Effect of Organic Coatings on Gas-Phase Nitrogen Dioxide Production

Article pubs.acs.org/JPCC

Effect of Organic Coatings on Gas-Phase Nitrogen Dioxide Production from Aqueous Nitrate Photolysis Dorea I. Reeser,† Nana-Owusua A. Kwamena,† and D. J. Donaldson*,†,‡ †

Department of Chemistry, University of Toronto, 80 St. George St., Toronto, Ontario Canada M5S 3H6 Department of Physical and Environmental Sciences, University of Toronto at Scarborough, 1265 Military Trail, Toronto, Ontario Canada M1C 1A4

‡

ABSTRACT: The influence of stearic acid, octanol, and octanoic acid monolayer coatings on the release of NO2 into the gas phase following aqueous NO3− photolysis was studied using incoherent broadband cavity-enhanced absorption spectroscopy (IBBC-EAS). The different organic compounds, when present at the aqueous surface, had varying effects on the gas-phase NO2 evolved. Stearic acid monolayers lowered the initial rate of appearance of NO2(g), and its steady-state concentration was the same as for uncoated solutions after ∼50 min. In the presence of octanol monolayers, both the steady-state [NO2(g)] and its rate of appearance decreased. A simple kinetic phase partitioning model suggests that the rate of NO2(g) evaporation from the aqueous surface is physically inhibited by the long uncompressed stearic acid chains, whereas both NO2 evaporation and steady-state NO2(g) concentration decrease when octanol is present at the aqueous surface, due to the enhanced solubility of NO2 in the less polar octanol environment. Despite its structural similarity to octanol, monolayers of octanoic acid showed a different effect and slightly increased the steady-state [NO2(g)]. We propose that octanoic acid enhances NO2(g) production because of an increase in solution acidity, which increases the quantum yield of NO2 production from nitrate photolysis.

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NO3− + hν → NO2 + O−

INTRODUCTION

Nitrogen dioxide is an important tropospheric pollutant since it contributes to photochemical smog and the formation of acid rain, and it influences the oxidation capacity of the troposphere via HOx/NOx cycles, where NOx = NO2 + NO and HOx = HO2 + OH.1 Furthermore, it is a respiratory irritant and reduces plant growth.2−4 Nitrogen dioxide in the troposphere is primarily produced via the reaction between nitric oxide and ozone or peroxy radicals. Combustion is the major source of NOx, and additional sources include lightning, microbial nitrification/denitrification in soils, oxidation of ammonia, and aqueous nitrate photolysis.1 There are several sources of nitrate in natural waters including the oxidation of ammonia to nitrite with subsequent oxidation to nitrate, nitrifying bacteria and wash off from wastewater, sodium nitrate, and ammonium nitrate fertilizers. In addition to the natural and anthropogenic sources of nitrate in natural waters, it is also produced as ammonium nitrate and nitric acid in the troposphere. The formation of nitric acid is a major sink for NOx species in the troposphere, and gas-phase HNO3 is easily sequestered by aerosols or other wet surfaces, forming aqueous nitrate.1,5 Aqueous HNO3 is also produced in the heterogeneous reaction between gas-phase dinitrogen pentoxide and water surfaces.1,6−8 The photolysis of aqueous nitrate can lead to the release of NOx, with yields which depend on the wavelength of light9−12 © 2013 American Chemical Society

−

→ NO2 + O −

→ ONOO

(1) (2) (3)

There are two primary pathways of nitrate photolysis due to two absorption bands of nitrate; a weak n → π* band at ∼305 nm (eqs 1 and 2) and a strong π → π* band at ∼200 nm eqs 1−3. Room-temperature nitrate solutions with a pH between 4 and 11 illuminated with ∼305 nm light show quantum yields of 0.01 and 0.001 for OH and O atom production, respectively,9−11 with OH from the reaction O− + H2O ⇌ OH + OH−. Similar solutions illuminated at ∼250 nm give quantum yields of 0.09 for OH and 0.1 for ONOO−.13,14 Understanding the chemical and physical processes that influence the evolution of NOx from nitrate photolysis is of interest because its aqueous deposition as nitric acid is considered to be a major loss process of tropospheric inorganic nitrogen due to the small quantum yield from nitrate photolysis under actinic radiation (λ ≥ 290 nm). Nitrate photolysis may occur in salt and fresh waters, in aqueous aerosols released from wave breaking of these waters, Special Issue: Ron Naaman Festschrift Received: February 12, 2013 Revised: April 4, 2013 Published: April 4, 2013 22260

dx.doi.org/10.1021/jp401545k | J. Phys. Chem. C 2013, 117, 22260−22267

The Journal of Physical Chemistry C

Article

free Xe arc lamp reflected by a mirror through a quartz window at the top of the chamber so that the lamp illuminated the solutions from above. Samples were exposed to a continuous flow of 0.3 LPM dry N2 passing through the reaction chamber and into the stainless steel IBBC-EAS cell. The gas-phase concentration of NO2, [NO2(g)], was measured using incoherent broadband cavity-enhanced absorption spectroscopy (IBBC-EAS); a diagram of the apparatus is shown in Figure 1.

