Reactive nitrogen in the troposphuere Chemistry and transport of NO, and PAN
Hanwant B. S i NASA Ames Research Center Moffen Field, Calif: 94035 During the middle part of this century, it became evident that human activities were altering the chemical composition of the Earth's ahnosphere. Photochemical smog was discovered and found to be. a byproduct of reactions involving hydrocarbons, nitrogen oxides (NOx = NO NOJ and sunlight. It was determined that the hydrocarbon-NO, precursors were principally emitted by anthropogenic sources and that ozone (0,) was a major component of smog. In addition, there is evidence that background levels of carbon dioxide, halocarbons, ozone, methane, carbon monoxide, and probably the hydroxyl radical (OH) are steadily changing in the Earth's atmosphere. These changes are. intricately linked with changes in climate and biosphere habitability. Furthermore, nitrogen oxides play a central role in many of these processes of local, regional, and global change. Our arrival at more complete knowledge of the chemistry and transport of reactive nitrogen species is one essential step toward understanding and controlling the effectsof human activity on the air environment.
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Sources and chemistry
Nitrogen oxides are emitted as nitric oxide (NO) from a variety of sources. The estimated emissions of NO, have large uncertainties associated with them, but lhble 1 shows that 7690% of emissions are from anthropogenic sources (1-3). Thus, substantial change to this natural cycle has already occurred. Although most of the emissions of reactive nitrogen occur as nitric oxide, it is typically converted to nitrogen dioxide in minutes by reaction with 320 Envimn. Sci. Technol., MI. 2.1. NO. 4. I987
1 This article not aub)en to U.S. oopyright. Published 1987 Amrican Chemical Society
ozone. Nitrogen dioxide in turn absorbs ultraviolet light at relatively long wavelengths and photolyzes rapidly, eventually reforming ozone (Equation 1). These reactions provide a basis for ozone formation even though no net ozone is produced by these reactions alone, and they are called a null reaction sequence. If there were no other reactive species, a stationary state could be reached in a matter of minutes (Equation 2). Equation 1 is not the only set of reactions that takes place in the troposphere (0-12 km above the Earth's surface). Nitric oxide also can react with species other than ozone. The most important of these is the hydroperoxy radical @IO2), which originates from photochemistry involving the hydroxyl radical and hydrogen atoms. Various organic peroxyradicals (RO,), produced in a variety of ways from hydrocarbon oxidation, are also important. The sequence shown in Equation 3 is typical of alkane (RH) oxidation. These reactions do produce ozone, and because nitric oxide is oxidized to nitrogen dioxide without removing mne, they are in fact responsible for producing much of the m n e associated with photochemical smog. This is because nitric oxide and nitrogen dioxide are cycled catalytically; the amount of m n e produced depends on the N@I NO ratio. Hence, the concentration of omne created is greater than the original concentration of NO,. Furthermore, breakdown of the resultant aldehydes and ketones yields other
products, the most important of which is the peroxyacetyl radical, CH3C(0)-0-0. This radical reacts quickly with nitrogen dioxide to form peroxyacetyl nitrate-PAN (Equation 4) ( 4 . 9 . The most common carbonyl compounds that yield the peroxyacetyl radical are acetaldehyde (CH3CHO), acetone (CH3COCH3), and biacetyl (CH3COCOCH3),all of which are hyproducts of hydrocarbon oxidation (Equation 5). Acetaldehyde is an oxidation product of virtually all n-alkanes and alkenes. Acetone is similarly produced from C3H8, iC4H,o, iC5H12; biacetyl is a major product of oxylene oxidation. Although PAN is probably the most abundant oxidation product, other organic nitrate prcducts, such as peroxy-
