Release of Nitrous Acid and Nitrogen Dioxide from ... - ACS Publications

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Release of Nitrous Acid and Nitrogen Dioxide from Nitrate Photolysis in Acidic Aqueous Solutions Nicole K. Scharko, Andrew E. Berke,† and Jonathan D. Raff* School of Public and Environmental Affairs and the Department of Chemistry, Indiana University, Bloomington, Indiana 47405-2204, United States S Supporting Information *

ABSTRACT: Nitrate (NO3¯) is an abundant component of aerosols, boundary layer surface films, and surface water. Photolysis of NO3¯ leads to NO2 and HONO, both of which play important roles in tropospheric ozone and OH production. Field and laboratory studies suggest that NO3¯ photochemistry is a more important source of HONO than once thought, although a mechanistic understanding of the variables controlling this process is lacking. We present results of cavity-enhanced absorption spectroscopy measurements of NO2 and HONO emitted during photodegradation of aqueous NO3¯ under acidic conditions. Nitrous acid is formed in higher quantities at pH 2−4 than expected based on consideration of primary photochemical channels alone. Both experimental and modeled results indicate that the additional HONO is not due to enhanced NO3¯ absorption cross sections or effective quantum yields, but rather to secondary reactions of NO2 in solution. We find that NO2 is more efficiently hydrolyzed in solution when it is generated in situ during NO3¯ photolysis than for the heterogeneous system where mass transfer of gaseous NO2 into bulk solution is prohibitively slow. The presence of nonchromophoric OH scavengers that are naturally present in the environment increases HONO production 4-fold, and therefore play an important role in enhancing daytime HONO formation from NO3¯ photochemistry.

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INTRODUCTION Nitrous acid (HONO) is a major source of hydroxyl radical (OH) and nitric oxide (NO), both of which play a role in ozone and aerosol formation in polluted air.1−5 Accumulation of HONO, up to several parts-per-billion (ppb) during the nighttime,6−9 has been attributed to reactions of nitrogen dioxide (NO2) on boundary layer surfaces,10−13 but field studies indicate there may be photochemical sources of HONO as well.14−18 Potentially important sources of HONO are the photolysis of surface-adsorbed nitrate (NO3¯) or nitric acid (HNO3), which are thought to play a role in HONO emissions from the surface of glass sampling manifolds,19 chambers,20 aerosols,21,22 and vegetation.23,24 High HONO emission rates observed during these studies have puzzled researchers since HONO is only expected to be a minor product when NO3¯ or HNO3 is photolyzed by actinic radiation.25,26 Numerous explanations have been offered, but the exact mechanisms responsible for these observations remain uncertain. Here, we study NO3¯ photodegradation in aqueous solutions with a focus on elucidating the reaction channels leading to the release of HONO under acidic conditions. The nitrate ion is ubiquitous in the air, soil, and water of both urban and rural areas.4 In regions polluted by nitrogen oxides (NOx ≡ NO + NO2), major sources of NO3¯ are the gas phase oxidation of NO2 by OH and hydrolysis of N2O5.4 Once © 2014 American Chemical Society

formed, HNO3 readily partitions into aqueous droplets or reacts with ammonia (NH3) to form aerosols consisting of ammonium nitrate (NH4NO3) in concentrations as high as 8− 26 M.27,28 Partitioning of HNO3 and disproportionation of NH4NO3 within aerosols means that the pH may drop below 2, with one study documenting aerosol pH in Southern California ranging from −2 to 0.2.29 Dry deposition of HNO3 and settling out of acidic aerosols invariably coats these surfaces with thin aqueous films that are acidic.30−32 Evidence that ground level surfaces are acidic stems from observations of HONO emissions from vegetation, soil, and urban surfaces.9,12,33 These emissions are only possible if the surface films coating these surfaces are below the pKa1 of HONO (∼3) since above this pH the dominant species present is aqueous nitrite (NO2¯).34,35 The photolysis of aqueous NO3¯ under actinic radiation (λ > 290 nm) yields either NO2 and O¯ (Reaction 1) or NO2¯ and O(3P) (reaction 2) with relatively low quantum yields.25,36 In addition, NO3¯ photodegradation is accompanied by OH Received: Revised: Accepted: Published: 11991