and in aerosols formed on ammonium nitrate seed particles. The uppermost hundreds of micrometers of natural fresh and salt water surfaces, called the surface microlayer (SML), are enhanced in dissolved organic matter (DOM) relative to the bulk surface and subsurface waters.15 Aerosols are known to be rich in organic matter,1,15,16 and there is recent evidence that this may partition to the aerosol airaqueous interface.17,18 Organic species present at airaqueous interfaces may influence the physical and chemical processes that occur there compared to a pure waterair interface and also the corresponding bulk phases.19,20 Therefore, understanding the influence of DOM present at the airwater interface has been a recent topic of interest.15 The presence of octanol coatings as a proxy for DOM can increase the reaction rates of O3 with bromide and polycyclic aromatic hydrocarbons (PAHs), whereas the presence of a similar species, octanoic acid, slows the reaction between O3 and PAHs.21−23 Another proxy for DOM, humic acid, has been shown to reduce the uptake of N2O5 onto aerosols.24 High surface coverages of short chain fatty acids and alcohols (C ≤ 6) may impede the evaporation of D2O as well as impede or enhance the uptake of HCl onto sulfuric acid depending on the functional groups present at the aqueous surface.25,26 Compressed long chain fatty acid monolayers (C ≥ 17) inhibit the evaporation of water and the uptake of acetic acid, N2O5, and O3 onto the aqueous phase.25,27−32 When surface coverage increases, insoluble organic molecules become more compact at the surface and further impede gas uptake or evaporation, and the degree of this effect depends on the structure of the components making up the monolayer.25,30−32 The present work examines the effect of surface active organic species on the release of NO2 into the gas phase from the photolysis of aqueous NO3− using stearic acid, octanol, and octanoic acid as proxies for DOM present at the aqueous surface.

Figure 1. Experimental setup of the photolysis experiments and IBBCEAS apparatus.

A 365 nm 10 W LED (LedEngin) was focused into the cell (length: 113 cm, diameter: 1 in.) using a 2 in. f/2 BCX lens. Each end of the cell was terminated by a highly reflective mirror (ATF Inc., >99.96% R at 362 nm); under typical conditions this creates a cavity with an effective path length of ∼3.3 km. The mirrors were continuously flushed with 25 sccm dry N2 to maintain cleanliness and thus reflectivity. The transmitted light was focused with a 1 in. f/1 lens onto a fiber optic bundle, which directed it into a 125 mm focal length spectrograph with a 1200 L/mm grating. The dispersed light was detected by a CCD. Wavelength calibration from 270.034 to 450.987 nm was performed using the known lines from a mercury lamp. Transmission spectra were collected for 38 s (either 3 averages with an integration time of 9.5 s or 10 averages with an integration time of 3.4 s) using software that accompanied the CCD. To calculate the concentration of NO2 from the measured spectra, the mirror reflectivity, the Rayleigh scattering crosssection of the carrier gas, and the absorption cross-section of NO2 all must be known as a function of wavelength. We used the approach previously described by Bodhaine et al.,36 Fiedler et al.,37 and Washenfelder et al.38 The reflectivity, R(λ), of the mirrors was determined as a function of wavelength using eq 4 ((A8) in Washenfelder et al.)38

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EXPERIMENTAL SECTION Solutions of 1.0 M NO3− were prepared using KNO3 (ACP, >99%) and 18 MΩ water; dilute HCl (Fisher Scientific) was added to adjust the pH to ∼3.8 for experiments performed under acidic conditions. Octanol (Aldrich, 99%) and octanoic acid (Acros Organics, 99%) monolayers were prepared by adding 7.9 or 9.5 μL of the organic liquid, respectively, to 20 mL of 1 M NO3− solution in an Erlenmeyer flask to make solutions of 2.5 mM octanol and 3 mM octanoic acid, which correspond to monolayer coverages on pure water.22,33 Addition of octanoic acid decreases the bulk pH from ∼5.4 to 3.6. These solutions were shaken gently for at least 2 min to ensure full dispersion of the organics, and then the solution of interest was poured into a shallow Pyrex dish inside the reaction chamber. Stearic acid coatings were prepared in a different manner because stearic acid is insoluble in water. A 0.5 mM stearic acid solution was made by dissolving stearic acid (Acros Organics, >97%) in chloroform (Caledon, ethanol stabilized, spectrograde) and stored in a refrigerator. Using a microliter syringe, 21.7 μL of this stearic acid/chloroform solution was gently added dropwise to the surface of the sample nitrate solution in the shallow Pyrex dish inside the reaction chamber. All solutions with organic species present were allowed to sit in the 1 L Teflon reaction chamber for at least 20 min prior to beginning an experiment to allow the organic monolayer to spread evenly across the aqueous surface.22,34,35 Solutions were photolyzed using the full output of a Newport/Oriel 75 W O3-

⎡⎛ I (λ ) ⎞ ⎤ ⎛ I(λ)N2 ⎞ 1 − R (λ ) N2 N2 ⎥ ⎜1 − ⎟⎟α(λ)ray ⎟⎟ = ⎢⎜⎜ − α(λ)He ray / ⎜ d I(λ)He ⎠ ⎥⎦ ⎝ ⎣⎢⎝ I(λ)He ⎠ (4)

where d is the length of the cell; I(λ)N2 and I(λ)He are the transmission intensities through N2 and He as a function of N wavelength; and α(λ)ray2 and α(λ)He ray are the calculated Rayleigh cross sections of N2 and He as a function of wavelength. It should be noted that the gas sample did not flow through the entire length of the IBBC-EAS cell because of the mirror purging; this was not accounted for in these calculations since it is expected to result in a very small correction. The Rayleigh cross sections were calculated from36 22261



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Figure 2. Example spectrum of NO2(g) (blue trace) and its fit using DOASIS (red trace) acquired at 20 min of 1 M NO3− photolysis.