nitric acid @I&N02), peroxypropionyl nitrate 0, and peroxybenzoyl nitrate (PBzN) are formed as well. The role of the hydroxyl radical is important in Equation 5 . This radical is believed to be ubiquitous and is formed hy ozone photolysis. However, the recycling of other free radicals and the photolysis of secondary products are additional important sources of the hydroxyl radical. The reactions shown in Equation 6 determine the hydroxyl radical's principal routes of formation. In addition to the importance for ozone and PAN formation, NO, greatly influences the atmospheric concentrations of hydroxyl radical. The recycling of the hydroperoxy radical via the H02 NO reaction is the most important in this regard. Nitrous acid @IN@) is relatively stable at night, but it is eas-
+
Cghtning Mtcmbia)activity in soils Input from thestratosphere Total
PAN)
I
-
Environ. k i. Technol., Vol. 21. No. 4, 1987 921
ily photolyzed and therefore l i l y to be a major source of early-morning OH. The hydroxyl radical is one of the most reactive species in the atmosphere and is responsible for the chemical removal of a large number of man-made and naturally occumng chemical species injected into the atmosphere. In the presence of hydrocarbons and NO,, net photosynthesis of ozone and PAN occurs. In the remote free troposphere (over the ocean, for example), the hydrocarbons involved may be simply methane and ethane or carbon monoxide. Significant ozone and PAN synthesis in the remote troposphere has
been postulated (1, 6, 7). In the polluted ahnosphere and in the boundary layer (the layer below the inversion layer, from the ground to 1-2 km above the Earth), highly reactive hydrocarbons of anthropogenic origin (such as a l h e s , alkenes, and aromatics) and of biogenic origin (such as isoprene and terpenes) are more important and dominate the reaction more than methane and ethane do. Along with ozone, PAN is present in polluted atmospheres and has been suggested as an indicator of photochemical smog (8).U d i i ozone, PAN does not have large natural sources in the strato-
Photochemically active nitrogen s p e c i e s in t h e troposphere
h e following are the important reactive nitrogen species found in the t r o w
sphere. Nitric oxide (NO) and nitrogen dioxide (NOZ). Primary emissions of these species occur nearly universally as NO. Both NO and NOz recycle easily and often are referred to as NO. (NO + NOZ).NO, is estimated to have a lifetime of ummer and a few days in winter. produced primarily from the oxidation of ime: Nighttime formation occurs through N205and liquid water. "Os is a sink for ry and wet deposition processes over a Peroxynitric acid (H02NOZ).H02N02has never been measured. Its lifetime depends largely on temperature and may be several months in the upper trow sphere. It is much tw unstable to be present at significant levels in the boundary layer. Paroxyacetyl nitrate (PAN, CH3C(0)-O-O-NOZ); peroxypmplonyl nlate (PPN, CzH5C(0)-0-0-NOz); peroxybenzoyl nitrate (PBzN. BH5C(0)-O-0-N02); peroxymethacrylyl nitrate (MPAN, CHz= (CH)sC(0)-O-O-NOz); peroxy(hydroxy)acetyI nitrate (HPAN, HOCH2C(0)-O-O-NOZ), PAN is the most abundant pernitrate species and has been shown to be globally ubiquitous. The concentration of PPN is only 510% that of PAN. The lifetime of PAN can range from a few hours to several months, depending on temperature. It is not a sink for NO, as it can rerelease NOzunder warmer conditions. Its stability greatly increases with falling temperatures. Nighttime PAN losses occur principally in the boundary layer through surface deposition. MPAN and HPAN are possible products of isoprene oxidation. NIrous acid ("03. HN02is easily photolyzed to release OH and NO during daytime. it has been measured only in highly polluted atmospheres at night ( 1 2 73). Because it can photolyze rapidly (1-2 h) it may be an important source of Barb-morning OH. Dlnitmgan pentoxide (N2O5).This species provides an important route to nighttime "Os formation, in which heterogeneous chemical processes are involved. N& h a s not been measured directly, and it is not expected to be present during daytime. Nitrate radical (NO3). NO, is stable only at