June 25, 2014 September 23, 2014 October 1, 2014 October 1, 2014 dx.doi.org/10.1021/es503088x | Environ. Sci. Technol. 2014, 48, 11991−12001

Environmental Science & Technology

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as a function of pH. Modeling of the measured data is used to examine the role of pH and secondary reactions on the amount of HONO and NO2 released. The results help explain an apparent enhancement in HONO released during NO3¯ photodegradation under acidic conditions and provide insight into the potential significance of NO3¯ as a photochemical source of reactive nitrogen in regions impacted by nitrate deposition.

formation stemming from the protonation of O¯ (reaction 3) at pH below 11.9.37 NO3− + hν → NO2 + O− ; Φ1 = 0.01

(1)

NO3− + hν → NO2− + O(3P); Φ2 = 0.001

(2)

O− + H3O+ ↔ OH + H 2O

(3)

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Under acidic conditions, protonation of NO2¯ will lead to HONO (pKa1 = 3.2) and H2ONO+ (pKa2 = 1.7):34,35 NO2− + H3O+ ↔ HONO + H 2O

(4)

HONO + H3O+ ↔ H 2ONO+ + H 2O

(5)

MATERIALS AND METHODS Photochemical Experiments. Experiments were conducted in a cylindrical glass cell (50 mm o.d. × 100 mm high) equipped with 6.35 mm i.d. inlet and outlet ports for gas and solution transport (Figure 1A). The top of the reservoir

A number of hypotheses have been offered to explain why adsorbed NO3¯ (or HNO3) is more efficient at producing HONO than what their known aqueous or gas phase chemistry predicts. One hypothesis is that HONO is generated from hydrolysis of NO2,23,38−40 possibly aided by the formation of surface complexes,41−43 or via sensitization by photoexcited organic chromophores.44−47 Photosensitization processes have also been suggested to play a role in directly reducing HNO3 or NO3− to HONO,48,49 although the mechanisms have proven difficult to quantify. It has also been suggested that HNO3 absorption cross sections are enhanced when nitric acid is adsorbed to a surface and this may result in higher HONO emission rates.50,51 We reason that pH likely plays an important role in many of the studies where HONO has been linked to NO3¯ and HNO3 photochemistry. Unfortunately, previous investigations of the pH dependence of NO3− photodissociation (both in the absence and presence of OH scavengers) have focused on aqueous nitrite formation,25,37,52−55 and not gas phase HONO formation. Early work by Daniels et al. showed that Φ2 is low at pH < 6, but increases above pH 6 in a step-like pattern, with plateaus at pH 9.5 and 12.55 These authors showed that the addition of OH scavengers resulted in a 5-fold increase in rate of NO2¯ formation at neutral pH, compared to when no scavenger was present.55 Warneck and Wurzinger also observed that Φ2 is higher in the presence of an OH scavenger and demonstrated that Φ2 is constant at pH 5−12, but drops steeply at pH < 5.54 Furthermore, the values of Φ2 based on measurements of O(3P) were an order of magnitude lower than those determined from measurements of NO2¯. On the basis of this, it was suggested that hydrolysis of NO2 may be partially responsible for the NO2¯ formed at pH > 5. Unfortunately, these previous studies did not measure HONO and were therefore unable to verify that HONO was emitted by these solutions or study the aqueous chemistry leading to HONO formation in these systems. This is not surprising as HONO is notoriously difficult to measure due to its tendency to partition between gas phase and surfaces, decompose, and to exist in ionic and neutral forms depending on pH. In addition, only recently have analytical methods been developed to measure HONO under atmospherically relevant conditions (e.g., at ppt−ppb levels in the presence of water vapor). In this work, we simultaneously measure gas phase HONO and NO2 using cavity-enhanced absorption spectroscopy (CEAS). This broadband spectroscopic method has been applied previously in laboratory studies of HONO sources and sinks.56−59 However, the present work is the first to make direct measurements of HONO and NO2 released into the gas phase from the UV−visible irradiation of aqueous solutions of NO3¯

Figure 1. (A) Schematics of the continuously stirred flow reactor used for NO3¯ photolysis experiments. (B) Emission spectrum of the HgXe arc lamp, noon-time solar actinic flux at ground level, and the absorption spectrum of aqueous nitrate.

