Transcript The Chemistry of the Stratosphere — EPS 200  (Ǻ) -1800 -3400 Penetration of sunlight ozone “cutoff” of solar ultraviolet light The stratospheric ozone layer prevents.

The Chemistry of the Stratosphere — EPS 200
20061026
 (Ǻ) <1200 1200
-1800
1800
-3400
Penetration of sunlight
ozone “cutoff” of
solar ultraviolet
light
The stratospheric ozone layer prevents light with < 3100 (Ǻ) from reaching
the earth’s surface. This provided the clue needed to deduce the presence of
ozone in the atmosphere, and it protects living organisms from the damaging
effects of UV light on DNA and other proteins.
http://www.ldeo.columbia.edu/dees/V1003/lectures/ozone/
Chapman mechanism for stratospheric
Ozone
•
The chemical processes responsible for production
and loss of O3 in an idealized O2-N2 atmosphere:
The Chapman formulation (1930):
hv + O2  O + O (13.1) ;
J 13.1
O + O2 +M  O3 + M (13.2); k 13.2
hv + O3  O + O2 (13.3);
J 13.3
O + O3  O2 + O2 (13.4);
k 13.4
O + O + M  O2 + M (important only in the mesosphere and up)
O3 + O3  2 O2
(does not proceed unless catalyzed)
Introduce concept of odd oxygen, O + O3.
hv + O2  O + O (13.1) ;
J 13.1
O + O2 +M  O3 + M (13.2); k 13.2
hv + O3  O + O2 (13.3);
J 13.3
O + O3  O2 + O2 (13.4);
k 13.4
(13.1) is the only source of odd O
(13.4) is the only sink of odd O.

Reactions (13.2), (13.3) recycle odd O: they
transform one type of odd O into another.
Inner and outer cycles for O3, oddOxygen
McElroy 2002
The concept of steady state
}
O3
(odd –O)
Steady Stat e O
dO
dt
 P r oduction- Loss
 2J13.1[O2 ]  J13.3 [O3 ]
 k13.2 [O][O2 ][M ] k 13.4[O][O3 ]
dO3
dt
 k13.2 [O][O2 ][M ]  J13.3 [O3 ] k 13.4[O][O3 ]
Since O has a very short lifetime, top equation
implies
J13.3
[O]

[O3 ] k13.2 [O2 ][M ]
d (O3)/dt  d (O + O3)/dt = 2 J13.1 [O2] – 2 k 13.4 [O][O3]
(odd oxygen equation)
Profile of O3 with height
(Chapman Mechanism)
k13.2 J13.1 1/ 2
Steady State [O3 ]  (
) f O2 [ M ]3 / 2
J13.3k13.4
J13.1 J13.3 1/ 2
[O]  (
) [ M ]1/ 2
k13.2 k13.4
Mole fraction [O3]/[M] decreases with height as M ½
[O]/[M] increases with height as M– 3/2
From Eq. 13.38 and 13.39, McElroy 2002
Comparison between calculated [O3] and observed
[O3] using this oxygen-only approach reveals very
large differences
What factors control ozone production and loss to
create such a discrepancy?
What other molecules or processes could be
involved?
Distribution of stratospheric ozone
[units: ppm (10-6 moles/mole) ]
Temperature (July)
Term
Symbols
Specify the energy state of reactant by
defining the
1
 (2 pz ) 
cos r e r 2
wavefunction
4 2
Principal quantum number n
Angular momentum QN
l
l = 0  s
n = 1, 2, 3, …
l = 1  p
l = 0, 1, 2, … n 1
l = 2  d
n and l designate a particular
orbital; the orientation of the
l = 3  f
orbital is given by the
magnetic quantum number, m
m = l , l 1, … 0, …  l +1,  l
Spin quantum number, ms = 1/2
Coulomb-Coulomb repulsion
dictates maximum distance
between electrons
What is the Ground State?
Term symbol
Total orbital angular momentum L   li
2 S 1
LJ
i
Total spin angular momentum S   si
i
Total angular momentum J  L  S
Steps to determine lowest energy (ground) state
Hund’s Rules:
• Ground state always has maximum spin multiplicity
• When two states have the same spin multiplicity, state with higher L is usually more stable
• For given S and L, minimum J value is most stable when subshell is less than half full, maximum
J more stable when subshell is more than half full
• When subshell is more than half full, disregard full subshells
Consider oxygen:
m =
ms =