night. It has been detected at concentrationsthat are always much lower than predicted and at levels of about 1% of NO2 present in the troposphere (14). it is a highly reactive radical that can play a role similar to that of OH during daytime. NOs has been observed in polluted and remote atmospheres (15). Total reacthre nitrogen (NO,) is the sum of NO, NOz, "Os, HOZNOz,the PANS, " 4 , 2NZO5,NOa and the gaseous nitrates. Other "odd nitrogen species, such as NHa,HCN, and CH3CN,are present in the troposphere as well. Ammonia is the only basic gas in the atmosphere, and it acts as an acid neutralizer. Ammonia may be a small source or sink of NO, but the chemistry involved is highly uncertain (1). Both CH&N and HCN are relatively inert (1-2-yr lifetimes) and do not appear to play any important role in tropospheric chemistry ( 1 4 77). The most abundant nitrogen oxide species, NzO, is not an odd nitrogen species. It is virtually inert in the troposphere. In the stratosphere, however, N20 photolyzes and reacts with free radicals to provide a major source of NO., which plays a criiical role in maintaining the stratospheric ozone layer. 22 Environ. Sci. Technol., Vol. 21, No. 4, 1987
sphere (12-50 km above the Earth's surface). It therefore offers advantages over ozone as a specific indicator of hydrocarbon-NO, photochemistry. PAN also has been associated with crop damage, most of which occurs during the daytime. Although potentially phytotoxic episodes (> 15 ppb PAN for 4 h during daytime) occur frequently in Southern California, PAN levels in much of the rest of the world are substantially lower (9). Lovelock has suggested that PAN also is likely to be involved in the epidemiology of skin cancer (10). Nitric acid, a key component of acid deposition, is for all practical purposes a product of photochemistry that is removed by precipitation and surface deposition over a period of 1-10 days (1).During the daytime, nitrogen dioxide reacts rapidly with the hydroxyl radical to produce nitric acid over a period of a few hours. Although nighttime gas-phase and liquid-phase reactions lead to acid formation, it is the heterogeneous reaction of Nz05 with liquid water that is the most important (F.quation 7). In theory, gas-phase nitric acid can further participate in the nitrogen chemistry of the atmosphere via the reactions shown in Equation 8. In practice, the rates of t h m reactions are so slow that this is a less important role in the lower troposphere. In the upper troposphere, however, HN03 can exist in equilibrium with nitrogen dioxide if air masses are isolated from precipitation scavenging for long periods of time (Quation 6. There is further reason to believe that there is an atmospheric equilibrium involving gaseous HNO,,gaseous NH3, and solid ammonium nitrate, although available data are insufficient to fully support this contention (11). This HN03-NH3-N&N03 equilibrium is highly sensitive to tempenlture, and equilibrium mass concentrations in the gas phase increase fourfold over a range of 20-30 OC (Equation 9).
-
__
Distribution and transport Although NO, is principally emitted as nihic oxide, it reacts quickly to transform into a variety of oxidized species. There are a number of important photochemically active nitrogen species liely to be present in the troposphere (12-17). A list of these species is shown in the sidebar, not all of them have been identified or measured. Based on reaction with the hydroxyl radical alone, the lifetime of NO, can be estimated to be less than one day in summer and about one week in winter. A number of field studies conducted across the continental United States support this estimate, although in sum-