σ(λ)ray =

24π 3(n(λ)2 − 1)2 × F (λ ) λ 4N 2(n(λ)2 + 2)2

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RESULTS The effects of stearic acid, octanol, and octanoic acid monolayer coatings on the evolution of NO2 in the gas phase from aqueous NO3− photolysis were studied by measuring [NO2(g)] as a function of time until it reached steady state; i.e., the rate of its production and volatilization into the gas phase was equal to its rate of loss from flowing out of the reaction chamber and through the IBBC-EAS cell. In the presence of stearic acid the observed steady-state [NO2(g)] is the same as for the uncoated solutions. The steady-state [NO2(g)] was smaller for octanolcoated and was larger for octanoic acid-coated solutions, compared to the uncoated solutions. Both the rate of NO2(g) evolution and its steady-state concentration showed a pH dependence. Photolysis of NO3− solutions with no organic coatings reached a steady-state [NO2(g)] of (1.4 ± 0.2) × 1012 molecules/cm3 in less than 20 min (black filled stars in all figures). In the presence of stearic acid monolayers, the rate of growth of [NO2(g)] was slower than for the uncoated solutions until ∼50 min, at which time it reached the same steady-state concentration. This behavior is illustrated as green filled triangles in Figure 3. When octanol monolayers were present at the surface of aqueous nitrate solutions, the steady-state [NO2(g)] was reduced compared to uncoated nitrate solutions. Figure 4a shows [NO2(g)] versus time with no coating (black filled stars) and with an octanol monolayer (red filled diamonds). The presence of octanol decreased the steady-state [NO2(g)] by ∼35% from ∼(1.4 ± 0.2) to (0.9 ± 0.3) × 1012 molecules/cm3. Photolysis of 100 mM NO3− solutions coated with an octanol monolayer also showed suppression of [NO2(g)] compared to 100 mM NO3− uncoated solutions, but these results are not shown because the NO2 released from the coated solutions was below our detection limit. Octanoic acid monolayers present at the aqueous surface enhanced both the evolution rate of NO2(g) and steady-state value of [NO2(g)] relative to uncoated solutions. This can be seen in Figure 5, where the blue filled circles show [NO2(g)] versus time for these conditions. The [NO2(g)] was enhanced at all times, and it reached steady state in the same amount of time

(5)

where the wavelength is in cm; N is the number density of the gas; n is the refractive index as a function of wavelength; and F(λ) is the depolarization term as a function of wavelength, accounting for molecular anisotropy and is therefore 1 for He. The refractive index of He, n(He), is 1.000035 at all wavelengths of interest, and the corresponding terms for nitrogen are36,39 F(N2) = 1.034 +

3.17 × 10−4 λ2

(6)

⎛ 32431.57 × 10−6λ 2 ⎞ n(N2) = 1 + ⎜68.5520 × 10−6 + ⎟ 144λ 2 − 1 ⎝ ⎠ (7)

The extinction due to absorption by NO2 and HONO, α(λ)abs, was calculated as a function of wavelength and reflectivity using eq 838 assuming Rayleigh scattering by N2 is the only other important loss of transmitted light within the cell during an experiment ⎛ 1 − R (λ ) ⎞⎛ I0(λ) − I(λ) ⎞ N2 α(λ)abs = ⎜ + α(λ)ray ⎟⎜ ⎟ ⎝ ⎠⎝ d I (λ ) ⎠

(8)

where I0 and I are the transmission of light through N2 and the reaction sample as a function of wavelength, respectively. The [NO2(g)] was calculated using DOASIS,40,41 a program which evaluates measured spectra by using a linear least-squares method to fit the rapidly varying component (narrow bands due to absorption by a trace species, e.g., NO2) to reference spectra, and to fit the slowly varying component (broad bands due to Rayleigh scattering, Mie scattering and other influences such as temperature drifts) to a polynomial. The DOASIS fit also includes terms that shift and squeeze the fit to account for differences in scaling between reference and measured spectra.40,41 In this study literature absorption spectra of HONO and NO2 were input as reference spectra.42,43 Figure 2 is an example of a measured spectrum (blue line) and a calculated fit (red line) using DOASIS. 22262



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Figure 5. Measured [NO2(g)] as a function of time from the photolysis of 1 M NO3− for solutions with no organic monolayer (black filled stars), no organic monolayer and a bulk pH adjusted to 3.8 (hollow black stars), and an octanoic acid monolayer (blue filled circles). In both experiments the error bars give 1 − σ uncertainties that are based on at least three experiments.

