Conversion of Fogwater and AerosolOrganic Nitrogen to Ammonium,Nitrate, and NOx during Exposure toSimulated Sunlight and OzoneQ I Z H A N G † A N D C O R T A N A S T A S I O *

Atmospheric Science Program, Department of Land, Air, andWater Resources, University of California, One Shields Avenue,Davis, California 95616-8627

Although organic nitrogen (ON) compounds are apparentlyubiquitous in the troposphere, very little is known abouttheir fate and transformations. As one step in addressingthis issue, we have studied the transformations of bulk(uncharacterized) organic nitrogen in fogwaters and aerosolaqueous extracts during exposure to simulated sunlightand O3. Our results show that over the course of severalhours of exposure a significant portion of condensed-phaseorganic nitrogen is transformed into ammonium, nitrite,nitrate, and NOx. For nitrite, there was both photochemicalformation and destruction, resulting in a slow net loss.Ammonium and nitrate were formed at initial rates on theorder of a few micromolar per hour in the bulk fogwaters,corresponding to formation rates of ∼10 and 40 ng m-3 h-1,respectively, in ambient fog. The average initial formationrate (expressed as ng (m of air)-3 h-1) of NH4

+ in theaqueous extracts of fine particles (PM2.5) was approximatelyone-half of the corresponding fogwater value. Initialformation rates of NOx (i.e., NO + NO2) were equivalentto ∼2-11 pptv h-1 in the three fogwaters tested. Althoughthe formation rates of ammonium and nitrate wererelatively small as compared to their initial concentrationsin fogwaters (∼200-2000 µM) and aerosol particles (∼400-1500 ng m-3), this photochemical mineralization and“renoxification” from condensed-phase organic N is apreviously uncharacterized source of inorganic N in theatmosphere. This conversion also represents a newcomponent in the biogeochemical cycle of nitrogen thatmight have significant influences on atmospheric composition,condensed-phase properties, and the ecological impactsof N deposition.

1. IntroductionOrganic nitrogen (ON) has been measured in precipitation(1-5), dry deposition (1, 5, 6), cloud waters (7), fogwaters (8),and atmospheric particles (4, 9-11). The ubiquity of thesecompounds suggests that they might play important roles inatmospheric chemistry and in the biogeochemical cycling ofN. In addition, since organic forms often represent asignificant and bioavailable portion of the total N indeposition (3-7, 12), atmospheric ON can be a nutrientburden to aquatic and terrestrial ecosystems (5, 13-16).

Past studies have shown that ON compounds are subjectto chemical and photochemical transformations in thetroposphere, forming products that might potentially influ-ence the properties of atmospheric condensed phases andthe bioavailability of N in deposition (1, 17-21). For instance,the more bioavailable, lower molecular weight compounds,such as amino acids and urea, usually account for less than20% of the atmospheric ON pool (8, 9, 11, 22, 23), whilelarger and less bioavailable species, such as humic substances,appear to be more abundant (24-26). Although biologicallyrefractory (27), humic substances are photochemically reac-tive (28, 29), and exposure to sunlight might cause them todecompose into smaller and more bioavailable molecules.Indeed, sunlight illumination of organic matter from surfacewaters leads to the formation of amino acids and ammonium(30-33). While similar reactions might occur in atmosphericdrops and particles, no such studies have been performed,and little is known about the transformation rates and fatesof atmospheric organic nitrogen.

The photochemical transformations of specific ON com-pounds, such as amino acids, nitrogen heterocycles, andnitroaromatics, have been studied in atmospheric samplesor under atmospherically relevant conditions (1, 19-21).However, studies of individual compounds are limited bythe fact that the bulk of ON in atmospheric drops and particlesis uncharacterized (8, 9, 11, 12). Furthermore, because thereactivity of individual ON compounds varies widely (1, 19),it is currently infeasible to extrapolate from single-compoundstudies to transformation rates of bulk (and largely unchar-acterized) organic nitrogen in the atmosphere. Therefore,we initiated this study to examine photochemical transfor-mations of bulk organic N in atmospheric fogwaters andaerosol particles. We describe here a preliminary set ofexperiments to examine whether bulk atmospheric organicN is transformed during exposure to sunlight and, if so,whether these reactions form inorganic N. In addition,because ozone can be a major sink for some atmosphericorganic nitrogen compounds (19, 22), the influence of ozoneon organic nitrogen transformations and inorganic N forma-tion was also investigated.

2. Experimental Methods2.1. Samples, Materials, and Analyses. Since details of oursampling dates and times, analytical methods, and samplecharacteristics have been presented in prior reports (8, 9,22), only the main points are given here. Fogwaters werecollected at the Davis, CA, NADP site (CA88; 38°33′ N, 121°38′W) during winters from 1997 to 2001 using a Caltech ActiveStrand Cloudwater Collector (CASCC2). Immediately aftercollection, samples were filtered (0.22 µm Teflon) and storedfrozen (-20 °C) in HDPE bottles. PM2.5 samples were collectedat the same location during August 1997-July 1998. Water-soluble particle components were extracted into Milli-Q water(g18.2 MΩ-cm) by sonication, followed by filtration (0.22µm Teflon) and frozen storage (-20 °C) in HDPE bottles.

Concentrations of NH4+, NO3

-, and NO2- were analyzed

using a Dionex DX-120 ion chromatograph (IC) with con-ductivity detection. Dissolved organic nitrogen (DON) wasdetermined as the difference in inorganic N concentrationsin a given fogwater before and after adjustment to pH ≈3and 24 h of illumination with 254 nm of light (to convertDON to inorganic forms): [DON] ) ([NH4

+] + [NO3-] +

[NO2-])after 254 nm hν - ([NH4

+] + [NO3-] + [NO2

-])before 254 nm hν.Concentrations of each N species in all samples wereconsiderably larger than the corresponding detection limits(i.e., ∼0.1 µM for NH4

+, NO3-, and NO2

- and 1.0 µM for DON).

* Corresponding author e-mail: canastasio@ucdavis.edu; tele-phone: (530)754-6095; fax: (530)752-1552.