mer lifetimes of as little as 4-8 h have been determined (28, 19). This short lifetime l i t s the transport scale to a few hundred kilometers. Although nitric acid can be m p o r t e d over long distances, its loss occurs by highly variable physical processes of dry and wet deposition, and it plays a small role in promoting photochemistry. Except for the pernitrates, all active nitrogen species listed in the sidebar are short-lived and are not likely to be transported over distances of more than a few hundred kilometers. In recent years, the role of PAN as a carrier and reservoir of NO, has received considerable attention. This possibility arose from the understanding that PAN exists in equilibrium with nitrogen dioxide, and this equilibrium is greatly affected by small changes in temperature. As a first approximation, PAN’S behavior can be illuminated by the assumption that all peroxyacetyl
radicals are produced by the oxidation of acetaldehyde. The reaftions of Equation 10 describe the principal formation and removal processes for PAN. AU rate conslilnts in Equation 1oSe in units of cm3/molecule/s unless otherwise specified. They are taken from the literature, and they apply to mid-latiNde conditions found in the Northern Hemisphere (20). Equations 1 1-13 are derived by assuming that the peroxyacetyl radical is in steady sate. The lifetime of PAN (rpAN) is mathematically defined by Equation 12. Conceptually, it is conas indicative of venient to think of rpAN the PAN decay rate undcr conditions in which no new PAN synthesis occurs. The thermal decomposition reaction rate of PAN m.3) is highly temperature dependent and occurs within expected times (Ilk,)of I h at 298 K, 2 days at 273 K, 148 days at 250 K, and 42 years at 230 K. These kinetic considerations
NigMfimf
led to the proposal of a mechanism in which the principal form of reactive nitrogen in the middle and upper troposphere was shown to be PAN (7). Figure 1 shows modeled distributions of reactive nitrogen species in the summertime marine (Figure la) and continental tropospheres (Figure lb) at 45 ON. Aircraft measurements are in general agreement (20, 21). From the equilibrium reaction (k3,k.3) it is obvious that a mass of cold air from the upper troposphere would release nitre gen dioxide when it warms because the equilibrium shifts rapidly toward that species. The upper tropospheric reservoir of PAN could thus transport nitrogen dioxide to lower latitudes and lower altitudes, where warmer temperatures prevail. Consistent with the temperature dependence of rpAN, the upper tropospheric abundance of PAN has been shown to favor a strong seasonal cycle (21).
m
Envimn. Sci. Technol.. MI. 21, No. 4, 1987 323
l b o aaamonal observauons on me such that the overall kte of loss, infate of PAN also must be noted. First, cluding production and destruction, is drastically reduced. Hov has simnlated PAN photolyzes and reacts with the hydroxyl radical (jPAN.ks). These reac- PAN transport in air masses traveling tions place an upper bound of about fim the United Kingdom toward Oslo, Norway (23). Figure 2a shows ozone three months on its mean lifetime (Equation 12). For temperatures above and PAN transport; Figure 2b is a map 273 K, ks and jpANdo not compete with of the associated back trajectories. k., and can be neglected. Contrary to Many previous studies have shown that some published estimates, PAN’S life- PAN, like ozone, is an excellent indicatime in the upper troposphere never a p tor of the transport of photochemical proaches years. Second, at night when air pollution (8, 24, 25). When PAN nitric oxide and peroxyl radical concen- synthesis during transport is included, trations approach zero, PAN is almost the overall *fold loss of PAN (that is, the amount of time required to reduce infhtely stable (Equation 12). Thus, PAN is transported freely during the PAN concentrations from 100% to night, and deposition processes con- 37%)is estimated to be 3 days at 20 OC and 35 days at 5 OC. tinue to OCCUT in the boundary layer. Thus, although PAN can always be Measurements show that PAN is insoluble in acidic aqueous solutions and transported distances of > 10,oOo lan that surface deposition over the ocean is in the upper troposphere over several not appreciable (22). Over land, how- months, its transport in the lower tmpever, a deposition velocity of 0.25 c d s osphere and in the boundary layer is a has been measured (22). Therefore, function of prevailing temperatures and some nighttime