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Section 7Stratospheric (Ozone) chemistry
1840: Discovery in Munich by German chemist C.F. Schönbein
ʻozoneʼ had a distinct smell and was almost colourless (ozone gas in high concentrations is light blue)
→ name taken from Greek ʻozeinʼ meaning to smell
late 19th century: studies of solar radiation spectrum revealed a layer of ozone above the tropopause
1930: Sydney Chapman attempted to explain the existence of the stratospheric ozone layer by the so-called ʻChapman cycleʼ
1970s: - budgeting work on stratospheric ozone including catalytic NOx reaction cycles by Paul J. Crutzen
- laboratory discovery of the role of CFCs concerning ozone depletion by Mario Molina and Sherwood Rowland
1984: Detection of ozone hole at the British Antarctic Survey station Halley Bay.
1995: Nobel price in chemistry for Crutzen, Molina and Rowland
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• Tropospheric ozone levels are increasing – Contributes to global warming – Air quality problems
• Stratospheric ozone levels are decreasing – More UV reaches the troposphere, affecting air chemistry
and human & ecosystem health • Both processes are primarily due to human activities.
Solar radiation spectrum: blackbody at 5900 K
Tropospheric ozone (northern mid-latitudes)
1870 1990 0
60
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Pinatubo
In the stratosphere • O2 and O3 absorb UV radiation → temperature increases with altitude → slow mixing → protects troposphere from UV
The ozone layer thus creates the stratosphere, and traps air in the troposphere!
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~tropopause
~3,000 ppb
<100 ppb The absorption of sunlight by ozone is the reason the stratosphere exists at all!
O3 + hν
O2 + hν Only this part reaches
the Earth’s surface (λ > 290 nm).
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Sydney Chapman proposed a chemical mechanism:
O2 + hν O + O (λ < 240 nm)
O + O2 + M O3 + M
O3 + hν O2 + O(1D) (λ < 320 nm)
O(1D) + M O + M
Net: O3 + hν O2 + O
O3 + O 2O2
Ground-state (but still reactive with O2) O(3P)
Lots of energy needed
O3 bonds weak
Loss of O3:
FAST
SLOW
SLOW
Why is there ozone in the stratosphere?
(R1)
(R2)
(R3)
(R4)
Chapman cycle
1) O2 + hν (λ< 240 nm) → 2O 2) O + O2 + M → O3 + M (*) 3) O3 + hν (240 nm<λ< 320 nm) → O + O2 4) O + O3 → 2O2 5) O + O + M → O2 + M (usually unimportant)
(*) Almost all free oxygen atoms participate in this reaction (due to large excess of O2)
Reactions 2-3 interconvert the two forms of odd oxygen (O and O3) and are much faster than reactions 1 and 4 which create and destroy odd oxygen.
The ozone balance is determined mainly by the total sum of odd oxygen (O + O3).
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Steady state solution
O O3 O2 slow
slow
fast
Odd oxygen family Ox = O + O3
Chapman mechanism
Chemical steady-state assumed for species if production and loss rate constant over lifetime Shortest-lived species (O):
similar for [O] between R3 and R2 (& neglecting slow R1 and R4)
R2
R3 R4
R1
O3
O
Observations agree closely with Chapman
[Ox] = [O] + [O3] ≅ [O3]
Steady state approximation: k2⋅[O]⋅[O2] ⋅[M]=k3 ⋅[O3]
Letʼs summarize
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[Ox] = [O] + [O3] ≅ [O3]
• [O3] controlled by slow production and loss by R1 and R4 (NOT fast production and loss of O3 from R2 and R3)
• Effective O3 lifetime ≅ τOx:
• τOx = [Ox]/2k4[O][O3] ≅ 1/ 2k4[O]
(factor of 2 can be derived formally from mass balance)
What have we learned?
upper stratosphere: τOx short enough that steady state can be assumed:
(applying the P.S.S. approximation for [O])
Values of τOx < 1 day in the upper stratosphere
several years in lower stratosphere
τOx
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• k1(z) and k3(z) are photolysis rate constants (J, not reaction rate constants)
k = ∫λ qX(λ)σX (λ)Iλd λ
Iλ = Iλ,∞ e-δ/cosθ
= ∫λ
∞(σO2 [O2] + σO3 [O3])dz
*: numerical solution obtained by starting from top of atmosphere and going downward incrementally
Actinic flux ∝ [O] & [O3]
Solar zenith angle Optical depth
*
CHAPMAN CYCLE provides qualitative agreement with observations
Maximum ∝ Ox production: 2k1[O2]
Lower stratosphere: s.s. not expected because of long τOx ALSO DYNAMICS
τOx
Upper stratosphere: flaw in theory…
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• At high altitudes, lots of UV light ⇒ reaction 1 efficient but total gas molecules concentrations are low ⇒ not much O2 around to produce O atoms.