† Present address: Cooperative Institute for Research in Envi-ronmental Sciences (CIRES), 216 UCB, University of Colorado,Boulder, CO 80309-0216.

Environ. Sci. Technol. 2003, 37, 3522-3530

3522 9 ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 37, NO. 16, 2003 10.1021/es034114x CCC: $25.00 2003 American Chemical SocietyPublished on Web 07/19/2003

Concentrations of total N (TN) were determined from thesum of N in the sample solution (([NH4

+] + [NO3-] +

[NO2-])after 254 nm hν, see section 2.2) plus, in the experiments

with ozone exposure, the concentrations of gas-derived NH4+,

NO3-, and NO2

- found on the denuders and reaction flask(section 2.3). Note that in the samples exposed to ozone, thesum of nitrate and nitrite includes photoproduced NOx thatwas converted to NO3

- and NO2- in the reaction flask.

2.2. Dissolved Inorganic N (DIN) Formation duringSimulated Sunlight Illumination. 2.2.1. Lyophilized Fog andPM2.5 Samples. Because concentrations of NH4

+ were gener-ally high in our samples (8, 9), we were initially concernedthat it would be difficult to distinguish between the photo-formation of DIN and analytical variations. To increase ourability to see the photochemical formation of NH4

+, the firstgroup of fogwaters and aerosol extracts were lyophilized (i.e.,freeze-dried; 34) to reduce NH4

+ levels prior to illumination.To do this, NaOH (1 M in Milli-Q) was added to form NH4OHstoichiometrically (based on the initial measurement ofNH4

+); the sample was rapidly frozen in liquid nitrogen; andthen lyophilized to dryness on a Virtis 2SL freeze drier.Afterward, the lyophilized residue was reconstituted withMilli-Q water and H2SO4 (0.1 M in purified water) to theoriginal sample dilution and pH.

Irradiations were conducted using a solar simulator(Spectral Energy) (35) with a sample chamber maintained at20 ( 0.2 °C by a Neslab RTE 211 water bath. A total of 8.0mL of the lyophilized and reconstituted sample was keptwithin a 2.0-cm airtight quartz cuvette (Spectrocell) andstirred continuously. The light beam of the solar simulatorwas unfocused to illuminate the entire solution. At measuredtime intervals, aliquots of sample were taken from the quartzcuvette and analyzed for NH4

+. With every illuminated sample∼3 mL of identical sample was placed in a stirred 1.0-cmquartz cuvette kept in the dark at 20 °C to monitor changesin [NH4

+] in the dark.2.2.2. Unaltered Fog Samples. For this set of experiments,

unaltered fog samples with relatively low initial NH4+ and

NO3- concentrations were tested using the conditions

described above. At measured time intervals, aliquots ofilluminated and dark samples were removed from the quartzcuvette; diluted 10-20 times with Milli-Q water (measuredgravimetrically); and analyzed for NH4

+, NO3-, NO2

- and, intwo samples, DON.

2.3. Formation of DIN and NOx during Exposure toSimulated Sunlight and O3. The apparatus for this set ofexperiments is illustrated in Figure 1. Unaltered fog samples(∼13 mL) were irradiated with simulated sunlight (see section2.2) in a 5-cm quartz cuvette at 20 °C while O3 (78-93 ppbv)was bubbled through the cuvette at a flow rate of 0.3 L min-1.O3 was generated by passing O2 (>99.6%, Matheson) throughone of two quartz tubes (GE 021, 5 mm i.d., 180 mm length)illuminated with an ozone-generating Mercury Analamp (620mm length, BHK, 81-1127-01). O3 from the “low-concentra-tion” tube (∼300 ppbv) was diluted with purified air (Aadco737-R) to a mixing ratio of ∼100 ppbv, saturated with watervapor (by passing through a water bubbler at 20 °C), and fed

into the quartz cuvette (Figure 1). O3 from the “high-concentration” tube (∼1 ppmv) was used to oxidize NO toNO2 (see below).

The air/ozone flow exiting the quartz cuvette was firstfiltered (5-µm unlaminated Teflon, Pall-Gelman) to removeany particles or drops generated during bubbling and thenpassed through a denuder coated with citric acid to collectNH3(g) and then a denuder coated with Na2CO3 to collectHNO2(g) and HNO3(g). Downstream of these two denuderswas a 1-L Pyrex reaction flask where NO was oxidized to NO2

(and more oxidized forms) by mixing with ∼1 ppmv O3 at aflow rate of 0.5 L min-1. The reaction flask was kept at ambienttemperature and pressure and was covered in aluminumfoil to keep it dark. NO2 and other oxidized forms werecollected downstream of the reaction flask by two denuderscoated with a solution of 10% guaiacol and 5% NaOH inmethanol (36).

At measured time intervals, aliquots of the fog samplewere taken from the quartz cuvette; diluted 10-50 timeswith Milli-Q; and measured for NH4

+, NO3-, NO2

-, and DON.At the same time, the exposed denuders and reaction flaskwere replaced by a new set and were each extracted with 2.0mL of Milli-Q water. The aqueous extract from the citricacid-coated denuder was analyzed for NH4

+; extracts fromthe other components were analyzed for NO3

- and NO2-.

Concentrations of NH4+, NO2

-, and NO3- formed in the

fogwaters were calculated by combining the amounts in thequartz cuvette with those collected on the citric acid-coatedand Na2CO3-coated denuder, respectively. Concentrationsof NOx were calculated from the sum of NO3

- and NO2-

measured on the two guaiacol-coated denuders and in thereaction flask. Controls were run in the same way as thosein section 2.2 (i.e., without O3 bubbling through the sample).

To determine the relative amounts of NO and NO2

produced, in two experiments we modified the experimentalsetup (Figure 1) so that one of the two guaiacol-coateddenuders was placed before the reaction flask. Under thisconfiguration, the NO2 concentration was calculated as thesum of NO3

- and NO2- measured on the first guaiacol-coated

denuder (which should have collected NO2 and not NO).The NO concentration was calculated from the sum of nitrateand nitrite measured in the reaction flask and on the secondguaiacol-coated denuder (after conversion of NO to NO2,NO2

-, and NO3- in the reaction flask).