loss of PAN occu~sin continued synthesis from precursors. In the boundary layer as a result of surface the absence of continued synthesis, its removal. Because the lifetime of PAN lifetime in the boundary layer is not is a few months in the upper tropo- likely to exceed one or two days. This sphere, it can be freely transported therefore limits PAN transport to diswithin each hemisphere. Little ex- tances similar to those for NO,--a few change between hemispheres is possi- hundred kilometers. The ratio of PAN to NO, increases as ble, however, because this process the air mass ages because PAN is a bytakes about a year. At warmer temperatures, PAN is product of hydrocarbon-NO, chemislonger lived, because of the reverse re- try. This is best demonmated in a smog action &), than its thermal dissociation chamber where no new emissions ocwould imply. For a typical cur. In one such experiment, the PAN/ rate (I.,) N o m a ratio of 0.2; T ~ A N= 3/k4 NO, ratio rose from 0 to 3 in 5 h (26). (Equation 12). The formation reaction A direct demonstration of this phenom(k,) causes PAN to be nearly three enon as it was measured in summer and times as stable as the decomposition re- fall in the Rocky Mountains is shown in action does. Despite this, rpAN has been Figure 3 (27). Depending on the prevailing winds, calculated to be only 4-5h at 2OoC and 2 1 h at 10DC. However, like this site can receive polluted air (easterozone, PAN is continually synthesized lies) to ai^ that has seen little or no recent pollution (westerlies). It is evident during transport. Equation 11 shows that daytime from Figure 3 that the PANINO, ratio long-range transport of PAN is nearly increases as the air becomes cleaner always associated with continued syn- (low NO, and westerly trajectories) and thesis. The magnitude of the first term indicates an aged air mass. As exon the right-hand side of the equation is pected, this role of PAN is much more 324 Environ Scl. Technol.. Val. 21. No. 4, I987
i m p o m t in autumn than it is in summer, presumably because of the much greater stability of PAN associated with lower temperatures. Even higher ratios have been reported in arctic air masses at 82 ON (28. For the higher NO, levels typical of moderately polluted conditions these ratios are comparable to those that have been reported in the literature (19.29, 30). In addition, PAN and PAN-lie compounds also are sources of organic peroxyradicals during the day and at night (Equation 13). For one experiment (23, it is estimated that daytime CH,C(O)-0-0 concentrations of 2-3 x lo7 molecules per cm3prevailed under relatively clean tropospheric wnditions associated with average temperatures of 290 K at 3 lan altitude (Equation 13). Measured PAN-NO, behavior implied that acetaldehyde was present in concentrations of 0.1-0.3 ppb. It also is clear from Equation 13 that the CH,C(O)-0-0 concentration would be highly dependent on temperature.
Reactive nitrogen budget Although a number of reactive nitrogen species are present in the air, the bulk of the odd nitrogen should he accounted for by NO, N@, H N a , PAN, and PPN. This assertion is tested by measuring these species individually and comparing them with direct measurements of total reactive odd nitrogen (NO,). The instrument used to make such measurements is basically an NO chemiluminescent insaument equipped with a catalytic converter that reduces all odd nitrogen species (except NH,, HCN, and CH3CN) to NO. Figure 4 shows the CNOy,INOyratio @NOyiis the sum of measured odd nitrogen species) based on measurements taken at Niwot Ridge, Colo., at an altitude of 3 km in the summer and fall of 1984 (31). It is clear from this figure that significant quantities of odd nitrogen (45% in summer and 10% in winter) are still unaccounted for. Other ex-
(02)
CH3CH0 + OH CH3C(O)-O-O CH&(O)-0-0
+
> CHsC(O)-O-O
NO2 , A CH,C(O)-O-O
+ NO
CH3
+ OH PAN + hv
PAN
+
H20
(PAN)
+ COP +
NO2
Products
k2 = 1.6 x IO-’’ = 6x 13330 IC3 = 1.12 x 10’6exp (--T k3
k4
= 1.4 X 10.”
651 k5 = 1.23 x lW’*exp (-T)
Products ,,j
= 1.2 x IO-’S-’
Envimn. Sci. Technol., MI. 21.