• At low altitudes, lots of O2, but little UV light ⇒ not enough photolysis events, low O production.
• Optimal O3 production thus occurs at intermediate altitudes. • Chapman functions (plotted on the previous page, see e.g.
Wayne pg 161 -> for some derivations) predict the altitude of the maximum reasonably well, but overestimate the concentration of ozone severely. – Also, horizontal O3 patterns are not correctly predicted.
The natural ozone layer Figure is compilation of available measurements from 1960s
Theory predict maximum O3 production in the tropics
But [O3] is not largest in the tropics
To explain this (and low strat. H2O) Brewer and Dobson suggested a circulation pattern
Region of largest production
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Other comments
O3 maxima occur toward high latitudes in late winter/early spring - the result of the descending branch of the B-D circulation
Virtually no seasonal change in the tropics
More accurate data has led to improvements in our understanding of this simple circulation pattern.
Brewer-Dobson circulation Observation
O3 columns are smallest in tropics despite this being the main stratospheric O3 production region
Explanation
Rising tropospheric air with low ozone
B-D circulation transports O3 from tropics to mid- and high latitudes
Recall that τOx is quite long in the lower stratosphere.
stratosphere
Brewer – Dobson circulation
From the tropics to the mid- & high latitudes
Highest UV in the tropics ⇒ highest ozone production
Circulation stronger in the northern hemisphere
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Chapman got it almost right, but did not account for…
Catalytic cycles for ozone loss: General Idea
O3 + X XO + O2
O + XO X + O2
Net: O3 + O 2 O2
X is a catalyst
The catalyst is neither created nor destroyed…but the rate for the catalytic cycle [odd-oxygen removal in this case] depends on catalyst concentrations
1) Hydrogen oxide (HOx) radicals (HOx = H + OH + HO2)
Initiation: H2O + O(1D) 2OH
Propagation through cycling of HOx radical family (examples):
OH + O3 HO2 + O2 H + O3 HO + O2
HO2 + O OH + O2 OH + O H + O2 Net: O + O3 2O2 Net: O + O3 2O2
Termination (example): OH + HO2 H2O + O2
Source from troposphere
HOx is a catalyst for O3 loss, but not the only one…
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2) Nitrogen oxide (NOx) radicals (NOx = NO + NO2)
• Initiation N2O + O(1D) 2 NO (N2O: “laughing gas”) • Propagation
NO + O3 NO2 + O2 NO + O3 NO2 + O2 NO2 + hν NO + O NO2 + O NO + O2 (O + O2 + M O3 + M) Null cycle – no OX removed Net O3 + O 2O2
• Termination Recycling NO2 + OH + M HNO3 + M HNO3 + hν NO2 + OH
NO2 + O3 NO3 + O2 HNO3 + OH NO3 + H2O NO3 + NO2 + M N2O5 + M NO3 + hν NO2 + O N2O5 + H2O 2HNO3 N2O5 + hν NO2 + NO3
O3 loss rate:
NO3, HNO3, N2O5 are reservoir species.
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What have we learned about NOy?
Production Natural NOy by
N2O + O(1D) - well understood source
Loss Dominant sink is deposition to troposphere. Residence time for air is 1-2 years. Loss rate well constrained
Cycling O3 loss related to NOy/NOx ratio. Under most conditions s.s. between different NOy species is a good approximation
NOx catalytic cycle reconciled Chapman theory with observations … 1995 Nobel Prize
• Initiation: (e.g.) CF2Cl2 + hν CF2Cl + Cl
• Propagation: Cl + O3 ClO + O2 ClO + O Cl + O2 Net: O3 + O 2O2
• Termination: Recycling: Cl + CH4 HCl + CH3 HCl + OH Cl + H2O ClO + NO2 + M ClONO2 + M ClONO2 + hν ClO + NO2 "
O3 loss rate:
Each chlorine atom destroys on the order of 100 000 ozone molecules before it is removed from the stratosphere.