1s
2s
0
1/2
0
1/2



2p
-1
1/2
0
1/2
+1
1/2
M L  ml 1  ml 2  0  1  1
L 1
S  1 2 1 2  1
3P
2,1,0
Potential energy curves of O2
…and
E
O3
O (1D) + O2(1)
O (3P) + O2(3S)
r( O + O2)
E
…and N2O
O (1D) + N2(1S)
O (3P) + N2(1S)
r( O + N2)

J    q e

  ndz
z
J1
O2  h 
O  O
k1
O  O2  M 
 O3  M
J2
O3  h 
 O  O2
k2
O3  O 
 2O2
d [O]
dt
[O] 
 2 J1[O2 ]  J 2 [O3 ]  k1[ M ][O2 ][O]  k2 [O3 ][O]
J1[O2 ]  J 2 [O3 ]
k1[ M ][O2 ]  k2 [O3 ]
d [O3 ]
Define new species
"Odd oxygen" = Ox  O  O3
d [O x ]
dt
dt
 k1[ M ][O2 ][O]  J 2 [O3 ]  k2 [O][O3 ]

d [O3 ] d [O]

 2 J 1[O2 ]  2k2 [O][O3 ]
dt
dt
Rate of Ox productionrate of breaking the O2 bondis just 2J1[O2]
Rate of O3 removal rate of reforming the O2 bondis 2k2[O][O3]
This event occurs at ~ 106 cm-3sec-1 in the stratosphere
But
occurs at 109cm-3sec-1 in the stratosphere, so
J 2 [O3 ]  J1[O2 ]
J [O ]
[O]  2 3
k1[ M ][O2 ]
k1[ M ][O2 ]  k2 [O3 ]
Catalyst
Chapman got the process almost right…
O3 + X  XO + O2
O + XO  X + O2
Net:
O 3 + O  2 O2
X is a catalyst
The concept of a rate limiting step: suppose XO can also react with Y,
XO + Y  XY + O
Then odd-o is not removed. The rate for odd-o loss due to catalyst X is
rate for odd-o loss = 2 kxo+o [O][XO]
xo+o is the "rate limiting" step in the catalysis
A more detailed look at the stratospheric
photochemistry of O, O2, and O3
O 2  h  O( P)  O( P)
J1
3
3
O  O2  M 
 O3  M
k1
O3  h  O( D)  O 2 (  )
J2a
1
1
J2b

O( 3 P)  O 2 ( 3 )
k2
O  O3 
 O2  O2
O(1D)
O(3P)
+M
kq
+M
d O  1 D 
dt
 J 2 a [O3 ]  kQ [ M ][O  1D ]  0
10@15 km ;1 @ 0 km
J [O ]
O  1D   2 a 3 ~ 100@ 25 km

 k [M ]
Q
1000@ 35 km
(cm-3)
HOx radicals as catalysts
O3 + hv  O (1D) + O2
Produces excited O (1D)
O(1D) +M  O(3P) +M
Quenches O (1D)
O(1D) + H2O  2 OH
Oxidizes water !
Free radicals have odd # of electrons)
Created (or removed) 2 at-a-time
OH  HO2
OH+HO2  H2O + O2
Regenerates H2O;
removes 2 radicals
OH and HO2 are HOx free radicals
Examples of catalytic cycles involving HOx radicals
The catalyst is neither created nor destroyed…but the rate for the
catalytic cycle [odd-o removal in this case] depends on catalyst
concentrations.
O3 + OH  HO2 + O2
HO2 + O  OH + O2
O3 + O  2 O2
O3 + OH  HO2 + O2
HO2 + O3  OH + 2O2
O3 + O3  3 O2
O3
O
H2O
OH
O (1D)
+CH4,
+HNO3
H2O
HO2
O3
O
H2O
HOx radicals
Inner and outer cycles of radical chain mechanism
CH4  O( 1 D)  CH3  OH
One of the most important is methane
Another source of HOx
HOx  OH  HO2  H
d[HOx ]
dt
 k [CH 4 ][O( 1 D)]  2kOH HO2 [HO2 ][OH]
OH  O  H + O2
H  O3  OH  O 2
net O  O3  O 2  O 2
d [H]
dt
 kO OH [OH][O]  k H O2 [ M ][O2 ][H]  k H O3 [H][O3 ]
[H] 
kO OH [OH][O]
k H O2 [ M ][O2 ]  kO OH [OH][O]

