Jacobson M.Z. Atmospheric Pollution

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Ozone absorbs wavelengths shorter than 0.35

m (strongly below 0.31 m and

weakly at 0.31 to 0.35 m). Ozone also absorbs weakly at 0.45 to 0.75 m.

Stratospheric ozone allows little UV-C, some UV-B, and most UV-A radiation to reach

the troposphere. Ozone in the background troposphere absorbs some of the UV-B and

UV-A not absorbed in the stratosphere. In polluted air, additional absorbers of UV-B

radiation include nitrated aromatic gases and aerosol particle components,

such as

black carbon (BC), nitrated aromatics, polycyclic aromatic hydrocarbons (PAHs), and

soil-dust components (Section 7.1.3.1).

The major UV-A-absorbing gas is nitrogen dioxide [NO

(g)]. Its mixing ratio in

clean air is too small to affect UV-A. In polluted air

, its mixing ratio is high only in the

morning, when UV-A intensity is low. Other UV-A absorbers in polluted air include

the same aerosol particle components that absorb UV-B radiation.

Gas- and aerosol particle absorption is not the only mechanism that reduces inci-

dent downward UV radiation. Gas and aerosol particle backscattering, ground reﬂection,

and cloud reﬂection return some incident UV radiation to space as well.

11.3. CHEMISTRY OF THE NATURAL OZONE LAYER

The chemistry of the natural ozone layer involves primarily oxygen-containing com-

pounds, but the shape of the stratospheric ozone vertical proﬁle is affected by nitrogen-

and hydrogen-containing compounds as well. Next, the chemistry of the natural ozone

layer is discussed.

11.3.1. The Chapman Cycle

The photochemistry of the natural stratosphere is similar to that of the free tro-

posphere, except that stratospheric ozone is produced after photolysis of molecular

oxygen, whereas tropospheric ozone is produced after photolysis of nitrogen dioxide.

Next, reactions naturally producing and destroying stratospheric ozone are described.

In the stratosphere, far-UV wavelengths (shorter than 0.245 m) break down

molecular oxygen by

(g)  h •O

•

(

D)(g)  •O

•

(g)   175 nm

Molecular Excited Ground-

oxygen atomic state atomic

(11.1)

oxygen oxygen

278 ATMOSPHERIC POLLUTION: HISTORY, SCIENCE, AND REGULATION

Table 11.1. Summary of Major Absorbers of UV Radiation

Far-UV 0.01–0.25 N

(g) Thermosphere, mesosphere

(g) Thermosphere, mesosphere,

stratosphere

Near-UV

UV-C 0.25–0.29 O

(g) Stratosphere

UV-B 0.29–0.32 O

(g) Stratosphere, troposphere

Particulate components Polluted troposphere

UV-A 0.32–0.38 NO

(g) Polluted troposphere

Particulate components Polluted troposphere

Name of Spectrum Wavelengths (m) Dominant Absorbers Location of Absorption

(g)  h •O

•

(g)  •O

•

(g) 175 <   245 nm

Molecular Ground-

oxygen state atomic

(11.2)

oxygen

The ﬁrst reaction is important only at the top of the stratosphere because wavelengths

shorter than 0.175 m do not penetrate lower. Neither reaction is important in the tropo-

sphere. Excited atomic oxygen from Reaction 11.1 rapidly converts to the ground state by

•O

•

(

D)(g)

•O

•

(g)

Excited Ground-

atomic state atomic

(11.3)

oxygen oxygen

Ozone then forms by

This reaction also occurs in the troposphere, where the O(g) in that case originates

from NO

(g) photolysis, not from O

(g) photolysis. Ozone is destroyed naturally in the

stratosphere and troposphere by

(g)  h O

(g)  •O

•

(

D)(g)   310 nm

Ozone Molecular Excited

oxygen atomic

(11.5)

oxygen

(g)  h O

(g)  •O

•

(g)   310 nm

Ozone Molecular Ground-

oxygen state atomic

(11.6)

oxygen

Stratospheric ozone is also destroyed by

•O

•

(g)  O

(g) 2O

(g)