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O + O catalytic cycle (example) OH + O H + O2 H + O2 + M HO2 + M HO2 + O OH + O2 Net: O + O + M O2 + M • Important at high altitudes
where [O]/[O3] is higher
O3 + O3 catalytic cycle (example) OH + O3 HO2 + O2 HO2 + O3 OH + O2 + O2 Net: O3 + O3 O2 + O2 + O2 • Important at low altitudes where
[O]/[O3] is low
Null cycle (example) NO + O3 NO2 + O2 NO2 + hv NO + O Net: O3 + hv O2 + O • No OX loss. Important because
the NOX tied up in null cycle is not removing OX in catalytic cycles.
Holding cycle (example) Cl + CH4 HCl + CH3 OH + HCl H2O + Cl Net: CH4 + OH CH3 + H2O • Does not involve OX directly,
but Cl atoms ”tied up” as HCl are not participating in catalytic cycles.
Holding cycles involve reservoir species.
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Formation of reservoir species
HOCl ClO + HO2 HOCl + O2 ClONO2 ClO + NO2 ClONO2 HO2NO2 HO2 + NO2 + M HO2NO2 + M HNO3 OH + NO2 + M HNO3 + M
Creation of odd oxygen!
OH + CO H + CO2 H + O2 + M HO2 + M HO2 + NO OH + NO2 NO2 + hv NO + O Net: CO + O2 + hv CO2 + O
Mixed null cycles
OH + O3 HO2 + O2 HO2 + NO OH + NO2 NO2 + hv NO + O Net: O3 + hv O2 + O
Cl + O3 ClO + O2 ClO + NO Cl + NO2 NO2 + hv NO + O Net: O3 + hv O2 + O
(Note that NO2 photolysis also competes with the reaction NO2 + O NO + O2 which would lead to a net loss of Ox.)
Wayne, pg 169: Each change in the rate coefficient for the HO2 + NO OH + NO2 reaction has necessitated drastic revision of stratospheric models.
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• Br – BrO catalytic cycle similar to Cl – ClO cycle. • However, Br is almost 60 times as effective as Cl in destroying ozone.
Why? • 1) BrONO2 is much less stable than ClONO2:
– BrONO2 + hv BrO + NO2 Effective also for long λ (visible light) – BrONO2 + hv Br + NO3 No such channel available for Cl
• 2) HBr formation thermodynamically unfavourable (unlike HCl formation) – Br + CH4 HBr + CH3 endothermic reaction
• Mixed bromine – chlorine cycle also strengthens ozone depletion: BrO + ClO Br + Cl + O2 Br + O3 BrO + O2 Cl + O3 ClO + O2
Net: O3 + O3 O2 + O2 + O2 • Fortunately, atmospheric bromine levels are quite low.
– The decision to use chlorine instead of bromine in CFCs is one of the luckiest decisions ever made. If bromine had been used instead, the ozone layer would probably have been gone before anyone noticed.
• Since Cl and Br are efficient at destroying oxygen, one might expect that also F (fluorine) and I (iodine) would participate in catalytic cycles.
• However, neither is very effective: • IO + hv I + O photolysis is very effective; so the IO + O I +
O2 reaction does not have time to occur ⇒ IOx catalytic cycle inefficient. – Also, stratospheric iodine levels are very low.
• Stratospheric fluorine loading is quite high, but HF is very stable and its formation reactions are fast ⇒ fluorine radicals are very rapidly removed from the stratosphere.
• ⇒ F and I contributions to ozone chemistry are negligible.
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• Increasing NOx weakens the ClOx cycle by increasing ClONO2 formation: the net effect of NOx + ClOx is less than the sum of its parts.
• Increasing HOx weakens the NOx cycle by increasing HNO3 formation: OH + NO2 + M HNO3 + M
• Increasing HOx strengthens the ClOx cycle by converting HCl back into active Cl:
OH + HCl Cl + H2O • BrOx and ClOx cycles strengthen each other due to the ClO + BrO
reaction: the net effect is greater than the sum of its parts.
Effect of increasing NOx levels on total ozone depletion (S&P pg 188). In low [NOx] conditions, adding more NOx decreases the total ozone loss due to ClONO2 and HNO3 formation. At higher [NOx] conditions, the NOx cycle starts to dominate and ozone depletion increases linearly with [NOx].
• Natural chlorine and bromine sources mainly HCl, CH3Cl, CH3Br (from e.g. oceans, algae, volcanoes, biomass burning): short lifetimes due to oxidation and deposition ⇒ only small amounts reach the stratosphere.