As a check on background concentrations, two completesets of freshly coated but unexposed denuders and a cleanreaction flask were extracted and analyzed for inorganic Nspecies. The average concentrations of NH4

+ and NO3-

detected on the citric acid- and Na2CO3-coated denuders,respectively, were less than 10% of the corresponding lowestsample values. N species on the other components were lessthan detection limits. We also performed one proceduralblank where Milli-Q water was used in place of a fog sampleand was purged with 70 ppb O3 and illuminated withsimulated sunlight for 2 h (the typical exposure time for agiven set of denuders and reaction flask during a kineticrun). Concentrations of NH4

+ in the quartz cuvette and onthe citric acid-coated denuders were <5% of the averagevalues measured in illuminated fog samples; concentrationsof NO3

- on the Na2CO3-coated denuder and the reactionflask were <10% of the average sample value. N species onthe other components were below detection limits. Theseblank levels of each species were subtracted from all samplevalues obtained after t ) 0.

2.4. Calculations and Tests of Statistical Significance.Initial reaction rates (Ri, where i refers to a given nitrogenspecies) were determined from plots of concentration versusillumination time using an appropriate fit performed withSigma Plot 7.0 (SPSS Inc). Values of Ri for NH4

+ and NO3-

formation were determined by fitting a three-parameter,

FIGURE 1. Equipment used to measure the photoformation ofinorganic nitrogen in experiments with ozone purging. Chemicalformulas (e.g., NH3 and NOx) indicate the species collected on eachcomponent.

VOL. 37, NO. 16, 2003 / ENVIRONMENTAL SCIENCE & TECHNOLOGY 9 3523

exponential rise to a maximum equation to the experimentaldata:

where [i] is the concentration of the species at time t, and[i]0, a, and b are parameters defined by the regression fit.Conceptually, the variable [i]0 is the initial concentration, ais the product of concentration of the parent DON compoundtimes the yield for formation of i, and b is the first-order rateconstant for formation of i. The initial rate (i.e., at t ) 0) ofNH4

+ or NO3- formation was calculated as

RNOx (i.e., the initial rate of NOx formation) was determinedusing the same equations except that the intercept was setto zero in eq 1.

Destruction rates of NO2- during illumination were

determined using a two-parameter, exponential decay fit:

where [NO2-]0 was a fitted parameter and jNO2

- was the fittedfirst-order rate constant for NO2

- destruction. The initialdestruction rate of nitrite was then calculated from

For DON loss, a single-exponential decay regression generallygave a poor fit, and so we used a double-exponential, four-parameter fit:

where, conceptually, [DON1]0 and [DON2]0 are the sizes (µM)of the two DON pools and jDON1 and jDON2 are the rate constants(min-1) for their destruction. The initial rate of DONdestruction was calculated using

For species where plots of concentration versus time werelinear (or where only two data points were available), thereaction rate was determined as the slope of a linearregression to the data. Linear fits were used for most of thedark control data.

Actinic flux values in each experiment were determinedby measuring the rate constant for loss of 2-nitrobenzalde-hyde (2NB) (j2NB,EXP) at the end of a given experiment (37,38). Cells were not air-purged during actinometry measure-ments since purging had a negligible effect on 2NB photolysisrates. Values of j2NB,EXP measured during this study were ∼2-6times higher than the rate constant for 2NB loss undermidday, winter-solstice sunlight in Davis (j2NB,WIN ) 0.00697s-1; 19). However, we did not adjust reported rates of Ntransformation (e.g., Ri) for differences in actinometry sinceour initial evidence indicates that the photochemical forma-tion/destruction rates for N species are not linearly pro-portional to j2NB,EXP. In the one fogwater tested, increasingthe actinic flux (i.e., j2NB,EXP) by a factor of ∼4 only increasedthe destruction rates of DON and NO2

- by factors of ∼2.0and 2.7, respectively, and increased the formation rates ofNH4

+, NO3-, and NOx by factors of 1.3, 3.3, and 1.4,

respectively. Thus our N transformation rates might be over-reported (relative to winter-solstice sunlight) in samples withhigh j2NB,EXP values. However, as discussed in sections 3.1and 3.3, there are also a few countervailing factors that could

have caused us to under-report rates of N transformation.On the basis of our calculations, light screening by thefogwater samples (i.e., the inner filter effect; 39) was relativelysmall (<13%); therefore, our reported reaction rates werenot increased to account for this effect.

To distinguish between actual changes in the concentra-tion of a given N species during an experiment and changesdue to analytical variations, we performed a student t-testwith Microsoft Excel assuming a two-tailed distribution tocompare the mean values of each pair of concentrations(i.e., before and after illumination). We set the error level at10% (i.e., p values <0.10 were considered significant whilethose >0.10 were considered insignificant). Changes that werenot statistically significant are flagged as such in the tablesof results. The p values for changes in the concentrations ofNH4

+, NO3-, and NO2

- were usually calculated based onstandard deviations of replicate injections. When no replicatemeasurements were performed (e.g., due to limited samplevolumes), p values were calculated using the average relativestandard deviation for each N species from our experimentaldata (1.0% for NH4

+ and NO3- and 2.3% for NO2

-). The pvalues for changes in the concentrations of NOx, DON, andtotal N (which were calculated from a linear combination ofindependent measurements of NH4

+, NO3-, and NO2

-) werecalculated using propagated standard deviations from theindependent DIN measurements (40).