Ground- Ozone Molecular (11.7)

state atomic oxygen

oxygen

In 1930, English physicist Sidney Chapman (1888–1970; Fig. 11.6) suggested

that ozone in the stratosphere must be produced from UV photolysis of molecular

oxygen. He further postulated that Reactions 11.2, 11.4, 11.6, 11.7, and the reaction

•O

•

(g)  •O

•

(g)

Ground- Molecular

state atomic oxygen

(11.8)

oxygen

GLOBAL STRATOSPHERIC OZONE REDUCTION 279

•O

•

(g)  O

(g)

Ground- Molecular Ozone

state atomic oxygen

(11.4)

oxygen

describe the natural formation and destruction of ozone in the stratosphere (Chapman,

1930). These reactions make up the Chapman cycle, and they simulate the process

fairly well. Some Chapman reactions are more important than are others. Reactions

11.2, 11.4, and 11.6 affect ozone the most. The non-

Chapman reaction, 11.5, is also important.

Some of the Chapman cycle reactions can be used

to explain why the altitudes of peak ozone concentra-

tion and mixing ratio occur where they do in Fig. 11.4.

Oxygen density, like air density, decreases exponen-

tially with increasing altitude. UV intensity decreases

with decreasing altitude. Peak ozone densities occur

where sufﬁcient radiation encounters sufﬁcient oxygen

density, which is near 25 to 32 km (Fig. 11.4). At high-

er altitudes, the oxygen density is too low for its

photolysis by Reactions 11.1 and 11.2 to produce peak

ozone densities; at lower altitudes, the radiation is not

intense enough for oxygen photolysis to produce peak

ozone densities.

11.3.2. Effects of Nitrogen on the Natural

Ozone Layer

Oxides of nitrogen [NO(g) and NO

(g)] naturally

destroy ozone, primarily in the upper stratosphere,

helping shape the vertical proﬁle of the ozone layer. In the troposphere, the major

sources of nitric oxide are surface emissions and lightning. The major source of NO(g)

in the stratosphere is transport from the troposphere and the breakdown of nitrous

oxide [N

O(g)] (laughing gas), a colorless gas emitted during denitriﬁcation by anaer-

obic bacteria in soils (Section 2.3.5). It is also emitted by bacteria in fertilizers,

sewage, and the oceans and during biomass burning, automobile combustion, aircraft

combustion, nylon manufacturing, and the use of spray cans. In the troposphere,

O(g) is lost by transport to the stratosphere, deposition to the surface, and chemical

reaction. Because its loss rate from the troposphere is slow, nitrous oxide is long lived

and well diluted in the troposphere, with an average mixing ratio of about 0.31 ppmv.

The mixing ratio of N

O(g) is relatively constant up to about 15 to 20 km, but decreases

above that as a result of photolysis. Throughout the atmosphere, N

O(g) produces

nitric oxide by

O(g)  •O

•

(

D)(g) N

•

O(g)  N

•

O(g)

Nitrous Excited Nitric oxide

oxide atomic

(11.9)

oxygen

Nitric oxide naturally reduces ozone in the upper stratosphere by

280 ATMOSPHERIC POLLUTION: HISTORY, SCIENCE, AND REGULATION

Figure 11.6. Sidney Chapman (1888–1970).

•

O(g)  O

(g) N

•

(g)  O

(g)

Nitric Ozone Nitrogen Molecular

oxide dioxide oxygen

(11.10)

The result of this sequence is that one molecule of ozone is destro

yed, but neither

NO(g) nor NO

(g) is lost. This sequence is called a catalytic ozone destruction cycle

because the species causing the O

(g) loss, NO(g), is recycled. This particular cycle is

the NO

(g) catalytic ozone destruction cycle, where NO

(g)  NO(g)  NO

(g), and

NO(g) is the catalyst. The number of times the cycle is executed before NO

(g) is

removed from the cycle by reaction with another gas is the chain length. In the upper

stratosphere, the chain length of this cycle is about 10

(Lary, 1997). Thus, 10

mole-

cules of O

(g) are destroyed before one NO

(g) molecule is removed from the cycle. In

the lower stratosphere, the chain length decreases to near 10. When NO

(g) is removed

from this cycle, its major loss processes are the formation of nitric acid and peroxyni-

tric acid by the reactions

•

(g)  O

•

H(g)