Figure 3. Measured [NO2(g)] as a function of time from the photolysis of 1 M NO3− for solutions with no organic monolayer (black filled stars) and a stearic acid monolayer (green filled triangles). In both experiments the error bars give 1 − σ uncertainties and are based on at least eleven experiments.

as the uncoated solutions (i.e., 20 min) at which point the [NO2(g)] was ∼(1.8 ± 0.3) × 1012 molecules/cm3. To explore whether the enhancement of [NO2(g)] in the presence of octanoic acid was due to an increased bulk aqueous acidity, we measured the [NO2(g)] evolved from uncoated and octanol-coated nitrate solutions adjusted to pHs of ∼3.8, to match that of the octanoic acid solutions. Figure 5 shows the results for acidified uncoated solutions as hollow stars, where [NO2(g)] reaches a steady state of (1.65 ± 0.25) × 1012 molecules/cm3 at about 20 min. The presence of acid in the uncoated solutions clearly does enhance the steady-state [NO2(g)] compared to the pH ∼ 5.4 case. Figure 4b shows the [NO2(g)] measured from the photolysis of aqueous nitrate solutions, adjusted to a pH of 3.8, with octanol coatings (hollow red diamonds). Figure 4b shows that the steady-state [NO2(g)] from acidified octanol-coated solutions is suppressed relative to the acidified uncoated solutions, but it is enhanced compared to octanol-coated solutions with unadjusted pHs (red filled diamonds), from (0.9 ± 0.3) to (1.3 ± 0.1) × 1012 molecules/cm3.

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DISCUSSION In aqueous solutions, NO3− is photolyzed, as shown in eqs 1−3, and the NO2 that is produced can partition to the surface and then into either the aqueous or gas phase. In this study we consider the uptake of NO2 from the gas phase to be negligible since the gases evolved are rapidly advected out of the reaction chamber and into the IBBC-EAS cell. We model the aqueous− surface−gas-phase partitioning behavior of NO2 using

Figure 4. Measured [NO2(g)] as a function of time from the photolysis of 1 M NO3− for (a) solutions with no organic monolayer (black filled stars) and octanol monolayers (red filled diamonds). In both experiments the error bars give 1 − σ uncertainties that are based on at least six experiments. (b) Solutions with no organic monolayer and a bulk pH adjusted to 3.8 (hollow black stars); octanol monolayers (red filled diamonds); octanol monolayers with a bulk pH adjusted to 3.8 (hollow red diamonds). In both experiments the error bars give 1 − σ uncertainties that are based on at least three experiments.

k1

k2

k pump

NO2(aq) XooY NO2(surf) → NO2(g) ⎯⎯⎯⎯→ Flow Out k −1

(9)

where k1 is the rate constant for the adsorption of aqueous NO2 to the aqueous surface; k−1 is the rate constant for the rate of desorption of NO2 from the aqueous surface and into the bulk aqueous phase; k2 is the rate constant for the evaporation of NO2 into the gas phase; and kpump is the inverse of the lifetime of NO2(g) in the chamber. Surface and gas-phase concentrations of NO2 are described by 22263


The Journal of Physical Chemistry C d[NO2(surf)] dt

Article

uncoated case.33 Assuming that the desorption of NO2 into either bulk phase (i.e., k2 and k−1) is affected by the same factor, compared to pure water, when an octanol monolayer is present, then Kads is ∼1 for NO2 in pure water. If, in addition, Kads is well described by Kow for NO2 in aqueous solutions with an octanol monolayer, then the [NO2(g)]S.S. is estimated to decrease by ∼30% in the presence of an octanol coating. This estimate agrees well with the measured change of ∼35% in [NO2(g)]S.S. in the present study. A similar observation has been made for the production of molecular iodine from the ozonation of aqueous iodide, where an octanol monolayer decreased the concentration of I2(g) by ∼50% while increasing the amount remaining in solution.48 It is also possible that additional loss process(es) may decrease the steady-state [NO2(g)] in the presence of an octanol monolayer. Enhanced photolysis of NO2 in the octanol monolayer at the aqueous surface may be one possible loss route of NO2

= k1[NO2(aq)] − (k −1 + k 2)[NO2(surf)] (10)

d[NO2(g)] dt

= k 2[NO2(surf)] − k pump[NO2(g)]

(11)

and thus their steady-state concentrations are [NO2(surf)]s.s. = [NO2(g)]s.s. =

k1 [NO2(aq)]s.s. k −1 + k 2

k2 [NO2(surf)]s.s. k pump

(12)

(13)

Equation 12 can be substituted into eq 13, giving the steadystate concentration of NO2(g) as [NO2(g)]s.s. =

k1k 2 [NO2(aq)]s.s. k pump(k −1 + k 2)

NO2 + hυ → NO + O

(14)

Recent work by Baergen and Donaldson has shown that nitrate photolysis is enhanced by ∼4 orders of magnitude on urban grime compared its rate in the aqueous phase.49 Nitrate photolysis is also enhanced by up to 2 orders of magnitude on Pyrex surfaces due to enhanced absorption cross sections.50−52 It may be that NO2 photolysis is also enhanced at an octanol monolayer on aqueous surfaces. In addition, NO2(aq) is in equilibrium with N2O4, and it is possible that NO2 and/or N2O4 reacts with octanol to form organic nitrites leading to further loss of NO2, although this process is slow in the gas phase.53−56 The release of NO2 into the gas phase was enhanced in the presence of octanoic acid monolayers. The photolysis of aqueous nitrate is known to be enhanced under acidic conditions at λ < 280 nm because the photolysis product peroxynitrate (3) reacts with a proton to produce additional NO212,57