3. Results and Discussions3.1. Photoformation of NH4

+ in Lyophilized Fog and AerosolSamples. We first examined the formation of NH4

+ in 7 fogsamples and 3 aerosol extracts that were lyophilized to reducethe high initial concentrations of NH4

+ so that we could moresensitively determine changes in NH4

+ during experiments(section 2.2). Lyophilization successfully reduced NH4

+

concentrations in these samples to approximately 1% of theoriginal values (Table 1). The photoformation of NH4

+

occurred in all 10 illuminated samples (Table 1). An averageof 8.9 µM NH4

+ (range ) 1.2-21 µM) was formed in thefogwaters during an average illumination time of 520 min,while 0.59-1.1 µM NH4

+ (average ) 0.83 µM) was formed inthe aerosol extracts during 240 min of illumination (Table 1).All of these changes were significant at a 90% significancelevel. The photochemical and thermal (dark) formation ofNH4

+ in the lowest, median, and most reactive fog samplesand in all three aerosol extracts are shown in Figure 2. Theaverage ((1σ) initial formation rates in the illuminated fogsamples and aerosol extracts (RNH4

+,hν) were 3.0 ( 4.0 and 0.4( 0.1 µM h-1, respectively (Table 1). In part, rates in theparticle extracts were lower than in the fogwaters becausethe extracts were much more dilute than the fogwatersamples: the average ((1σ) concentrations of ON prior tolyophilization were 522 ( 480 and 94 ( 43 µM in fogwaterand aerosol extracts, respectively; the average [NH4

+] priorto lyophilization in the fogwaters was ∼7 times greater thanthat in the particle extracts (8, 9). Converting the solutionvolume-based rates of NH4

+ formation to air volume-basedrates (using measured fog liquid water contents, volumes ofair sampled, and volumes of aerosol extract solution) yieldsvalues of 4.0 ( 5.6 and 1.9 ( 0.95 ng of NH4

+ (m of air)-3 h-1

in fogwaters and PM2.5, respectively (Table 1).The fact that the average air volume-based formation rate

of NH4+ in the fogwaters was ∼2 times larger than that in the

PM2.5 suggests that the fogwaters contained greater amountsof photochemically labile DON precursor for NH4

+. This mightbe because (i) the fog samples contained DON from bothfine and coarse (>2.5 µm) particles (since both can act ascondensation nuclei for fog drop formation) and/or (ii) mostof the fogwaters were collected during the night (when theDON is protected from photochemical degradation) whilethe aerosol samples were typically collected over a continuous

[i] ) [i]0 + a(1 - e-bt) (1)

Ri ) ab (2)

[NO2-] ) [NO2

-]0 exp(- jNO2-t) (3)

RNO2- ) [NO2

-]0 jNO2- (4)

[DON] )[DON1]0 exp(- jDON1

t) + [DON2]0 exp(- jDON2t) (5)

RDON ) [DON1]0 jDON1+ [DON2]0 jDON2

(6)

3524 9 ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 37, NO. 16, 2003

2-day period (8, 9). A strong correlation (R 2 ) 0.83) betweenRNH4

+,hν (ng (m of air)-3 h-1) and the concentration (ng of N(m of air)-3) of combined amino N in these fog and PMsamples (8, 9, 22) suggests that combined amino compoundssuch as proteins and peptides might be an importantprecursor of photochemically formed NH4

+. In contrast,RNH4

+,hν was only very weakly correlated with initial concen-trations of NH4

+, NO2-, NO3

-, total organic N, and free aminoN (R 2 ) 0.13, 0.01, 0.07, 0.01, and 0.01, respectively).

The formation of NH4+ was observed in the dark in both

fogwaters and PM2.5 extracts, indicating that NH4+ is formed

thermally as well as photochemically (Figure 2). However,the thermal (i.e., dark) reaction rate (RNH4

+,dark) was only ∼10%of the corresponding rate in illuminated samples (Table 1).In addition, samples with higher photoformation rates ofNH4

+ (RNH4+,hν) typically had higher dark rates (RNH4

+,dark),although the correlation is relatively weak (R 2 ) 0.51).

It should be noted that the measured rates described above(and in subsequent sections) might overestimate or under-estimate actual rates of inorganic N formation in ambientfog and PM2.5. There are several reasons for this. First, whileour actinometry indicates that the actinic flux in this set ofexperiments was approximately 2.5 times higher than thatat midday on the winter solstice in Davis, initial data indicatethat rates of NH4

+ are only weakly dependent on actinic flux(section 2.4). In any case, this difference in actinic flux wouldbe largely eliminated if we were to take into accountnumerical modeling results, which indicate that the actinicflux in an aqueous drop is ∼2 times greater than in thesurrounding gas phase (41). Second, although our sampleswere stored frozen between collection and the experiments(a period of ∼1-2 yr), past studies have shown that ON candecay during this storage (8). This loss of labile ON likelywould cause our measured rates of IN formation to under-

TABLE 1. Formation of NH4+ in Lyophilized Fog Samples and Aerosol Extracts

illumination conditions formation ratesc

sample no.a

[NH4+] prior to

lyophilization(µM)

[NH4+] after

lyophilization(µM)

t(min)

j2NB(s-1)

∆[NH4+]b

(µM)RNH4

+,hν(µM h-1) RNH4

+,dark

RNH4+,hν

(ng m-3 h-1)d

Fog SamplesDA97-05F 473 1.2 600 0.019 5.2 0.9 0.4 1.9DA98-09F 1287 3.1 306 0.023 6.4 3.2 0.1 4.7DA98-13F 719 1.9 360 0.020 1.2 0.2 0.0 0.27DA99-01F 514 1.5 690 0.023 3.0 0.5 0.1 0.88DA99-04F 3151 24 660 0.022 15.4 2.7 1.0 2.8DA99-05F 1437 32 300 0.023 21 12 0.9 16DA99-07F 1128 5.5 715 0.023 10 1.4 0.3 1.0mean ( 1σ 1244 ( 921 9.9 ( 13 519 ( 188 0.022 ( 0.002 8.9 ( 7.1 3.0 ( 4.0 0.4 ( 0.4 4.0 ( 5.6

Aerosol ExtractsADA97-49/50e 87 3.1 240 0.024 0.59 0.4 0.02 0.91ADA97-57/58f 200 1.9 240 0.024 1.1 0.5 0.03 2.8ADA98-09 249 2.3 240 0.024 0.86 0.4 0.04 2.1mean ( 1σ 178 ( 83 2.5 ( 0.8 240 ( 0 0.024 ( 0.000 0.8 ( 0.2 0.4 ( 0.1 0.03 ( 0.01 1.9 ( 0.95

a Information on sampling dates and times are given in refs 8 and 9. pH values of the samples ranged from 5.7 to 7.0 for the fogwaters andfrom 5.6 to 6.3 for the aerosol extracts. b Changes in NH4