HNO

(g)

Nitrogen Hydroxyl Nitric (11.13)

dioxide radical acid

•

(g)  N

•

(g)

Hydroperoxy Nitrogen Peroxynitric (11.14)

radical dioxide acid

Nitric acid and peroxynitric acid photolyze back to the reactants that formed them,

but such processes are slow. Peroxynitric acid also decomposes thermally, but

thermal decomposition is slow in the stratosphere because temperatures are low

there.

The natural NO

(g) catalytic cycle erodes the ozone layer above ozone’s peak alti-

tude shown in Fig. 11.4. Although the NO

(g) catalytic cycle is largely natural, an

unnatural source of stratospheric NO

(g) and, therefore, ozone destruction, is stratos-

pheric aircraft emission of NO(g) and N

O(g). In the 1970s and 1980s, scientists were

concerned that the introduction of a ﬂeet of supersonic transport (SST) jets into the

stratosphere would enhance N

O(g) emissions sufﬁciently to damage the ozone layer

by Reactions 11.9 through 11.12. This concern disappeared because the plan to intro-

duce a ﬂeet of stratospheric jets never materialized.

11.3.3. Effects of Hydrogen on the Natural Ozone Layer

Hydrogen-containing compounds, particularly the hydroxyl radical [OH(g)] and the

hydroperoxy radical [HO

(g)], are responsible for shaping the ozone proﬁle in the

lower stratosphere. The hydroxyl radical is produced in the stratosphere by one of sev-

eral reactions

GLOBAL STRATOSPHERIC OZONE REDUCTION 281

•

(g)  •O

•

(g) N

•

O(g)  O

(g)

Nitrogen Ground- Nitric Molecular

dioxide state atomic oxide oxygen

oxygen

(11.11)

•O

•

(g)  O

(g) 2O

(g)

Ground- Ozone Molecular

state atomic oxygen

oxygen

(net process) (11.12)

(11.15)

The hydroxyl radical participates in an HO

(g) catalytic ozone destruction cycle,

where HO

(g)  OH(g)  HO

(g). HO

(g) catalytic cycles are important in the lower

stratosphere. The most effective HO

(g) cycle, which has a chain length in the lower

stratosphere of 1 to 40 (Lary, 1997), is

OH(g)

OH(g) +

(g)

Hydroxyl

radical

Excited

atomic

oxygen

D)(g) +

O(g)

(g)

Methane

Molecular

hydrogen

Water

vapor

Hydroxyl

radical

Methyl

radical

Atomic

hydrogen

(g) species can be removed temporarily from catalytic cycles by Reactions 11.13

and 11.14 and by the reaction

•

(g)  O

•

H(g) H

O(g)  O

(g)

Hydroperoxy Hydroxyl Water Molecular (11.19)

radical radical vapor oxygen

This mechanism is particularly efﬁcient at removing HO

(g) from catalytic cycles

because it removes two HO

(g) molecules at a time.

11.3.4. Effects of Carbon on the Natural Ozone Layer

Carbon monoxide and methane produce ozone by the reaction mechanisms shown in

Sections 4.2.4 and 4.2.5, respectively. The contributions of CO(g) and CH

(g) to ozone

production are small in the stratosphere. A by-product of methane oxidation in the

stratosphere is water vapor, produced by

(g)  O

•

H(g) C

•

(g)  H

O(g)

Methane Hydroxyl Methyl Water (11.20)

radical radical vapor

282 ATMOSPHERIC POLLUTION: HISTORY, SCIENCE, AND REGULATION

•

H(g)  O

(g) HO

•

(g)  O

(g)

Hydroxyl Ozone Hydroperoxy Molecular

radical radical oxygen

(11.16)

•

(g)  O

(g) O

•

H(g)  2O

(g)

Hydroperoxy Ozone Hydroxyl Molecular

radical radical oxygen

(11.17)

(g) 3O

(g)

Ozone Molecular

oxygen

(net process) (11.18)

Because water vapor mixing ratios in the stratosphere are low and transport of water

vapor from the troposphere to stratosphere is slow, this reaction is a relatively impor-

tant source of water vapor in the stratosphere.