was sealed with a quartz window. Solutions were irradiated from above using a Xe-Hg arc lamp (Spectra Physics, Oriel 200 W) fitted with a UV cutoff filter (λ > 290 nm) and a dichroic mirror to eliminate infrared radiation. A lamp intensity of 1.2 × 1016 photons cm−2 s−1 was determined in the range 290 < λ < 420 nm using NO2 actinometry.60,61 Figure 1B shows that the lamp output is comparable in light intensity to noon-time solar actinic flux in the region of 290−350 nm, which corresponds to the n → π* electronic transition of aqueous NO3¯. A peristaltic pump was used to deliver and extract the solution from the reactor at a rate of 5−6 cm3 min−1 to allow for a continuous replenishment of reactants during the experiment. Both the solution in the reactor (∼8 mL) and reservoir used to supply the reactor were magnetically stirred during all experiments. A continuously stirred flow reactor design was used to ensure that the amount of HONO and NO2 emitted reached steady-state within 10 min of starting each 11992

dx.doi.org/10.1021/es503088x | Environ. Sci. Technol. 2014, 48, 11991−12001


Article

experiment. The flowing experiments were performed at ambient temperature (∼22 °C) and pressure. Solutions of NaNO3 (Sigma-Aldrich, ACS reagent, ≥99%) were prepared using nanopure water (resistivity >18 MΩ cm, Milli-Q) that was acidified by adding concentrated H2SO4 (Fluka, trace select, ≥95%) to achieve a desired pH. For scavenger experiments, either ethylene glycol (>99%, Macron Fine Chemicals) or sodium benzoate (>99%, Fisher Scientific) was added to the pH-adjusted sodium nitrate solution. The acidity of the solution was measured using a pH meter (Thermo Scientific Orion Star A211). A carrier gas of dry zero air was flowed through the reactor headspace at a rate of 1000 cm3 min−1 and then directed to the cavity-enhanced absorption spectrometer (CEAS) for quantitation of the HONO and NO2 emitted from solution. Detection of HONO and NO2. Nitrous acid and NO2 concentrations were measured using a home-built cavityenhanced absorption spectrometer (CEAS); see Figure 2A.

dry air is provided to each mirror. Although the gas sample does not flow through the entire length of the cavity, the effect of the mirror purge gas on the final measured concentration has been shown to be negligible.67 Light exiting the cavity is focused by a 100 mm plano-convex lens, filtered to remove spurious light at λ > 380 nm, and then directed into a monochromator and temperature-controlled charge-coupled device (Ocean Optics, QE65000). Transmission spectra are collected for 10 min (30 averaged spectra with integration time of 20 s). This, combined with the mirror reflectivity gives a typical detection limit of slightly better than 600 ppt for both HONO and NO2. Data analysis is based on the differential optical absorption spectroscopy (DOAS) retrieval software DOASIS,68,69 from which simultaneous HONO and NO2 concentrations are obtained. An example result of a DOAS fit is shown in Figure 2B, highlighting the ability of the DOAS retrieval method to deconvolute multiple overlapping absorbers. The final, dark blue fit spectrum at the top of Figure 2B results from using the following fit equation, which is a modification of the Beer−Lambert expression: I(λ) = exp[∑ ciσi(si + tiλ) + P(λ)] + O(λ)

(6)

where si, ti, and ci are the shift and squeeze fit parameters and the fit coefficient (nominally, the concentration), respectively, for each absorber, i. The shift and squeeze are reference spectrum adjustment parameters within the retrieval software for center wavelength and peak width, respectively. These values are typically set to their default settings and are not allowed to vary. The terms P(λ) and O(λ) are arbitrary fitting polynomials that account for any broad, featureless background signal not related to the absorption spectra of interest.68,69 The DOASIS fitting software simultaneously adjusts user-provided reference spectra and the background polynomials to find the best final fit to the measured spectrum. The individual components of the final fit can be seen in Figure 2, as well as the fit residual. Reference spectra for both HONO and NO2 were obtained using our CEAS, with known concentrations of each gas. Nitrogen dioxide was diluted from a tank provided by Matheson Tri-Gas (20.9 ppm of NO2 in N2), while nitrous acid was generated using the method of Febo et al. via the reaction of HCl with NaNO2.70,71