No. 4, 1987 a25
-_edwith a characteristic time ( l / k ) of 31 days and 10 days, respectively. Assuming steady state, methyl nitrate concentrations can be estimated at 15-50% those of PAN, [CH,0N02] = ks [PAN]/jcH,oNoz. at temperatures of 290-298 K. This fraction may increase if the CH30N@ photolysis quantum yield is less than unity, thus resulting in a longer estimated lifetime. A second possibility is the presence of other PAN-lie compounds that have not yet been measured. It is now known that PPN abundance is only about 5% that of PAN. The possibility should be considered that biogenic hydrocarbons such as isoprene and a-pinene also can sequester reactive nitrogen. Such biogenic hydrocarbons are highly reactive, are emitted most often in summer, and typically are present only in the boundary layer. The chemistry of isoprene clearly suggests that its major oxidation products (methyl vinyl ketone and methacrolein) can react further to produce PAN and a variety of PAN-like compounds (35).None of these compounds has been measured to date, but it is probable that they contributeto the NO, defipiency in the boundary layer during warm atmospheric conditiw. A third possibility arises from the formation of alkyl nitrates (RON@) directly from the reaction of organic peroxyradicals with nitric oxide @quation 15). Experimental data show that for alkanes of carbon number less than 3, no significant alkyl nitrate formation occurs. However, at carbon numbers of 4 or more alkyl nitrate formation becomes significant (36). Ratios for k7/b of 0.1, 0.25, and 0.5 have been reported for C4, C6, and C8 alkanes, respectively. The alkyl nitrate yield is lower at lower pressures and higher temperatures. In the boundary layer, sizable 32s Environ. Sci. Technol., Vol. 21, No. 4. 1987
concentrations of C4-CI0 anthropogenic hydrocarbons and biogenic hydrocarbons (>C& such as isoprene and other terpenes, could produce significant amounts of alkyl nitrates. High concentrations of organic peroxyradicals have been detected in moderately polluted to relatively clean environments (37). It is expected that many of these radicals have high carbon numbers. In the winter, reduced temperatures, slower photochemistry, and reduced biogenic emissions should combine to make alkyl nitrates far less abundant. Although the proposed presence of alkyl nitrates and peroxynitrates seems plausible, the possibility that some asyet unknown odd nitrogen species are present cannot be ruled out. Possible deficiencies of reactive nitrogen in the upper troposphere have not yet been investigated.
Laoking ahead Although our knowledge of nitrogen chemistry has improved greatly over the past two decades, a number of important questions must be answered. Among the most basic of these is the disagreement between measured and predicted N@/NO ratios @pation 4), especially in clean atmospheres. Most models suggest N@/NO ratios that are significantly lower than the measured values (1). Is this a problem associated with measurements, or is the simplest step in nitrogen chemistry not well understood? It is clear that there are undiscovered reactive nitrogen species present at least in the boundary layer. What are these species? Are they present only in the lower troposphere, or are they globally pervasive? What is the chemistry of these missing nitrogen compounds? What role do they play? If NO, has a lifetime of only a few hours,
how is it transported globally? Does the upper tropospheric PAN reservoir pre vide a principal mechanism for transport of NO,? No measurements of pernitric acid have been made to date, but it is expected to be important in the upper troposphere. The vertical distribution of the PANINO, and H@N@lNO, ratios have not been determined in the troposphere. Theory suggests that these ratios should increase with height. This further suggests that in the middle and upper troposphere most of the reactive nitrogen exists as PAN rather than as N@. Nighttime chemistry is only beginning to be understood. Nitrate radicals (NO,) are found to be less abundant than predicted. What are some of the processes that remove NO, from the troposphere? Measurements of N20, have not been possible to date, although it is recognized as a key intermediate in nighttime nitric acid formation. Nitrogen chemistry is receiving a g r a deal of attention both here and abroad because of the important role it plays in the troposphere. In one recent NASA-sponsored program (the Global Tropospheric Experiment-Chemical Instrumentation Est and Evaluation), aircraft measurements were made of PAN, NO, N@, HNO,, NH,, O,, and other important species. Although the results are not in, it is hoped that these and other planned experiments will go a long way toward answering many of these important questions of tropospheric chemistry. Acknowledgment Portions of thls work are supported by the NASA Global ’lkopospheric Experiment and by the National Scienfe Foundation under grant ATM8606269 to SRI International. We are grateful to F. Fehsenfeld and coworkers at the National Oceanic and Atmospheric Administration in Boulder, Colo., and to B. Ridley and coworkers at the National Center for Atmospheric RG seaxh, also in Boulder, for their particiption in many of the field sNdies from which this article draws extensively. This article has been reviewed for ouitability as an ES&T feature by D. H.Stedman, University of Denver, Denver, Colo. 80208; and by Thomas J. Kelly. Environmental Chemistry Division, Bmokhaven National Laboratory, Upton, N.Y. 11973. Reference ( I ) Logan, 1. A. 1. Gcophys. Res. 1983, 88, 10785-807. (2) Cantrell, B.; Ludwig, EL.; Siogh, H.B. “A Review of the Fate of NO, and its Role in Rural Ozone Formation,” SRI Project 3&23; SRI-Inlernational: Menlo Park, Calif.,
(3) “National Air Qu$ty and Emissions ltends Report, 1984, EPA-45014-86-001; EPA. Research Triangle Park, N.C., 1986. (4) Hendry, D.0.; Kenley, R. A. 1. Am. Ckm. Soc. 15”/7,99,3198-99. (5) Con, R. A.; Roffey, M.1. Emiron. Sci. Tcchnol. 1977,II. 900-906.