11.4. RECENT CHANGES TO THE OZONE LAYER

Changes in stratospheric ozone since the early 1970s can be divided into global stratos-

pheric changes, Antarctic stratospheric changes, and Arctic stratospheric changes.

11.4.1. Changes on a Global Scale

Figure 11.7 shows that between 1979 and 2000, the global stratospheric ozone column

abundance decreased by approximately 3.5 percent (from 304.0 to 293.4 DU). Unusual

decreases in global ozone occurred following the El Chichón (Mexico) volcanic erup-

tion in April 1982, and the Mount Pinatubo (Philippines) eruption in June 1991 (Fig.

11.8). These eruption injected particles into the stratosphere. On the surfaces of these

particles, chemical reactions involving chlorine took place that contributed to ozone

loss. Over time, however, the concentration of these particles decreased, and the global

ozone layer partially recovered. Because volcanic particles were responsible for only

temporarily ozone losses, the net loss of ozone over the globe from 1979 to 2000 was

still about 3.5 percent. The decrease between 60S and 60N was 2.5 percent (298.08 to

290.68 DU), that between 60°N and 90N was 7.0 percent (370.35 to 344.29 DU),

and

that between 60°S and 90°S was 14.3 percent (335.20 to 287.23 DU).

11.4.2. Antarctic Stratospheric Changes

Between 1950 and 1980, no measurements from three ground-based stations in the

Antarctic showed ozone levels less than 220 DU, a threshold for deﬁning Antarctic

ozone depletion. Every Southern Hemisphere spring (September–November) since 1980,

GLOBAL STRATOSPHERIC OZONE REDUCTION 283

-10

-5

1980 1985 1990 1995 2000

Percentage difference in global ozone

from 1979 monthly average

Year

Mount Pinatubo

(June 1991)

El Chichón

(April 1982)

Figure 11.7. Percentage change in the monthly averaged global (90S to 90N) ozone column

abundance between a given month and the same month in 1979. Data were obtained from

the satellite-based Total Ozone Mapping Spectrometer (TOMS) and made available by NASA

Goddard Space Flight Center, Greenbelt, Maryland. No data were available from December

1994 to July 1996.

measurements of stratospheric ozone have shown a depletion. Farman et al. (1985) ﬁrst

reported depletions of more than 30 percent relative to pre-1980 measurements. Since

then, measurements over the South Pole have indicated depletions of up to 70 percent of

the column ozone for a period of a week in early October. The largest average depletion

for the month of September from 60 to 90°S since 1979 was 32.8 percent and occurred

in 2000. The largest depletion for the month of October from 60 to 90°S since 1979 was

38.3 percent and occurred in 1998 [Fig. 11.9(a)]. Most ozone depletion has occurred

between altitudes of 12 and 20 km. The large reduction of stratospheric ozone over the

Antarctic in the Southern Hemisphere spring each year is the Antarctic ozone hole. The

areal extent of the ozone hole is no

w greater than the size of North America.

Figure 11.9(b) shows the zonally and October-averaged ozone column abundance

versus latitude for 1979, 1999, and 2000. The ﬁgure shows that in 1999, the October

average over the South Pole was 131 DU, which compares with 286 DU in 1979. The

October average was slightly higher in 2000 than in 1999, but the September average

(not shown) was lower in 2000 than in 1999. October ozone levels over 45S (the lati-

tude of southern New Zealand, Chile, and Argentina) were 5 to 6 percent lower in

284 ATMOSPHERIC POLLUTION: HISTORY, SCIENCE, AND REGULATION

Figure 11.8. Mount Pinatubo eruption, June 12, 1991. Three days later, a larger eruption, the

second largest in the twentieth century, occurred. Photo by Dave Harlow, USGS.