kO OH [OH][O]
k H O2 [ M ][O2 ]
~ 104 cm 3  k H O2 [ M ]  kO OH [OH][O]
 [OH],[HO2 ]
OH  O3  HO 2 + O 2
RLS HO 2  O3  OH  2O 2
(1)
net O3  O3  3O3
OH  O3  HO 2 + O 2
RLS HO 2  O  OH  O 2
(2)
net O  O3  O2  O 2
OH  O3  HO2 + O2
HO2  NO  OH  NO 2
NO2  h  NO  O
O + O2  M  O3  M
net
NULL
Cycle 1 and 2 are catalytic with respect to removal of Ox by reactive nitrogen, but what is the
quantitative contribution to d[Ox]/dt?
IDENTIFICATION OF RATE LIMITING STEP:
1. Only when both reactions take place in sequence is O2 bond reformed
2. So rate of net reaction is controlled by: Fast step? Slow step?
3. Thus
d [Ox ]
dt
 2kO3
Lower stratosphere
HO2
[HO2 ][O3 ]  kHO2 O [HO2 ][O]
Upper stratosphere
(3)
NO x  NO  NO2  NO3
NO y  NO x  HONO2  N 2O5
But we remain focused on the catalytic destruction of O3
NO  O3  NO 2  O 2
NO  O3  NO 2  O 2
NO 2  O  NO  O 2
NO 2  h  NO  O
O3  O  O 2  O 2
d [O x ]
dt
 2kO NO2 [NO 2 ][O]
net NULL
Other cycles!
O3
N2O
NO
O
(1D)
OH h
NO2
NO3, M, aerosol
hv
O
OH
NOx radicals
Odd Nitrogen
acid
HNO3
Cl  O3  ClO  O 2
ClO  NO  Cl  NO 2
NO2  h  NO  O
Cl  O 3  ClO  O 2
RLS ClO  O  Cl  O 2
net O3  O  O 2  O 2
Cl  O 3  ClO  O 2
RLS ClO  HO 2  HOCl  O 2
HOCl  h  OH  Cl
OH  O3  HO 2  O 2
net 2O3  3O 2
d [O3 ]
dt
 2kClO O [O][ClO]  kClO HO2 [ClO][HO2 ]
At temperatures below 210 K
O  O 2  M  O3  M
net
NULL
ClONO2  h  Cl  NO3
NO3  h  NO  O 2
Cl  O3  ClO  O 2
NO  O3  NO2  O2
ClO  NO 2  M  ClONO 2  M
net O3  O3  3O2
ClO  ClO  ClOOCl
ClOOCl  h  Cl  ClOO
ClOO  M  Cl  O 2  M
Cl  O3  ClO  O 2
net 2O3  3O2
h, (OH)
CFCs,
MeCCl3,
etc
O3
{Cl + CH4}
Cl
ClO
O
inorganic
Cl
HCl
OH
Bromine Chemistry
Catalytic Cycles
Br  O3  BrO  O 2
RLS BrO  HO 2  HOBr  O 2
HOBr  h  OH  Br
OH  O3  HO 2  O 2
net 2O3  3O 2
RLS ClO  BrO  Cl  Br  O 2
Result
BrO
Bry
BrONO 2
~ 0.5
Bry
~ 0.2
HOBr
Bry
HBr
Cl  O3  ClO  O 2
Br  O3  BrO 2  O 2
net 2O3  3O 2
Br  O3  BrO  O 2
RLS
BrO  O 2  Br  O 2
net O3  O3  O 2  O 2
BrO  NO 2  BrONO 2  M
BrONO 2  h  Br  NO 3
NO3  h  NO  O 2
NO  O3  NO 2  O 2
Br  O3  BrO 2  O 2
net 2O3  3O 2
Bry
~ 0.2
~ 0.1
Iodine Chemistry
Virtually all Iodine s as IO
and I because of the rapid
photolysis of HI
IO  ClO  I  Cl  O 2
IO  BrO  I  Br  O 2
I  O3  IO  O 2
I  O3  IO  O 2
Cl  O3  ClO  O 2
Br  O3  BrO2  O 2
net O3  O 3  O 2  O 2  O 2
net 2O3  3O 2
Interaction between radical cycles
NOx radicals
NOy
Interaction between radical cycles: HOx
In the current atmosphere, NOx radicals are more important
as inhibitors of ozone loss by ClO than as radical catalysts
on their own!
What does this distill down to
with respect to ozone loss?
d [O x ]
dt
 2kO O3 [O][O3 ]  k NO2 O [NO2 ][O]
 2k HO2
O3
[HO2 ][O3 ]  2kHO2 O [HO2 ][O]
 2kClO O [ClO][HO2 ]  2kClO HO2 [ClO][HO2 ]  2kClO ClO [ClO]2
 2k BrO HO2 [BrO][HO2 ]  2kBrO ClO [BrO][ClO]
What is the evidence?
1. Each of these rate limiting free radicals have been observed in the
stratosphere: O, OH, HO2, ClO, BrO, NO2
2. With these catalytic cycles, and the observed radical concentrations,
[O3]calculated and [O3]observed converge
3. But the most important issue is prediction; how will the ozone loss rate
respond to changes in
– Chemical boundary conditions
– T, H2O in response to forcing by CO2, etc.?
O3 column amounts observed with a Dobson spectrometer
(1 Dobson unit (DU) = 2.5 x 1016 molecules cm-2)
satellite-
agung
El Chichon
Pinatubo
O3 over
Arosa
1980
1990
backscatter from
stratospheric
aerosols
Heterogeneous chemistry in the midlatitude stratosphere
The principal reaction affecting NOx converts free radicals to HNO3:
NO2 + O3  NO3 + O2
NO3 + NO2 + M  N2O5 .
The first (rather slow) reaction occurs day or night, but in the day
NO3 is almost instantly photolyzed back to NO2. Therefore
production of N2O5 (nitric acid anhydride) is confined to the night. If