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RESULTS AND DISCUSSION pH Dependence of Nitrate Photodissociation. Figure 3 shows gas phase concentrations of HONO and NO2 generated from the photolysis of aqueous NaNO3 as a function pH. The amount of HONO produced sharply increases as the pH decreases below the pKa1 of HONO, with the highest amount generated (8.1 ppb) at pH 1.86. Above the pKa1, the HONO levels drop to below the detection limit of the CEAS system (∼600 ppt). This pH profile is not surprising since nonvolatile NO2¯ is expected to be the dominant species in solution above pH 3.2, while HONO is more abundant below this pH.34,35 The amount of NO2 emitted from solution was between 6 and 12 ppb, with no statistically significant trend observed over the entire pH range studied. Previous photochemical studies conducted on nitrate solutions and using wavelengths in the actinic region also reported that product yields of reaction 1 are pH-independent.37,72 Zellner et al. reported OH quantum yields for the photolysis of aqueous nitrate that were independent of pH between pH 4 and 9.37 Since OH is generated from the O¯ photoproduct in reaction 1, OH

Figure 2. (A) Schematic of the cavity-enhanced absorption spectrometer (CEAS) used for NO2 and HONO quantitation and (B) example fit results from the DOAS retrieval procedure. The top spectrum (dark blue) is the final fit to the measured data (green). Below that are the individual contributions to the measured spectrum from NO2 (orange fit on blue) and HONO (red fit on black). Not shown is the polynomial used to fit the background. See text for more details.

Our setup is functionally similar to several other published designs.62−66 Briefly, we focus a temperature-controlled incoherent, broadband light-emitting diode (350 mW, Nichia Corporation) with a center wavelength of 368 nm into a 1.013 m long, 1 in. diameter high-finesse cavity using an aspheric condenser lens. Two highly reflective mirrors (R ∼99.986%) form the ends of the cavity, resulting in a path length of approximately 7.1 km. Mirror reflectivity is monitored regularly using the method of Washenfelder et al.63 To maintain reflectivity, a constant flow of approximately 250 cm3 min−1 of 11993



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Figure 3. Measured gas phase NO2 or HONO concentrations as a function of pH from the photolysis of 0.125 M NaNO3 solutions. Error bars represent uncertainty of the least-squares fit of reference spectra to the measured spectra. Lines are simulated concentrations determined from model simulations using reactions R1−R7 and R15− R20 listed in Table 1. Green dotted line is the pKa1 of HONO for reference. The limit of detection for HONO and NO2 is ∼0.6 ppb.

Figure 4. Measured NO2 or HONO concentrations as a function of pH from the photolysis of 0.125 M NaNO3 solutions in the presence of either 0.01 M ethylene glycol or 0.005 M benzoate. Error bars represent uncertainty of the least-squares fit of reference spectra to the measured spectra. Lines are simulated concentrations determined from model simulations using reactions R1−R20 listed in Table 1. Green dotted line is the pKa1 of HONO for reference. The limit of detection for HONO and NO2 is ∼0.6 ppb.

quantum yields are an indicator of the NO2 production efficiency. Richards and Finlayson-Pitts observed that the rate of NO2 production was independent of pH between 0.5 and 7.5 when deliquesced films of NaNO3-sea salt mixtures were irradiated with UV−visible light.72 Figure 3 also provides evidence of more complicated secondary reactions involving both NO2 and NO2¯/HONO once they are generated. In particular, the amount of NO2 and HONO emitted from solution are similar at pH 2. This is surprising since the measured quantum yields of reaction 1 and 2 would predict a [NO2]/[HONO] ratio of 10:1.25,54 pH Dependence in the Presence of OH Scavengers. Aqueous NO3¯ solutions were photolyzed in the presence of an OH scavenger between pH 2−6 to determine whether OH radical influences the measured NO2 and HONO levels. Nonvolatile ethylene glycol was chosen because it does not interfere with the spectroscopic measurements. Several experiments using the more commonly used OH trap benzoate were also carried out;73 both ethylene glycol and benzoate gave similar results. In order to determine the concentration of scavenger necessary to trap all the nascent OH, the amount of NO2 and HONO formed from nitrate photodegradation was measured as a function of scavenger concentration. As shown in Supporting Information Figure S2, the amount of HONO and NO2 present in the system did not change for ethylene glycol concentrations above 0.01 M (0.005 M for sodium benzoate), indicating that OH was completely scavenged. Measured concentrations of HONO and NO2 evolved from solutions of irradiated nitrate in the presence of an OH scavenger are shown in Figure 4. The amount of HONO formed is strongly pH-dependent and significantly higher than was observed during experiments conducted when OH is present. At pH 2 the concentrations of HONO are four times higher than those of NO2 and it is only at pH > 4 that the amount of HONO emitted from solution drops to below that of NO2. The amount of NO2 formed in the presence of OH scavengers is 2−3 times less than what is formed in the presence of OH, approaching 4 ppb at pH 6.

