layer would erode

Why are ClO Concentrations So High?
During Polar winter Special vortex conditions + Low temperature Denitrification of ClONO2 @H2O crystal Cl2 + HNO3 sunlight Stratospheric ‘containment vessel’ over S. pole •Cl

Activation of Cl On Ice Particles (Polar Stratospheric Clouds)
Cl resides in stable "reservoir" compounds, HCl and chlorine nitrate (ClONO2) PSC’s ‘denitrify’ (remove NO2 from the atmosphere) as HNO3, which prevents the newly formed ClO from being converted back into ClONO2 HCl + ClONO2 → Cl2 + HNO3 Cl2 + hν → 2 •Cl The crystals persist in the polar season even in springtime due to low temperature in the lower stratosphere (-80 °C) Exposure of sunlight in the early spring initiates destruction of O3 Heterogeneous reactions denitrify the stratosphere

HCl + ClONO2 → Cl2 + HNO3 Cl2 + hν → 2 •Cl

Sink for CFC’s…the Stratosphere
Actual and projected concentration of stratosphereic chlorine (bromine as chlorine equivalent) versus time Contributions from various gases

Future Antarctic and mid-latitude depletion will likely disappear around 2040 Unusually cold Arctic winter-spring seasons of recent years stand as a critical challenge to our understanding that could affect these projections Eruptions in the current altered chemical state also are a source of great uncertainty

What if anything is scientifically incorrect in this picture?
1. UV wavelengths of SW energy will enter the troposphere – more UV radiation leads directly to an increase in the Greenhouse Effect. 2. If stratospheric ozone is depleted, then the Greenhouse Effect in the stratosphere will be reduced, leading to global cooling. 3. Stratospheric Ozone depletion will not cause Global Warming from an enhanced Greenhouse effect directly. What if anything is scientifically incorrect in this picture?

Movies CNN - A Hole in the Sky (1992)
HORIZON - Time of Darkness (1992) NOVA - The Hole in the Sky (1987) Ozone - The Hole Story (1992)

Further Reading Biever, G. (2003) Bring Me Sunshine. New Scientist ( 9 August), pp Molina and Rowland, (1974) Stratospheric Sink for Chlorofluromethanes: Chlorine Atom-Catalyzed Destruction of Ozone. Nature, Vol. 249, pp Montzka, S.A. et al. (1999) Present and Future Trends in the Atmospheric Burden of Ozone Depleting Halogens. Nature, Vol. 398, pp Solomon, S. (1999) Stratospheric Ozone Depletion: A Review of Concepts and History. Journal of Geophysics, Vol. 37, pp Taylor, J.S (1990) DNA, Sunlight, and Skin Cancer. Jounal of Chemical Education, Vol. 67, pp Walker, G. (2000) The Hole Story. New Scientist (25 March), pp Folkins, I. and Brasseur, G. (1992) The Chemical Mechanisms Behind Ozone Depletion. Chemistry and Industry, pp Phillips, P. (1990) Ozone in the atmosphere: I The upper atmosphere. School Science Review, Vol. 71, pp Gribbin, J. (1988) Inside science: The ozone layer. New Scientist. Vol. 118, pp Scott, E. and Rowland, F.S. (1987) Chlorofluorocarbons and stratospheric ozone. Journal of Chemical Education. Vol. 64, No. 5, pp

Text Books Campbell, I.M. (1995) Energy and the Atmosphere, first edition. Wiley. Graedel, T.E. and Crutzen, P.J. (1993) Atmospheric Change: An Earth System perspective. Freeman. Wayne, R.P (1991) Chemistry of Atmospheres, second edition. Oxford. Jacob, D.J. (1999) Introduction to Atmospheric Chemistry. Princeton University Press.

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