1999 and 2000 than in 1979. Ozone levels over 30S (the latitude of Australia, South

Africa, Chile, Argentina, and southern Brazil) were about 3 percent lo

wer in 1999 and

2000 than in 1979. Temporary losses of ozone over these countries have caused con-

cern due the effects of the resulting enhanced UV-B radiation on health (Section 11.9).

11.4.3. Arctic Stratospheric Changes

Since 1979, the stratospheric ozone layer over the North Pole has declined during

the Northern Hemisphere late winter and spring (March–May). This reduction is

the Arctic ozone dent. Figure 11.9(a) indicates that the Arctic ozone dent in March

has consistently been less severe than has the corresponding springtime Antarctic

ozone hole in October. Figure 11.10, however, shows that ozone levels over the

Arctic during March 2000 were nearly 16 percent lower than were those during

GLOBAL STRATOSPHERIC OZONE REDUCTION 285

-40

-20

1980 1984 1988 1992 1996 2000

Percentage difference in ozone

from 1979 monthly average

Year

Mount Pinatubo

(June 1991)

60–90°N

March

El Chich´on

(April 1982)

60–90°S

October

100

150

200

250

300

350

400

450

500

-90 -60 -30 0 30 60 90

Ozone (DU)

Latitude (degrees)

October zonal average

1999

1979

2000

(a) (b)

Figure 11.9. (a) Percentage difference in March-averaged 60 to 90N and October-averaged 60 to 90S

ozone column abundances between the given year and 1979. (b) Variation with latitude of October month-

ly and zonally averaged column abundances of ozone in 1979, 1999, and 2000. Data were obtained from

the satellite-based Total Ozone Mapping Spectrometer (TOMS) and made available by NASA Goddard

Space Flight Center, Greenbelt, Maryland. No data were available from December 1994 to July 1996.

200

250

300

350

400

450

500

-90 -60 -30 0 30 60 90

Ozone (DU)

Latitude (degrees)

March zonal average

2000

1979

1999

Figure 11.10. Variation with latitude of March monthly and zonally averaged column abun-

dance of ozone in 1979, 1999, and 2000. Data were obtained from the satellite-based Total

Ozone Mapping Spectrometer (TOMS) and made available by NASA Goddard Space Flight

Center, Greenbelt, Maryland.

March 1979. Ozone levels over the Arctic in March 1999 were nearly the same as

those in March 1979.

11.4.4. Summary of Effects of Ozone and Air Pollution Changes

on UV Radiation

The yearly global reduction and seasonal regional destruction of stratospheric ozone

affect the amount of UV radiation reaching the surface. In comparison with their levels

in the 1970s, levels of ground UV-B radiation in 1998 were about 7 percent higher in

Northern Hemisphere midlatitudes in winter and spring, 4 percent higher in Northern

Hemisphere midlatitudes in summer and f

all, 6 percent higher in Southern Hemisphere

midlatitudes during the entire year, 130 percent higher in the Antarctic in the Southern

Hemisphere spring, and 22 percent higher in the Arctic in the Northern Hemisphere

spring (Madronich et al., 1998). Localized measurements of UV at Lauder, New Zealand

(45S), for example, show that surface UV-B radiation doses during the summer of

1998–1999 were 12 percent higher than they were in the ﬁrst years of the 1990s

(McKenzie et al., 1999).

11.5. EFFECTS OF CHLORINE ON GLOBAL OZONE REDUCTION

Ozone reductions since the late 1970s correlate with increases in chlorine and bromine

in the stratosphere. Molina and Rowland (1974) ﬁrst recognized that anthropogenic

chlorine compounds could destroy stratospheric ozone. Since then, scientists have

strengthened the links among global ozone reduction, Antarctic ozone depletion, and

the presence of chlorine- and bromine-containing compounds in the stratosphere.