• The use of CFC (Chlorinated Fluorinated hydroCarbons), a.k.a. freons since the 1930s have significantly increased the concentration of chlorine in the stratosphere. Similarly, NOx concentrations have increased due to N2O emissions.
• CFC:s are not water soluble, and are not oxidized in the troposphere, their only sink is photolysis by UV radiation in the stratosphere:
CFCl3 + hv CFCl2 + Cl • CFCs used as refrigerants, lubricants, solvents, aerosol propellants etc. • In 1973, Jim Lovelock et al. reported that the quantity of CFCs in the
atmosphere was ”equal, within experimental error, to the total amount ever manufactured”.
• CFC lifetimes measured in several decades or centuries. • Bromine – containing CFC analogs (”halons”) used e.g. as fire retardants
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• As a result of increasing stratospheric chlorine and bromine loading, stratospheric ozone concentrations are declining. – ⇒ UVB and UVC radiation levels are increasing. – Possible/probable effects include:
• Increases in skin cancer cases. • Degradation of various tissues in the eyes . • Damage to ecosystems (e.g. ocean plankton). • Damage to crops. • Faster degradation of many building materials (plastic, wood). • Increased D-vitamin production.
– It should be noted that lifestyle changes (sunbathing etc.) have increased the average European & North American human UV radiation exposure more than ozone depletion to date.
• ”Ozone depletion” really refers to two quite different phenomena: – 1) A slow (and more or less steady) decline of mid-latitude ozone
concentrations. – 2) Formation of a ”ozone hole” with extremely low ozone
concentrations over the Antarctic in spring (and similar but weaker phenomenom in the Artic).
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Montreal protocol
CFC production Is banned
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Concentration increases with altitude, does not depend on latitude
Chlorine or dynamics?
Pinatubo
Stabilization of chlorine?
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Greater ozone depletion over polar regions than over the tropics or mid-latitudes, and greater ozone depletion over the antartic than the artic.
British Antarctic Survey, 2005
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DU
Southern hemisphere ozone column seen from TOMS, October
1 Dobson Unit (DU) = 0.01 mm O3 STP = 2.69x1016 molecules cm-2
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Almost all ozone in the lower stratosphere destroyed in the Antarctic spring!
Upper stratosphere ozone levels almost unchanged.
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• Abnormally high ClO concentrations in the polar vortex allow a new catalytic cycle which increases ozone destruction:
ClO + ClO + M ClOOCl + M ClOOCl + hv ClOO + Cl ClOO + M Cl + O2 + M (2x) Cl + O3 ClO + O2 Net: O3 + O3 O2 + O2 + O2 • Also, high ClO concentrations increase the rate of the ClO + BrO
reaction, liberating both Cl and Br. • But why is ClO so high in first place?
• When temperatures are low enough, so-called polar stratospheric cloud (PSC) particles form. – Classified into different types based on chemical composition;
contain H2SO4 and/or HNO3 and/or H2O • At the surface of PSC:s, reservoir species of the chlorine cycle are
efficiently converted back into active chlorine. ClONO2(g) + HCl(s) Cl2(g) + HNO3(s)
Very fast reaction, but happens only at the surface of PSC particles! Other surface reactions include e.g. : HOCl(g) + HCl(s) Cl2(g) + H2O(s)
ClONO2(g) + H2O(s) HNO3(s) + HOCl(g) • Back in the gas phase, Cl2 + hv Cl + Cl HOCl + hv OH + Cl Cl + O3 ClO + O2 (and so on)
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Pola
r stra
tosp
heric
clo
uds
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Spring (SH):
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• Sedimentation of PSC:s removes HNO3, thus lower NOx levels in polar stratosphere in spring/summer
• This decreases the efficiency of the ClO + NO2 ClONO2 reaction, and ozone depletion by Cl/ClO is more effective.
• Hence, the ozone hole persists for quite some time after the PSC:s have all evaporated.
• Type I, formed at 190-195 K – Diameter typically < 1 µm – Ia: solid particles, mostly nitric acid trihydrate (NAT). – Ib: supercooled liquid droplets of HNO3-H2SO4-H2O ternary solution.
• Type II, formed at 188-190 K. – Diameter up to 1 mm! – Composed mostly of water ice.
• All PSCs catalyse the chlorine-freeing reactions (albeit at different rates), but type II PSCs are much more efficient at denitrification.