The presence of an uncompressed stearic acid monolayer decreased the rate of appearance of NO2 in the gas phase, but it did not affect the steady-state [NO2(g)] compared to the case of no coating. Consideration of eq 11 suggests that stearic acid lowers the evaporation rate of NO2 from the interface by reducing k2. The k2 was estimated to decrease by a factor of ∼4 in the presence of stearic acid, as determined by fitting both plots in Figure 3 to single exponential growth curves and assuming the rise rate is proportional to k2. A kinetic model using the same factor of 4 decrease in k2 supports the observed results and gives the same steady-state [NO2(g)] for both cases. Studies have shown that compressed stearic acid can inhibit the evaporation of water and the uptake of acetic acid, N2O5, and O3.27−31 In the present work the surface coverage of stearic acid was approximately 30 Ǻ 2/molecule, which is uncompressed on a pure water surface;44 thus, its structure at the aqueous surface was disordered. A decrease in the evaporation rate constant seems feasible since the stearic acid carbon chains are long, are disordered in the uncondensed state, and can cross each other. This decreased evaporation rate may be related to the decreased uptake coefficients reported by McNeill et al. for N2O5 onto water surfaces covered with uncompressed films.45 Therefore, the disordered and long stearic acid chains may inhibit the rate of NO2 evaporation from the aqueous surface, but they have no effect on the steady-state concentration of gasphase NO2. Both the rate of NO2 release and the steady-state [NO2(g)] were smaller when octanol monolayers were present at the aqueous surface. Octanol monolayers at the aqueous surface are known to enhance the reaction between PAHs and ozone due to the increased partitioning of these species to the less polar environment octanol provides.22,23,46 The presence of octanol may enhance the solubility of other nonpolar compounds such as NO2; indeed the estimated octanol to water partition coefficient of NO2, Kow, is 2.7.47 Taking these ideas and defining a surface partitioning coefficient, Kads = k1/k−1, eq 14 is modified to [NO2(g)]s.s. = K ads

k −1k 2 [NO2(aq)]s.s. k pump(k −1 + k 2)

(16)

ONOO− + H+ → → NO2 + OH

(17)

An enhancement in the rate of NO2(g) evolution was also observed for uncoated and octanol-coated solutions with pH adjusted to ∼3.8, compared to the same solutions with unadjusted pH. It is likely that the presence of an octanoic acid monolayer enhances the measured [NO2(g)] because the octanoic acid reduces the pH. Using pKa = 4.89 for octanoic acid58 and the measured water pH of 5.4 prior to adding nitrate and octanoic acid, the estimated percent dissociation of acid in the bulk solution would be about 6%. The measured pH of the bulk solution after adding octanoic acid was about 3.8, and the percent dissociation of acid using this pH and the initial pH is about 5%. The good agreement between the estimated and measured percent dissociation suggests that the presence of nitrate has no influence on the dissociation of octanoic acid. Furthermore, previous work in our group shows that the change in pH at the aqueous surface is the same as that in the bulk aqueous phase.59 Thus, we expect that the presence of nitrate at the aqueous surface has no effect on the dissociation of octanoic acid. An octanoic acid monolayer has a lower surface concentration than an octanol monolayer,22,33 despite their similar size, because the anionic functional groups of the octanoic acid repel each other, which create more gaps in the octanoic acid monolayer compared to octanol. The enhanced rate of NO2(g)

(15)

Mmereki et al. found a factor of 4 increase in the uptake of O3(g) to octanol-coated aqueous surfaces compared to the 22264



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production when octanoic acid is present rather than octanol, at the same pH, could be related to the gaps created at the aqueous surface by the repelling carboxylate groups. Figure 5 shows that steady-state [NO2(g)] from photolysis in solutions with an octanoic acid monolayer is similar to that from uncoated solutions with a pH of ∼3.8 when t ≥ 45 min. However, the initial rate of NO2(g) evolution does appear to be slightly enhanced. Recent studies examining nitrate photolysis in the presence of inorganic salts have suggested that the presence of halides at the aqueous surface may attract bulkphase cations toward the interface, and these may subsequently attract nitrate anions to the water surface. The nitrate may not be fully solvated at the water surface, which could make the NO2 more readily available to evaporate.60,61 If octanoic acid is enhancing the production of NO2(g), it may be because the carboxylate anions have a similar effect.

Article

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The authors would like to thank Dr. Steven Brown for lending the highly reflective 362 nm mirrors; Prof. Cora Young and Prof. Tara Kahan for useful discussions regarding the IBBCEAS; and NSERC and CFCAS for funding this project.