+ concentration during illumination. All changes were significant at p < 0.10 (see section2.4). c RNH4

+,hν ) NH4+ formation rate in illuminated cell; RNH4

+,dark ) NH4+ formation rate in dark cell. d Formation rate of NH4

+ in the ambient fogor aerosol, calculated using measured aqueous formation rates and ambient sample information (liquid water content for fog or the volumes ofair sampled and aqueous extraction volumes for aerosol samples). e Mixture of equal amounts of extracts from ADA97-49 and ADA97-50. f Mixtureof equal amounts of extracts from ADA97-57 and ADA97-58.

FIGURE 2. NH4+ formation in lyophilized fog samples and fine particle extracts illuminated with simulated sunlight (plots a and c) and

kept in the dark (b and d). Samples: 1, DA99-05F; 2, DA98-09F; 3, DA98-13F; 4, ADA97-57 + ADA97-58; 5, ADA98-09; and 6, ADA97-49 +ADA97-50. Symbols with error bars represent the average value ((1σ) from two replicate injections. Lines represent the regression fitsto each set of experimental data.

VOL. 37, NO. 16, 2003 / ENVIRONMENTAL SCIENCE & TECHNOLOGY 9 3525

estimate the actual rates in the atmosphere. The loss of ON(and accompanying underestimation of IN formation rates)appears to have been somewhat exacerbated by lyophilizationon the basis of the one sample tested where lyophilizationreduced DON concentrations by 30% but did not significantlyaffect levels of NO3

-. Finally, filtration of our samples mighthave also reduced sample reactivity since our filtered samplestypically have ∼30-40% less absorbance (λ ) 290-500 nm)as compared to unfiltered samples. However, a loweredsample absorbance might not necessarily lead to a reducedsample reactivity, since according to a recent review (33),there is no correlation between light absorptivity and thephotoformation rate of ammonium in aquatic samples.

3.2. DIN and DON Dynamics during Simulated SunlightIllumination in Unaltered Fog Samples. Although thelyophilization process allowed us to more sensitively measureNH4

+ formation rates, it also apparently reduced concentra-tions of ON as described above. To avoid this possible bias,we examined the photochemical reactions of N species infive fogwater samples that were unaltered (i.e., withoutlyophilization). To enhance our ability to measure DINformation, we chose samples with relatively low initialconcentrations of inorganic N (on average ∼5 times lowerthan the typical value in Davis fog samples; 8). Three of thefive unaltered samples showed a statistically significantincrease in [NH4

+] (p < 0.10), while the change in one samplewas significant at a slightly larger value (p < 0.11, Table 2).On average 15 µM NH4

+ was produced during ∼490 min ofirradiation (Table 2) and initial rates ranged from 1.8 to 8.4µM NH4

+ h-1 (average (1σ ) 5.7 ( 2.7; Table 2). This valuewas roughly two times higher than the average initial rate ofNH4

+ formation in the lyophilized samples (Table 1), sug-gesting that lyophilization significantly reduced samplereactivity.

NO3- was formed in all five of the illuminated samples;

in four samples the increase was statistically significant atp < 0.10, while in the fifth was nearly so (p < 0.11; Table 2).The initial rates of NO3

- formation were very similar to thoseof NH4

+, with an average ((1σ) value of 6.3 ( 2.5 µM h-1

(range ) 3.7-9.9 µM h-1; Table 2). In contrast to NH4+ and

NO3-, we observed a net loss of NO2

- in these fogwatersduring illumination (average (1σ ) -1.1 ( 0.7 µM h-1, Table2). The average lifetime of NO2

- (τΝÃ2-) during these experi-

ments was 15 ( 6.7 h (range ) 6.8-22 h; Table 2), which isconsiderably larger than the lifetime of nitrite measured inMilli-Q water (τΝÃ2

- ) 5.8 h) under similar conditions (pH5.6; j2NB,EXP ) 0.022 s-1). Since direct photolysis accountedfor ∼90% of NO2

- loss in the illuminated Milli-Q and sincethe inner-filter effect in the fogwaters was negligible (section2.4), the slower destruction rates of NO2

- in the fog samplesindicate that NO2

- was photochemically formed in thefogwaters in addition to being photochemically destroyed.The net loss of NO2

- in the illuminated fog samples indicatesthat the photochemical destruction of nitrite was more rapidthan its photochemical formation.

Although all five fog samples studied had significantamounts of DON (average of 238 ( 70 µM in the originalsamples; 8), due to limited sample volumes we followed DONloss in only two samples (DA98-13F and DA99-02F). Bothsamples lost more than half of their initial DON over thecourse of ∼10 h of illumination (p < 0.10; Table 2). In bothsamples a single exponential regression gave a poor fit to theloss of DON during illumination but a double-exponentialfit (eq 5) worked well (e.g., Figure 3d). This suggests that theDON compounds in fogwaters fall broadly into two pools:one that is relatively rapidly mineralized and another that ismore resistant to mineralization. On the basis of the curve-fitting parameters determined for these DON decays (Table2), the sizes of the more reactive DON pools (i.e., [DON]1)were up to 50% smaller than those of the less reactive pools TA

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)R N

O3-

,hν

(µM

h-1 )

[NO

2-] 0

(µM

)∆

[NO

2-]

(µM

)R Ν

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,hν

(µM

h-1 )

τ NO

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(h)

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(µM

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[DO

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(µM

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(1σ

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(3.

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itia

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sect

ion

2.4)

.Ri,h

ν,fo

rmat

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or

des

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ated

cell.

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sin

gd

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ther

no

tmea

sure

d(f

or

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N)o

rre

pre

sen

ted

con

cen

trat

ion

chan

ges

that

wer

en

ots

tati

stic

ally

sig

nif

ican

t(fo

rN

H4+

and

NO

2-).