(6) Fishman, 1.; Crutzen, !F 1. Norure 1978. 274. 855-58. (7) SinKh. H. B.;Hanst. I? L. Geoohvs. Res. k r r . 1981, 8, 941-44. (8) Stephens. E. R. Ad”. Environ. Sci. Echno!. 1%9,1. 11946. (9) Temple, P.I.; Taylor, 0.c. ~ m s Envi. ron. 1983, 17, 1583-87. (10) Lovelock, 1. E.Ambio 1977.6, 131-33. (11) Steison. A. W.; Friedlander, S.K.; Seinfeld. I . H. Armos. Environ. 1979, 13, 369-
.,
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7 , I..
(12) Harris, 0. W. et al. Environ. Sei. Techno/. 1982,16,414-19. (13) Sjodin, A.; Ferm, M. Arms. Environ. 1985, 19.985-92. (14) Noxon. I. E 1. Geophys. Res. 1983, 88, 11017-21. (15) Atkinson. R.; Winer, A. M.; Fins,J. N., Ir. Armos. Environ. 1986.20, 331-39. (16) Cicerone, R.1.; Zellner. R. 1. Geophys. Res. 1983.88, 10689-96. (17) Arils, E.;Brasseur, G . 1. Geophys. Res. 1986,91.4003-16. (l!i Spicer, C. W. Science 1982, 215. 10957,.
(19) Altshuller, A. !F Armos. Environ. 1986, 20,245-68. (20) Kasting. 1. F.; Singh. H. B. 1. Geophys. Res. 1986.91, 13239-56. (21) Singh. H. B.; Salas, L.1.; Viezee. W. Nature 1986,321, 588-91. (22) Garland, 1. A,; Penkett, S. A. Armos. Enw o n . 1976, IO. 1127-31. (23) Hov, 0. 1. Amtos. Chem. 1984, 1. 187202. (24) Nielsen, T et al. Nnruie 1981, 293. 55355. (25) Brice, K. A. et al. Armos. Environ. 1984, 18. 2691-2702. (26) Spicer, C. W. Environ. Sci. Technol. 1983, 17. 112-20. (27) Singh, H. E. el al. Norwe 1985, 318. 347-49. (28) Bottenheim. I. W.; Gallant, A. 0.; Brice, K. A. Ceophys. Res. Len. 1986.13, 113-16. (29) Lnnneman, W. E.; Bufalini. 1.1.; Seila, R. L. Environ. Sci. Technol. 1976, IO. 37480.
(30i~Spicer.C. W. Sei. Torol Environ. 1982, 24. 183-92. (31) Fahey, D.W. et al. 1. Geophys. Res. 1986.91, 9781-93. (32) Senum. G . I.; Fajer. R.; Gaffney, 1. S. 1. Phys. Chem. 1986.90, 152-56. (33) Gaffney, J . S. et ai. Inr. 1. Chem. Kinel, 1986.18. 399-407. (34) Taylor, W. D.et al. Im. 1. Chrm. Kim. 1980,IZ. 231. (35) Lloyd, A. C. etal. Amos. Environ. 1983, 17 1931.50 ., . ... ...
~
(36) Carter. W. L.; Atkinson, R. 1. Arms. Chem. 198537, 377-405. (37) Miheicic, D.;Musgen, F!; Ehhalt, D.H. 1. Armos. Chem. 1985.3. 341-61.
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