an aerosol surface is available,
N2O5 + H2O(aerosol surface)  2 HNO3 .
The role of heterogeneous chemical reactions
A chemical reaction is said to be "heterogeneous" when the reactants (or
the products) are in different phases. For example, when one reactant is
on the surface of a particle, and the other is in the gas phase.
NO2 + NO3 + M -> N2O5 (homogeneous—gas phase)
N2O5 is the anhydride of nitric acid: N2O5 + H2O (liq) -> 2 HNO3
The hydrolysis of N2O5 takes place efficiently and rapidly on the surface
of particles, the first important heterogeneous reaction recognized in the
stratosphere.
The hydrolysis of N2O5 reduces NOx radical concentrations. Since
currently NOx is most important as an inhibitor of ozone loss by Cl
radical cycles, adding more aerosols (volcanoes) leads to more loss of
O3 and thus lower ozone in the atmosphere.
Heterogeneous chemistry in the midlatitude stratosphere
(continued)
A number of important reactions that affect halogens were discovered
later:
ClNO3 + H2O(aerosol)  HOCl + HNO3
BrNO3 + H2O(aerosol)  HOBr + HNO3 .
These reactions convert relatively stable species (ClNO3, BrNO3) to
molecules that rapidly photolyze to radicals, HOCl + h  HO + Cl.
Two stable species are converted to one readily photolyzed by
HCl + ClNO3aerosol
 Cl2 + HNO3 .
Note that all of the reactions promoted by aerosols tend to speed up
catalytic losses due to the halogen (Cl, Br) radical chain reactions.
•NOx radicals remove odd-o through their own catalytic cycle
•NOx radicals inhibit catalysis by ClO—the dominant effect
•Hydrolysis of N2O5 on the surfaces of aerosol particles is an
important sink for NOx radicals
•When a major volcanic eruption occurs, a great mass of
aerosol particles enters the stratosphere, gradually decaying
over ~ 2 years.
•Ozone levels decline after a volcanic eruption, because the
volcanic aerosols keep NOx low and allow ClO to be a more
efficient catalyst for removal of ozone!
The Antarctic Ozone "Hole" phenomenon
Total ozone observed over Halley Bay, Antarctica, by the British
Antarctic Survey using a Dobson spectrometer (Farman et al., 1985).
Ozone Loss in Northern Hemisphere
ozone sonde
data
(Hofmann)
Scientists were at first completely surprised that dramatic ozone losses
appeared over Antarctica just after polar night, when chemical reactions
were thought to be negligibly slow. Soon it was realized that the very
cold temperatures over Antarctica in polar night set in motion a bizarre
series of chemical processes leading to virtually complete removal of
ozone from a deep segment of the stratosphere. The steps are as follows:
•1. Temperatures get cold enough to condense solid particles containing
nitric acid and water--about 195 K (-78 C (!)) for HNO3·3H2O
Polar stratospheric clouds (PSCs).
•2. Reactions involving HCl and another non-reactive chlorine species,
ClNO3 occur on the surfaces of these particles, releasing reactive Cl
atoms and converting NOx to HNO3.
•3. These particles then sediment out of the atmosphere (like a very fine
snow of nitric acid). This step removes NOy: otherwise when the sun
returned in the spring, Cl radicals would be converted back into nonreactive chlorine species, especially ClNO3.
The Strange Chemistry of the
Antarctic Ozone “Hole”
HCl + ClNO3
cold
particle
surface
Cl2 + HNO3 (solid phase)
Cl2 + hv  2Cl
sedimentation
Cl + O3  ClO + O2
(cold)
ClO + ClO  (ClO)2
“dimer”
(ClO)2 + hv  Cl + ClOO  2Cl + O2
springtime
2O3 + hv  3O2
ER-2
"Denitrification"
ER-2 flight track at
20 km
2 Sep
'87
ClO
O3
1.8
ppm
O3
ClO
1200
ppt
16 Sep '87
ClO
O3
1.8
ppm
Hypothesis 1: The observed long-term trend in mid-latitude Northern Hemisphere ozone in the lower
stratosphere over the last two decades is the result of increased catalytic loss of ozone in the lower