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REFERENCES

(1) Finlayson-Pitts, B. J.; Pitts, J. N. Chemistry of the Upper and Lower Atmosphere: Theory, Experiments and Applications; Academic Press: San Diego, Calif., 2000. (2) Samet, J. M.; Utell, M. J. The Risk of Nitrogen Dioxide: What Have We Learned From Epidemiological and Clinical Studies. Toxicol. Ind. Health. 1990, 6, 247−262. (3) Hasselblad, V.; Eddy, D. M.; Kotchmar, D. J. Synthesis of Environmental Evidence: Nitrogen Dioxide Epidemiology Studies. J. Air Waste Manage. 1992, 42, 662−671. (4) Taylor, O. C.; Eaton, F. M. Suppression of Plant Growth by Nitrogen Dioxide. Plant Physiol. 1966, 41, 132−135. (5) Kleffmann, J.; Becker, K. H.; Wiesen, P. Heterogeneous NO2 Conversion Processes on Acid Surfaces: Possible Atmospheric Implications. Atmos. Environ. 1998, 32, 2721−2729. (6) Atkinson, R.; Winer, A. M.; Pitts, J. N., Jr. Estimation of NightTime N2O5 Concentrations From Ambient NO2 and NO3 Radical Concentrations and the Role of N2O5 in Night-Time Chemistry. Atmos. Environ. 1986, 20, 331−339. (7) Wahner, A.; Mentel, T. F.; Sohn, M. Gas-Phase Reaction of N2O5 with Water Vapor: Importance of Heterogeneous Hydrolysis of N2O5 and Surface Desorption of HNO3 in a Large Teflon Chamber. Geophys. Res. Lett. 1998, 25, 2169−2172. (8) Ljungström, E.; Hallquist, M. Nitrate Radical Formation Rates in Scandinavia. Atmos. Environ. 1996, 30, 2925−2932. (9) Warneck, P.; Wurzinger, C. Product Quantum Yields for the 305nm Photodecomposition of Nitrate in Aqueous Solution. J. Phys. Chem. 1988, 92, 6278−6283. (10) Zellner, R.; Exner, M.; Herrmann, H. Absolute OH Quantum Yields in the Laser Photolysis of Nitrate, Nitrite and Dissolved H2O2 at 308 and 351 nm in the Temperature Range 278−353 K. J. Atmos. Chem. 1990, 10, 411−425. (11) Zepp, R. G.; Hoigne, J.; Bader, H. Nitrate-Induced Photooxidation of Trace Organic Chemicals in Water. Environ. Sci. Technol. 1987, 21, 443−450. (12) Goldstein, S.; Rabani, J. Mechanism of Nitrite Formation by Nitrate Photolysis in Aqueous Solutions: The Role of Peroxynitrite, Nitrogen Dioxide, and Hydroxyl Radical. J. Am. Chem. Soc. 2007, 129, 10597−10601. (13) Mark, G.; Korth, H.-G.; Schuchmann, H.-P.; von Sonntag, C. The Photochemistry of Aqueous Nitrate Ion Revisited. J. Photochem. Photobiol. A. 1996, 101, 89−103. (14) Herrmann, H. On the Photolysis of Simple Anions and Neutral Molecules as Sources of O−/OH, SOx− and Cl in Aqueous Solution. Phys. Chem. Chem. Phys. 2007, 9, 3935−3964. (15) Liss, P. S.; Duce, R. A. The Sea Surface and Global Change; Cambridge University Press: Cambridge, U.K., 2005. (16) Jimenez, J. L.; Canagaratna, M. R.; Donahue, N. M.; Prevot, A. S. H.; Zhang, Q.; Kroll, J. H.; DeCarlo, P. F.; Allan, J. D.; Coe, H.; Ng, N. L.; et al. Evolution of Organic Aerosols in the Atmosphere. Science 2009, 326, 1525−1529. (17) You, Y.; Renbaum-Wolff, L.; Carreras-Sospedra, M.; Hanna, S. J.; Hiranuma, N.; Kamal, S.; Smith, M. L.; Zhang, X. L.; Weber, R. J.; Shilling, J. E.; Dabdub, D.; et al. Images Reveal that Atmospheric

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CONCLUSIONS Stearic acid, octanol, and octanoic acid monolayers were used as proxies for different types of DOM present at aqueous surfaces to better understand how organic films might influence the amount of gas-phase NO2 released from the photolysis of aqueous NO3−. Stearic acid decreases the rate of production of gaseous NO2 because it is physically blocking the aqueous surface, consequently decreasing the rate of evaporation of NO2 from the water surface. Octanol and octanoic acid have similar structures, but they have different effects on the release of NO2 into the gas phase. Octanol enhances the solubility of NO2 at the aqueous surface, and this combined with possible additional loss processes of NO2 in octanol at the aqueous surface limits the steady-state concentration of NO2 in the gas phase compared to no coating. However, octanoic acid increases the evolved NO2(g) because NO3− photolysis is enhanced in acidic conditions. The various properties at the aqueous surface can influence the outcome of NO2 released into the gas phase from aqueous NO3− photolysis. Stearic acid, octanol, and octanoic acid monolayers pack differently at the aqueous surface and subsequently have very different effects on the measured NO2 released into the gas phase. Stearic acid is 18 carbons in length, and these molecules can cross each other when coating a surface uncompressed, which may lead to physically blocking adsorption sites at the aqueous surface. Octanol and octanoic acid are the same carbon length, but they have different functional groups. Octanol packs the aqueous surface tighter than octanoic acid, and the anionic carboxylates of octanoic acid may repel and create gaps at the aqueous surface which could enhance the evolution of NO2 into the gas phase. In addition, studies suggest that NO3− photolysis is enhanced at surfaces compared to bulk phases because the absorption cross sections are up to 2 orders of magnitude greater, and photolysis is up to 4 orders of magnitude greater on grime films. Water surface microlayers are complex systems of biota, organic, and inorganic species, and the present study suggests that the effect organics have on NO2 released into the troposphere from nitrate photolysis in water systems depends on the physical and chemical nature of the organic species present. Further research is needed to gain a better understanding of the influence that different types of aqueous organic coatings at the interface may have on aqueous processes because this will affect the fate of gaseous tropospheric species formed and their environmental implications. 22265