11.5.1. CFCs and Related Compounds

The compounds that play the most important role in reducing stratospheric ozone are

chloroﬂuorocarbons (CFCs). Important CFCs are identiﬁed in Table 11.2. CFCs are

gases formed synthetically by replacing all hydrogen atoms in methane [CH

(g)] or

ethane [C

(g)] with chlorine and/or ﬂuorine atoms. For example, CFC-12

[CF

(g), dichlordiﬂuoromethane] is formed by replacing the four hydrogen atoms

in methane with two chlorine and two ﬂuorine atoms.

11.5.1.1. Invention of CFCs

CFCs were invented on a Saturday afternoon in 1928 by Thomas Midgley and his

assistants, Albert L. Henne (1901–1967) and Robert R. McNary (1903–1988), at the

Thomas and Hochwalt Laboratory, 127 North Ludlow Street, Dayton, Ohio. Midgley

is the same scientist who invented tetraethyl lead (Ethyl) gasoline (Section 3.6.9).

Some argue that Midgley’s inventions led to the two greatest environmental disasters

of the twentieth century.

Midgley and his assistants developed CFC-12 effectively on the same day that a

representative of General Motors’ Frigidaire division asked Midgley to ﬁnd a nontoxic,

nonﬂammable substitute for an existing refrigerant, ammonia, a ﬂammable and toxic

gas. CFC-12 and subsequent CFCs were inexpensive, nontoxic, nonﬂammable, nonex-

plosive, insoluble, and chemically unreactive under tropospheric conditions; thus, they

became popular. Midgley demonstrated the nontoxic and nonﬂammable properties of

286 ATMOSPHERIC POLLUTION: HISTORY, SCIENCE, AND REGULATION

GLOBAL STRATOSPHERIC OZONE REDUCTION 287

Table 11.2. Mixing Ratios and Lifetimes of Selected Chlorocarbons,

Bromocarbons, and Fluorocarbons

CFCl

(g) CFC-11 Trichloroﬂuoromethane 270 45

(g) CFC-12 Dichlorodiﬂuoromethane 550 100

CFCl

Cl(g) CFC-113 1-Fluorodichloro, 2-diﬂuorochloroethane 70 85

ClCF

Cl(g) CFC-114 15 220

ClCF

(g) CFC-115 5 550

ClH(g) HCFC-22 Chlorodiﬂuoromethane 130 11.8

CFCl

(g) HCFC-141b 6 9.2

Cl(g) HCFC-142b 2-Diﬂuorochloroethane 8 18.5

CCl

(g) Carbon tetrachloride 100 35

CCl

(g) Methyl chloroform 90 4.8

Cl(g) Methyl chloride 610 1.3

HCl(g) Hydrochloric acid 10–1,000 1

Estimated

Tropospheric Overall

Mixing Ratio Atmospheric

Chemical Formula Trade Name Chemical Name (pptv) Lifetime (yrs)

Br(g) H-1301 Triﬂuorobromomethane 2 65

ClBr(g) H-1211 Diﬂuorochlorobromomethane 2 11

BrCF

Br H-2402 1-Diﬂuorobromo, 2-diﬂuorobromoethane 1.5 22–30

Br(g) Methyl bromide 12 0.7

FCF

(g) HFC-134a 1-Fluoro, 2-triﬂuoroethane 4 13.6

(g) Perﬂuoroethane 4 10,000

(g) Sulfur hexaﬂuoride 3.7 3,200

Sources: Shen et al. (1995); Singh (1995); WMO (1998); Mauna Loa Data Center (2001).

Other Chlorocarbons

Other Chlorinated Compounds

BROMOCARBONS

Halons

Other Bromocarbons

FLUOROCARBONS AND FLUORINE COMPOUNDS

Hydroﬂuorocarbons (HFCs)

Perﬂuorocarbons (PFCs)

Other Fluorinated Compounds

Hydrochloroﬂuorocarbons (HCFCs)

CHLOROCARBONS AND CHLORINE COMPOUNDS

Chloroﬂuorocarbons (CFCs)