stratosphere by the rate limiting halogen radicals ClO, BrO, and IO. The observed variability of ozone in
the lower stratosphere is the result of factors that modulate those halogen radicals, specifically water
vapor, aerosol loading and temperature.
This hypothesis addresses a central tenet in the latest
WMO report: Scientific Assessment of Ozone Depletion:
2002. In the Scientific Summary for the key Chapter 4,
under, “Attribution of Past Changes in Ozone,” on page
4.2, it is stated:
“The vertical, latitudinal, and seasonal characteristics of
changes in mid-latitude ozone are broadly consistent
with the understanding that halogens are the primary
cause of these changes, in line with similar conclusions
from the 1998 Assessment.”
Strange Chemistry of the Antarctic Ozone “Hole”:
4. When the sun returns to the polar region in September
(springtime), the temperatures remain cold enough so that ClO combines
with itself to make a rapidly-photolyzed dimer, (ClO)2.
HCl + ClNO3
cold
particle
surface
Cl2 + HNO3 (solid phase)
Cl2 + hv  2Cl
Cl + O3  ClO + O2
(cold)
ClO + ClO  (ClO)2
“dimer”
(ClO)2 + hv  Cl + ClOO  2Cl + O2
2O3 + hv  3O2 (net reaction)
2 Sep
'87
ClO
2.0
ppm
O3
1.8
ppm
1.8
O3 ppm
ClO
1200
ppt
16 Sep '87
ClO
O3
1.0
ppm
The NASA ER-2 stratospheric research aircraft in Kiruna, January 2000.
(dip)
Arctic ozone and ClO from Kiruna
(landing)
Polar stratospheric clouds over Kiruna, Sweden,
during the "SOLVE" scientific expedition in 2000.
Snow sculpture in Kiruna, world championship competition, 2000.
Lessons to be learned:
•We don't know enough to predict when or where very dramatic changes
may occur in the global environment due to human activities. Particles
are present in the stratosphere outside the polar regions, but their
surfaces happen to be unreactive towards HCl and ClNO3. What would
have happened if these particles were reactive?
•Dramatic changes may occur in the global environment due to humans
and be undetected until they are essentially irreversible. The ozone
"hole" first appeared in about 1975; it wasn't detected until 1985, and
CFCs weren't phased out until 1996. The "hole" continue beyond 2075.
•We don't know enough to predict the impact of even rather dramatic
changes in the global environment arising from human activities. There
are extremely high rates of skin cancer in Australia, declining fisheries,
and possible effects on vegetation, but it is not possible to "prove" that
changes in global ozone have contributed significantly to any of these.
•Consider these lessons in the context of climate change…
Montreal Protocol
In 1978 the US Congress passed a bill barring the use of CFCs in aerosol
spray cans, a major use deemed inessential. A few Scandinavian
countries also did this. Only theories, and a few rough measurements of
ClO, related CFCs to possible loss of stratospheric O3.
As better data accumulated, the US led an international diplomatic effort
to convince the countries of the world to cease use of these compounds.
In 1987, at the urging of Secretary of State George Schultz, President
Reagan signed the Montreal Protocol mandating the 50% reduction in
CFC use in developed countries by the year 2000. The treaty required
triennial scientific reviews with the authority to modify the targets.
In 1990, after the ozone hole was conclusively shown to result from
global anthropogenic halogen pollution, the mandate was changed to
complete phase-out of CFCs, CH3CCl3 in developed countries by 1996.
Phase-out of CH3Br was added later.
Montreal
Protocol
European
acceleration
European
acceleration
Austin and Wilson, 2006
Antarctic Ozone "hole"
(minimum value) for
different scenarios
= Cly + 50 Bry
Austin and Wilson, 2006