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ed

ark

ther

ew

ere

no

det

ecta

ble

chan

ges

inD

ON

or

NO

2-,w

hile

NO

3-ch

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edin

on

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ne

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ple

(DA

99-0

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ark

)0.

67µM

h-

1 ),b

ut

this

valu

ew

asst

atis

tica

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sig

nif

ican

t(p

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19).

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plin

gd

ates

and

tim

esar

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iven

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fs8

and

9.p

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lues

of

the

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ple

sra

ng

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om

6.6

to7.

5.b

Val

ues

of

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ob

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ga

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eter

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fit

toth

ed

ata.

RD

ON

)[D

ON

1]0×

j DO

N1

+[D

ON

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j DO

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(eq

5).V

alu

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j DO

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0,an

dj D

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nd

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×10

-4

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rD

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and

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was

stat

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gn

ific

ant

atp

<0.

10:

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DA

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(p)

0.11

).

3526 9 ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 37, NO. 16, 2003

(i.e., [DON]2), while the lifetime of the more reactive pool(1/jDON1 ∼ 1 h) was roughly 30 times shorter (1/jDON2 ∼ 33 h).The net result was a quick depletion of the “more reactive”ON compounds followed by a slower photodestruction ofthe remaining DON (e.g., Figure 3d). Initial destruction ratesof DON in DA98-13F and DA99-02F were 52 and 56 µM h-1,respectively, much higher than the sum of the initialformation rates of NO3

- and NH4+ in these samples (Table

2). It should also be noted that even those DON compoundsthat were not mineralized likely underwent significantchemical transformations as a result of direct and/or indirectphotooxidation reactions to form products such as alkylnitrates (42) and a suite of other compounds.

The average ((1σ) air volume-based initial formation ratesfor NH4

+ and NO3- in the unaltered fog samples were 12 (

8.4 ng m-3 h-1 (range ) 3.3-23) and 42 ( 12 ng m-3 h-1

(range ) 25-54), respectively. The initial destruction ratefor NO2

- was 6.4 ( 5.6 ng m-3 h-1 (range ) 3.1-15). Assumingan average molecular mass of DON of 100 Da per N atom(9), we calculate that the destruction rates of DON in DA98-16F and DA99-02F were 490 and 610 ng m-3 h-1, respectively.Although these samples were not lyophilized, as discussedat the end of section 3.1 there are still a number ofuncertainties that prevent us from quantitatively using our

experimentally measured values to determine actual reactionrates in ambient Davis fogwaters. In addition to the uncer-tainties described earlier, given that concentrations ofinorganic N and DON are positively correlated in fogwaters(8) and that high DON levels suggest a larger reservoir ofNH4

+ and NO3- precursors, our results in Table 2 might be

biased to lower rates since we specifically chose sampleswith low initial concentrations of DIN.

The loss of DON and formation of DIN in one fog sample(DA99-02F) are plotted in Figure 3. As seen in the other fogsamples, the formation rates of NH4

+ and NO3- in DA99-02F

were similar, while the NO2- destruction rate was smaller

and the DON destruction rate was much faster. In the darkcell, none of the species showed concentration changesdistinguishable from analytical variations (i.e., p > 0.10;section 2.4). As shown in the bottom panel of Figure 3, theamount of total N in the quartz cuvette decreased by ∼10 µMduring illumination, that is , the amount of DON lost was∼10 µM greater than the sum of NH4

+, NO3-, and NO2

- formed(p < 0.10). In contrast, there was no statistically significantchange of total N in the dark (p > 0.10). These results suggestthat approximately 10 µM DON was converted into one ormore volatile N species that escaped from the liquid phaseduring illumination and were not accounted for. This“missing” N species likely was produced from the morereactive DON pool since it was apparently produced duringthe initial stage of illumination (Figure 3d,e). As describedin section 3.3, it appears that NOx was this missing inorganicN species.

To examine the effects of solution pH on DON miner-alization, we characterized N photochemistry in DA99-02Fat its original pH (6.4) and at pH 3.2 after adjustment withH2SO4. Acidification did not significantly influence the rateof NH4

+ formation (Figure 4a), but it did accelerate the lossof N(III) (i.e., NO2

- and HNO2) and the formation of NO3-

(Figures 4b,c). The rates for N(III) destruction and NO3-

formation were 20 and 6 times faster, respectively, at pH 3.2than at pH 6.4. The more rapid transformation of N(III) atthe lower pH was likely caused by the increase in theabundance of HNO2 at lower pH (pKa ) 3.25; 43), which ismuch more photochemically reactive than its conjugate base(NO2

-) in sunlight (44). The faster NO3- formation at lower

pH was likely due to a more rapid transformation of N(III)rather than a faster DON mineralization, as suggested by thefact that the sum of N(III) and NO3

- was almost the same inthe different pH solutions throughout the course of il-lumination (Figure 4b). This experiment indicates that thephotochemical conversion of dissolved organic N speciesinto NO3

- (and NH4+) was independent of pH and that a

portion of the photoproduced NO3- resulted from N(III)

photooxidation, while the remainder came from DONmineralization.

3.3. Photoformation of DIN and NOx in Unaltered FogSamples Exposed to Simulated Sunlight and O3. In an effortto determine whether NOx was the missing product formedfrom DON destruction in DA99-02F, we modified ourexperimental apparatus to be able to collect NOx (section 2.3and Figure 1) and exposed the sample to simulated sunlightduring purging with 93 ppbv ozone. During 420 min ofexposure, 8.9 µM NO3

- and 20 µM NH4+ were formed while

3.9 µM NO2- and 33 µM DON were destroyed (Figure 5).