• Thermodynamics and nucleation kinetics of the HNO3-H2SO4-H2O system at low temperatures (as well as the detailed kinetics of the surface reactions) are quite complicated and still being researched…
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Ia Ib
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• Polar vortex over the artic is much weaker due to meteorology (presence of continents closeby). – More mixing, more
transport of ozone via the B-D circulation.
– Higher temperatures, less formation of PSCs.
• At worst, spring ozone levels over Finland have been ∼35% lower than ”normal”.
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• Average mid-latitude ozone depletion has been about 3-4% per decade, which is more than what gas-phase only models would predict.
• Reactions on stratospheric aerosol surfaces are needed to quantitatively explain observed ozone concentrations and trends.
• The most important mid-latitude heterogeneous reaction is removal of NOx (which decreases ClONO2 formation):
N2O5(g) + H2O(s) → 2 HNO3(s) • Some chlorine activation (analogous to PSC reactions) probably also
occurs. • Mid-latitude stratospheric aerosol consist mainly of H2SO4 + H2O.
Sulfate emissions to the stratosphere (from tropospheric COS + volcanic SO2) thus play an important role in determining stratospheric ozone concetrations.
• Major volcanic eruptions have a strong, but short-lived, ozone-depleting effect (e.g. Pinatubo led to a 25% decrease in mid-latitude ozone in 1991).
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N2O5 hydrolysis also increases HOx-catalyzed ozone loss (why?).
• Bromine – containing molecules are worse than molecules containing only chlorine.
• Hydrogen – containing halocarbons (HCFCs) have much shorter lifetimes (due to tropospheric oxidation) and are thus less harmful.
• All other things being equal, more chlorine/bromine means more ozone depletion (e.f. CFCl3 is worse than CF2Cl2).
• No chlorine or bromine = no problem. – For the ozone layer, that is. But fluorine – hydrogen (HFCs)
compounds may have their own problems; e.g. they may be strong greenhouse gases, and some oxidation products may be very toxic, like CF3COOH.
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• Analogous to global warming potential presented earlier. CFC-11 (CFCl3) is used as the reference compound for which ODP = 1.
• In itʼs simplest form, ODP for a halocarbon X is:
Where CLP is the Chlorine Loading Potential of X, the f:s are efficiency factors, M are molar masses, τ are lifetimes and n is the number of chlorine atoms in one molecule of X.
• The efficiency factors can be used to e.g. account for different photodissociation efficiencies, or the greater ozone-destroying potential of bromine compared to chlorine.
• The efficiency factors, and thus the ODPs, are somewhat dependent on the details of the stratospheric chemistry model used to compute them. They also typically depend somewhat on the concentrations of other reactants.
• The ODP as defined above corresponds to the total ozone destroyed by a halocarbon over its entire lifetime (which can be quite long!); more advanced treatments can be used to compute time- and concentration dependent ODPs.
Halocarbon τ, years ODP Contribution to ozone loss 1998*
CFCl3 53 1.0 21.0
CF2Cl2 100 0.82 31.2
CF2ClCFCl2 83 0.90 8.1
CHFCl2 11.8 0.04 2.3
CF3Br 69 12 7.3
CF2BrCl 36 5.1 7.3
CH3CCl3 4.8 0.12 2.5
CCl4 46 1.2 5.9
CH3Br 0.7 0.34 14.3
*: relative values, calculated only for the compounds in this table
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• Legally binding international treaty from 1987 to phase out CFC emissions. – Additional targets set in 1990, 1992, 1997, 1999.
• The current agreement requires the phase-out of all production and use of ozone-depleting chemical (CFCs, halons and HCFCs). – Some exceptions for medical and research use.
• Extra time for developing countries. • Widely considered one of the most succesful international
(environmental) treaties ever.
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• Also tropical and mid-latitude ozone concentrations would likely have dropped catastrophically at some point. (P. A. Neuman et al., ACPD 8, 20565, 2009, "Business-as-usual" simulation)
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• 1930s-1950s: Chapman mechanism – Explains existence of ozone layer, qualitative description of vertical
behavior. With Brewer-Dobson circulation, also horizontal patterns qualitatively described.
• 1960s-1970s: Gas-phase catalytic cycles – Explains (semi-quantitatively) mid-latitude ozone concentrations
and long-term ozone depletion.
• 1980s-1990s: Heterogeneous chemistry – Explains polar ozone holes, quantitative explanation for mid-latitude
ozone concentrations.