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Particles Can Undergo Liquid-Liquid Phase Separations. Proc. Natl. Acad. Sci. U.S.A. 2012, 109, 13188−13193. (18) Ciobanu, V. G.; Marcolli, C.; Krieger, U. K.; Weers, U.; Peter, T. Liquid-Liquid Phase Separation in Mixed Organic/Inorganic Aerosol Particles. J. Phys. Chem. A 2009, 113, 10966−10978. (19) Donaldson, D. J.; Vaida, V. The Influence of Organic Films at the Air−Aqueous Boundary on Atmospheric Processes. Chem. Rev. 2006, 106, 1445−1461. (20) Donaldson, D. J.; Valsaraj, K. T. Adsorption and Reaction of Trace Gas-Phase Organic Compounds on Atmospheric Water Film Surfaces: A Critical Review. Environ. Sci. Technol. 2010, 44, 865−873. (21) Clifford, D.; Donaldson, D. J. Direct Experimental Evidence for a Heterogeneous Reaction of Ozone with Bromide at the Air-Aqueous Interface. J. Phys. Chem. A 2007, 111, 9809−9814. (22) Mmereki, B. T.; Donaldson, D. J.; Gilman, J. B.; Eliason, T. L.; Vaida, V. Kinetics and Products of the Reaction of Gas-Phase Ozone with Anthracene Adsorbed at the Air-Aqueous interface. Atmos. Environ. 2004, 38, 6091−6103. (23) Henderson, E. A.; Donaldson, D. J. The Influence of Organic Coatings on Pyrene Ozonolysis at the Air-Aqueous Interface. J. Phys. Chem. A 2012, 116, 423−429. (24) Badger, C. L.; Griffiths, P. T.; George, I.; Abbatt, J. P. D.; Cox, R. A. Reactive Uptake of N2O5 by Aerosol Particles Containing Mixtures of Humic Acid and Ammonium Sulfate. J. Phys. Chem. A 2006, 110, 6986−6994. (25) Burden, D. K.; Johnson, A. M.; Nathanson, G. M. HCl Uptake through Films of Pentanoic Acid and Pentanoic Acid/Hexanol Mixtures at the Surface of Sulfuric Acid. J. Phys. Chem. A 2009, 113, 14131−14140. (26) Glass, S. V.; Park, S. C.; Nathanson, G. M. Evaporation of Water and Uptake of HCl and HBr through Hexanol Films at the Surface of Supercooled Sulfuric Acid. J. Phys. Chem. A 2006, 110, 7593−7601. (27) Gilman, J. B.; Tervahattu, H.; Vaida, V. Interfacia Properties of Mixed Films of Long-Chain Organics at the Air−Water Interface. Atmos. Environ. 2006, 40, 6606−6614. (28) Rouvière, A.; Ammann, M. The Effect of Fatty Acid Surfactants on the Uptake of Ozone to Aqueous Halogenide Particles. Atmos. Chem. Phys. 2010, 10, 11489−11500. (29) Archer, R. J.; Mer, V. K. L. The Rate of Evaporation of Water Through Fatty Acid Monolayers. J. Phys. Chem. 1955, 59, 200−208. (30) Cosman, L. M.; Knopf, D. A.; Bertram, A. K. N2O5 Reactive Uptake on Aqueous Sulfuric Acid Solutions Coated with Branched and Straight-Chain Insoluble Organic Surfactants. J. Phys. Chem. A 2008, 112, 2386−2396. (31) Cosman, L. M.; Bertram, A. K. Reactive Uptake of N2O5 on Aqueous H2SO4 Solutions Coated with 1-Component and 2Component Monolayers. J. Phys. Chem. A 2008, 112, 4625−4635. (32) Park, S. C.; Burden, D. K.; Nathanson, G. M. Surfactant Control of Gas Transport and Reactions at the Surface of Sulfuric Acid. Acc. Chem. Res. 2009, 42, 379−387. (33) Mmereki, B. T.; Donaldson, D. J. Direct Observation of the Kinetics of an Atmospherically Important Reaction at the Air− Aqueous Interface. J. Phys. Chem. A 2003, 107, 11038−11042. (34) Benderskii, A. V.; Eisenthal, K. B. Effect of Organic Surfactant on Femtosecond Solvation Dynamics at the Air−Water Interface. J. Phys. Chem. B 2000, 104, 11723−11728. (35) Zhao, X.; Goh, M. C.; Subrahmanyan, S.; Eisenthal, K. B. Studies of Surface Diffusion by Second Harmonic Fluctuation Spectroscopy. J. Phys. Chem. 1990, 94, 3370−3373. (36) Bodhaine, B. A.; Wood, N. B.; Dutton, E. G.; Slusser, J. R. On Rayleigh Optical Depth Calculations. J. Atmos. Oceanic Technol. 1999, 16, 1854−1861. (37) Fiedler, S. E.; Hese, A.; Ruth, A. A. Incoherent Broad-Band Cavity-Enhanced Absorption Spectroscopy. Chem. Phys. Lett. 2003, 371, 284−294. (38) Washenfelder, R. A.; Langford, A. O.; Fuchs, H.; Brown, S. S. Measurement of Glyoxal Using an Incoherent Broadband Cavity Enhanced Absorption Spectrometer. Atmos. Chem. Phys. 2008, 8, 7779−7793.