Calculated initial rates of NO3- and NH4

+ formation in thisexperiment were 7.4 and 11 µM h-1, respectively, while thatof NO2

- destruction was 1.3 µM h-1 (τΝÃ2- ) 12 h). These

values were ∼2-2.5 times higher than those measured inthe same sample during simulated sunlight illuminationwithout ozone (Figure 3, Table 2). Surprisingly, the initialrate of DON loss in the presence of O3 was slower than inthe experiment without ozone (Figures 3d and 5d). Thekinetics of the DON decay were also different in the two

FIGURE 3. Concentration changes of nitrogen species in an unalteredfogwater (DA99-02F) during simulated sunlight illumination (opensymbols) and in the dark (solid symbols). The value of j2NB,EXP/j2NB,WIN

for this experiment was 2.5. Each data point represents an averageof two replicate injections, with error bars representing (1σ. Solidlines represent regression fits to the data (plots a-d). In plot e,[Total N] is the sum of the concentrations of NH4

+, NO3-, NO2

-, andDON in the fogwater in the quartz cuvette. The dashed line in plote is the average of [Total N] in the dark and light solutions att ) 0.

VOL. 37, NO. 16, 2003 / ENVIRONMENTAL SCIENCE & TECHNOLOGY 9 3527

experiments: in the presence of light only (no ozone) theDON decay was well described by a double-exponential fit(Figure 3d), while the combination of light + ozone gave aDON decay that was well-described by a single-exponentialregression (Figure 5d). These differences in rates and kineticsmight be due to a change in the DON transformationmechanism, possibly due to the presence of ozone, and/orbecause of the higher actinic flux employed in the secondexperiment (j2NB,EXP ) 0.044 s-1 as compared to 0.018 s-1 inthe previous one). In addition, although the initial DONconcentrations in DA99-02F were nearly the same in the twoexperiments (Figures 5d and 3d), the bulk reactivity of DONmight also have changed because the second experimentwas performed ∼1 yr after the first (during which time thesample was stored frozen).

A total of 7.5 µM NOx (expressed as an equivalent aqueousconcentration in the fogwater sample) was formed in DA99-02F during the 420 min of illumination with O3 exposure(Figure 5e). The initial formation rate of NOx was 3.7 µM h-1,corresponding to an air volume-based formation rate of 10pptv NOx h-1 (liquid water content ) 0.11 g (m of air)-3). NO2

accounted for 31% of the NOx formed during this experiment.The inclusion of NOx in the N budget resulted in a betterbalance of total N (i.e., [NO3

-] + [NO2-] + [NH4

+] + [DON]+ [NOx]) throughout the experiment (Figure 5f).

Extrapolation of this measured NOx formation rate to theambient foggy air is complicated by several factors (inaddition to those discussed earlier): (i) the experimental

actinic flux was higher than that of wintertime sunlight(j2NB,EXP ≈ 6.5 j2NB,WIN); (ii) this fog sample (DA99-02F) mighthave suffered significant losses of reactive ON during storage;and (iii) the mixing ratio of O3 used (93 ppbv) was muchhigher than typical values in the interstitial air of winter fogsat Davis (45). In addition, the characteristic time forevaporation of NOx in the bulk fog solution during theexperiment (τevap ∼ 1 s at a purging rate of 0.3 l min-1) wasmuch longer as compared to that in ambient fog drops (τevap

∼ 0.022 s in a 20 µm diameter drop). One consequence ofthe slow evaporation in our experiments is that a significantportion of the aqueous NOx was likely converted into NO3

-

and other products (46) before it could be purged out ofsolution. We calculate that the dominant sinks for aqueousNOx were reactions with superoxide (•O2

-) and self-reactions(e.g., NO + NO2), both of which eventually produce NO3

-

FIGURE 4. Influence of pH on inorganic nitrogen formation ratesin (unaltered) sample DA99-02F during simulated sunlight illumina-tion (j2NB,EXP/j2NB,WIN ) 2.5). Filled symbols (and solid lines) representresults from the original, unaltered sample (pH 6.4), while opensymbols (and dashed lines) represent results from the sample afteradjustment to pH 3.2. In plot a, the data point at t ) 36 min wasomitted from the regression fit for the pH 3.2 data. In plot b, theinverted triangles represent the changes in the total concentrationof NO3

- and N(III) [i.e., ∆([NO3-] + [N(III)])] at pH 6.4, while the

upright triangles represent changes at pH 3.2. The dotted curve isan exponential regression fit to all of the ∆([NO3

-] + [N(III)]) data.(Note that N(III) ) NO2

- + HNO2.) In plot c, the regression fit to thepH 6.4 data was performed with the data point at t ) 0 min omitted.

FIGURE 5. Changes in nitrogen speciation in (unaltered) fogwaterDA99-02F during exposure to simulated sunlight and purging withO3 (93 ppbv). Each data point for NH4

+, NO3-, and NO2

- representsthe average ((1σ) of two replicate injections. Error bars on eachDON, NOx, and total N data point represent (1σ calculated frompropagated standard deviations of the corresponding NH4

+, NO3-,

and NO2- measurements. Open symbols represent samples exposed

to simulated sunlight and O3, while filled symbols represent samplesin dark, unpurged controls. In plot f, [Total N] is the sum of theconcentrations of NH4

+, NO3-, NO2

-, and DON in the fogwater plusthe IN amounts from the denuders and reaction flask. For thisexperiment j2NB,EXP/j2NB,WIN ) 6.5.

3528 9 ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 37, NO. 16, 2003

(46). Therefore, the upper bound for RNOx would be the sumof the measured formation rates of NOx and NO3

-, that is,11 µM h-1 (or 30 pptv NOx h-1 on an air volume basis) inDA99-02F.

To understand the fraction of photoproduced NOx thatcame from DON reactions, we subtracted the rate of NOx

formation due to the direct photolysis of NO3- and N(III)

using actinometry-normalized rate constants reported pre-viously (38, 44). On the basis of these data and an estimatedevaporation efficiency for NOx of 50%, we determined thatapproximately half of NOx production in DA99-02F duringthe light + ozone exposure was due to NO3

- and N(III)photolysis and that approximately half of the NOx was fromDON photodestruction. It is interesting to note that eventhough nitrate is photochemically reduced to NOx (47, 48),in our samples we generally saw an increase in NO3

- duringillumination (although sometimes we could measure no netchange), indicating that aqueous photochemical reactionswere generally a net source of nitrate.

The formation of NOx was also detected in the other twosamples tested with simulated sunlight and ozone. Initialformation rates during the experiments were 1.5 µM h-1

(DA99-04F, j2NB,EXP/j2NB,WIN ) 6.1, NO2:NOx ) 0.55) and 2.1µM h-1 (DA01-01F, j2NB,EXP/j2NB,WIN ) 5.9), corresponding toinitial rates of NOx release of 2.1 and 5.1 pptv h-1, respectively,in ambient foggy air. Due to high initial concentrations,neither sample had detectable changes in NO3

- and NH4+

during illumination. Unlike the case for DA99-02F, in theseexperiments we calculate that all the NOx collected camefrom photolysis of NO3

- and N(III), probably because theinitial concentrations of these inorganic chromophores werevery high.

3.4. Possible Reaction Mechanisms for Conversion ofOrganic N to Inorganic Forms. Although a very incompleteunderstanding of DON composition hampers our under-standing of the mechanisms for the formation of inorganicN in our experiments, we can make some suggestions forpossible mechanisms. For example, it has been reported thattransformations of amino compounds can produce NH4

+,either through hydrolysis or through reactions with inter-mediates such as hydroxyl radicals (18, 19, 31, 32, 49, 50).Given that amino compounds typically accounted for 20%of the organic N in the samples examined in this study (8,9, 22), it is possible that these compounds were a significantprecursor for the photoproduced NH4

+. The strong correla-tion between the initial formation rates of NH4

+ and theconcentrations of amino N in fogwaters (section 3.1) supportsthis hypothesis. In addition to amino compounds, humicsubstances are probably also present at significant concen-trations in the Davis fogwaters and aerosol samples (37, 51)and could be significant photochemical sources of NH4

+ (30-32).

The chemistry and formation of NO3- is intricately linked

with reactions of NO2- and NOx. For example, NO3

- is botha product from oxidation reactions of NO2

- and NOx as wellas a photochemical precursor of these two species (47, 48).Nonetheless, our results indicate that the mineralization oforganic nitrogen was primarily responsible for the formationof NO3

- and NO2- and could be a significant source of NOx.

One possible source of NOx, NO2-, and NO3

- in our samplesis nitro compounds (e.g., nitrophenols), which can bephotolyzed in particles and water drops to form NO and NO2

that can be further oxidized into NO2- and NO3

- (20, 21, 52,53). In addition, Chang and Novakov (54) have suggestedthat atmospheric ON can be formed by reactions of gaseousNH3 and NOx with organics or soot; if these reactions arereversed in the presence of light they might form NH4

+,NO3

-and NOx.3.5. Environmental Implications. Our results indicate

that organic nitrogen compounds in atmospheric condensed

phases are converted into NH4+, NO3

-, NO2- and NOx during

exposure to sunlight and O3. This transformation of ON toinorganic N is a previously unrecognized source of theseinorganic nitrogen species in the troposphere. Although theformation rates of ammonium and nitrate were relativelysmall as compared to their initial concentrations, the absoluterates of formation were significant (i.e., up to nearly 10 µMh-1). As we note in several previous sections, quantitativelyextrapolating from these laboratory-derived rates to rates inthe ambient atmosphere is not yet possible. Furthermore,quantitatively describing or modeling the influence of organicnitrogen transformations as a source of inorganic N in theatmosphere will require more extensive measurements ofON in tropospheric particles and drops.

Despite these uncertainties, our results indicate that thephotochemical conversion of atmospheric organic N intoinorganic N is significant and could have important impacts.For example, given the widespread occurrence and abun-dance of tropospheric ON (1, 5, 7-9, 11, 12) and the fact thatinorganic N is typically much more bioavailable than organicN (55), these transformations might significantly increasethe bioavailability of N in deposition. In addition, these samereactions likely occur after deposition (e.g., in surface waters),thereby further increasing the ecological impacts of depositedatmospheric organic N. Such effects might be relatively moreimportant in remote areas where ON accounts for a largerfraction of the total N in the atmosphere (4, 12).

The conversion of ON into inorganic N species likely alsoinfluences the physical, chemical, and toxicological propertiesof atmospheric condensed phases. For example, due to thehigh water solubility of NH4

+ and NO3-, formation of these

species from organic precursors likely alters the hygroscop-icity and surface tension of atmospheric particles and waterdrops. Since NOx is a key participant in the photochemicalreactions that form ozone and thereby alters the oxidationcapacity of the atmosphere, the photochemical formation ofNOx from condensed-phase organic N is especially interestingin terms of its potential influence on tropospheric chemistry.In addition, the observed rapid loss of organic nitrogencompounds during exposure to light (or light + ozone) mightprofoundly influence the many important properties ofatmospheric condensed phases that are affected by organiccompounds (e.g., light absorption, acid-buffering capacity,and complexation/speciation of transition metals). Thesesame properties will also be altered by photochemicallyinduced changes that occur in the organic N pool that remainsafter exposure. Although this ON is not mineralized, it likelyundergoes alterations in composition, for example, to formmore oxidized forms of ON, including alkyl nitrates (42).

AcknowledgmentsThe authors thank Ingrid George and Jeff Chan for help withexperiments and Keith McGregor and Mike Jimenez-Cruzfor assistance with fog and aerosol sampling. This work wassupported by the U.S. EPA (R819658 and R825433) Centerfor Ecological Health Research at the University of Californiaat Davis. Although the information in this document hasbeen funded in part by the United States EnvironmentalProtection Agency, it may not necessarily reflect the viewsof the Agency, and no official endorsement should be inferred.Additional funding for this work was provided by a graduatefellowship from the University of California Toxic SubstancesResearch and Teaching Program (ecotoxicology component),by the California Agricultural Experiment Station (ProjectCA-D*-LAW-6403-RR), and by a Jastro-Shields GraduateResearch Award from the University of California at Davis.

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Received for review February 7, 2003. Revised manuscriptreceived May 28, 2003. Accepted June 3, 2003.

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