(39) Weber, M. J. Handbook of Optical Materials; CRC Press: USA, 2002; pp 445−455. (40) Kraus, S. DOASIS: A Framework Design for DOAS; University of Mannheim: Mannheim, Germany, 2006. (41) Lehmann, T. DOASIS - DOAS Intelligent System; https:// doasis.iup.uni-heidelberg.de/ (accessed January 29, 2013). (42) Stutz, J.; Kim, E. S.; Platt, U.; Bruno, P.; Perrino, C.; Febo, A. UV-Visible Absorption Cross Sections of Nitrous Acid. J. Geophys. Res. 2000, 105, 14585−14592. (43) Vandaele, A. C.; Hermans, C.; Simon, P. C.; Carleer, M.; Colins, R.; Fally, S.; Mérienne, M. F.; Jenouvrier, A.; Coquart, B. Measurements of the NO2 Absorption Cross-Sections from 42000 cm−1 to 10000 cm−1 (238−1000 nm) at 220 and 294 K. J. Quant. Spectrosc. Radiat. Transfer 1998, 59, 171−184. (44) Kwok, D. Y.; Tadros, B.; Deol, H.; Vollhardt, D.; Miller, R.; CabrerizoVilchez, M. A.; Neumann, A. W. Axisymmetric Drop Shape Analysis as a Film Balance: Rate Dependence of the Collapse Pressure and Molecular Area at Close Packing of 1-Octadecanol Monolayers. Langmuir 1996, 12, 1851−1859. (45) McNeill, V. F.; Patterson, J.; Wolfe, G. M.; Thornton, J. A. The Effect of Varying Levels of Surfactant on the Reactive Uptake of N2O5 to Aqueous Aerosol. Atmos. Chem. Phys. 2006, 6, 1635−1644. (46) Mmereki, B. T.; Chaudhuri, S. R.; Donaldson, D. J. Enhanced Uptake of PAHs by Organic-Coated Aqueous Surfaces. J. Phys. Chem. A 2003, 107, 2264−2269. (47) Signorelli, S.; Möller, M. N.; Coitiño, E. L.; Denicola, A. Nitrogen Dioxide Solubility and Permeation in Lipid Membranes. Arch. Biochem. Biophys. 2011, 512, 190−196. (48) Reeser, D. I.; Donaldson, D. J. Influence of Water Surface Properties on the Heterogeneous Reaction Between O3(g) and I(aq)−. Atmos. Environ. 2011, 45, 6116−6120. (49) Baergen, A. M.; Donaldson, D. J. Photochemical Renoxification of Nitric Acid on Real Urban Grime. Environ. Sci. Technol. 2013, 47, 815−820. (50) Du, J.; Zhu, L. Quantification of the Absorption Cross Sections of Surface-Adsorbed Nitric Acid in the 335−365 nm Region by Brewster Angle Cavity Ring-Down Spectroscopy. Chem. Phys. Lett. 2011, 511, 213−218. (51) Zhou, X. L.; Gao, H. L.; He, Y.; Huang, G.; Bertman, S. B.; Civerolo, K.; Schwab, J. Nitric Acid Photolysis on Surfaces in LowNOx Environments: Significant Atmospheric Implications. Geophys. Res. Lett. 2003, 30, 12-1−12-4. (52) Zhu, C.; Xiang, B.; Zhu, L.; Cole, R. Determination of Absorption Cross Sections of Surface-Adsorbed HNO3 in the 290− 330 nm Region by Brewster Angle Cavity Ring-Down Spectroscopy. Chem. Phys. Lett. 2008, 458, 373−377. (53) White, E. H.; Feldman, W. R. The Nitrosation and Nitration of Amines and Alcohols with Nitrogen Tetroxide. J. Am. Chem. Soc. 1957, 79, 5832−5833. (54) Fairlie, A. M.; Carberry, J. J.; Treacy, J. C. A Study of the Kinetics of the Reaction between Nitrogen Dioxide and Alcohols. J. Am. Chem. Soc. 1953, 75, 3786−3789. (55) Yoffe, A. D.; Gray, P. 325. Esterification by Dinitrogen Tetroxide. J. Chem. Soc. 1951, 1412−1412. (56) Niki, H.; Maker, P. D.; Savage, C. M.; Breitenbach, L. P. An FTIR Study of the Reaction Between Nitrogen Dioxide and Alcohols. Int. J. Chem. Kinet. 1982, 14, 1199−1209. (57) Mack, J.; Bolton, J. R. Photochemistry of Nitrite and Nitrate in Aqueous Solution: A Review. J. Photochem. Photobiol. A 1999, 128, 1− 13. (58) Haynes, W. M.; Bruno, T. J.; Lide, D. R. CRC Handbook of Chemistry and Physics, 93rd ed., Internet Version 2013 ed.; CRC Press: Boca Raton, FL, 2013. (59) Clifford, D.; Bartels-Rausch, T.; Donaldson, D. J. Suppression of Aqueous Surface Hydrolysis by Monolayers of Short Chain Organic Amphiphiles. Phys. Chem. Chem. Phys. 2007, 9, 1362−1369. (60) Richards, N. K.; Wingen, L. M.; Callahan, K. M.; Nishino, N.; Kleinman, M. T.; Tobias, D. J.; Finlayson-Pitts, B. J. Nitrate Ion 22266



Article

Photolysis in Thin Water Films in the Presence of Bromide Ions. J. Phys. Chem. A 2011, 115, 5810−5821. (61) Wingen, L. M.; Moskun, A. C.; Johnson, S. N.; Thomas, J. L.; Roeselová, M.; Tobias, D. J.; Kleinman, M. T.; Finlayson-Pitts, B. J. Enhanced Surface Photochemistry in Chloride−Nitrate Ion Mixtures. Phys. Chem. Chem. Phys. 2008, 10, 5668−5668.

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Effect of Organic Coatings on Gas-Phase Nitrogen Dioxide Production

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