The enhancement of HONO levels and suppression of NO2 formation in the presence of a scavenger can be explained in part by a mechanism where OH reacts with NO2¯ and HONO to form NO2 according to reactions 7 and 8.26,74 It is also possible that OH reacts with NO2 to form nitric acid (reaction 9),74 NO2− + OH → NO2 + OH−; k 7 = 1 × 1010 M−1 s−1 (7) 9

HONO + OH → NO2 + H 2O; k 8 = 3 × 10 M

−1 −1

s

(8) +

NO2 + OH → H +

NO3− ;

9

k 9 = 4.5 × 10 M

−1 −1

s

(9)

Dubowski et al. reached a similar conclusion after finding that the amount of NO2¯ formed when frozen solutions of NO3¯ were irradiated increased 5-fold in the presence of formate relative to when formate was not added to the ice.38 Although the above mechanism helps rationalize the apparent role of OH in controlling the NO2 and HONO levels, it does not explain why the amount of HONO formed is so high relative to NO2, especially under acidic conditions. To explain this, we consider three possibilities: (1) the molar absorptivity of NO3− is enhanced under acidic conditions, (2) the effective quantum yield of NO2− formation is enhanced under acidic conditions, and (3) secondary reactions of NO2 efficiently lead to HONO/NO2− in this system. Considering Enhanced Absorptivity or Effective Quantum Yields. There have been suggestions that the absorptivity of NO3− and HNO3 in certain environments may be enhanced relative to what is measured in solution or (in the case of HNO3) in the gas phase, and that this may influence the amount of photoproducts formed.50,51,75 To further investigate whether the n → π* electronic transition (λmax ∼ 310 nm) for aqueous NO3− is enhanced under acidic conditions, the UV− visible absorption spectra of 0.0625 M NaNO3 at pH = 1.94, 3.02, and 5.67 were collected. The resulting spectra 11994



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NO2 Hydrolysis as a Source of HONO. A more intuitive explanation for the relative levels of HONO and NO2 observed in Figure 4 is that secondary reactions of NO2 are responsible for some of the HONO measured. Hydrolysis of NO2 (reaction 10)79−83 is thought to involve the dimerization of NO2 to N2O4,84−86 which subsequently rearranges to an asymmetric form and undergoes autoionization.11,87−92 Theory predicts that the formation of the ion pair (NO+)(NO3¯) in large water clusters is thermodynamically feasible and considered an intermediate step in the production of HONO and HNO3.89,92

(Supporting Information Figure S1) indicate that there is no enhancement in the molar absorption coefficients over the indicated pH range. We also considered the possibility that Φ2 is enhanced under acidic conditions. As described by reactions 1 and 2, nitrate’s absorption of a photon leads to dissociation of a N−O bond within a solvent cage comprised of water molecules. For both channels, there is an attractive force between the O¯···NO2 or O···NO2¯ photofragments that will reduce their escape velocity and hence, diminish the quantum yields.25 Further decreases in dissociation efficiency occur when electronically excited NO3¯ undergoes deactivation via internal conversion upon collision with solvent molecules. Reactions of the photofragments with reactive scavengers present in high concentrations are known to compete with geminate recombination, thereby increasing the fraction of products that escape the solvent cage.76 For example, it was shown that H3O+ acts as a scavenger to increase the effective quantum yields of I¯ photolysis.77 In this case, excitation of I¯ generates a free electron in the solvent cage that can be scavenged by H3O+ before recombination occurs. We considered whether [H3O+] at low pH is high enough to act as a scavenger and protonate the NO2¯ photofragment before it can recombine with O(3P) and regenerate NO3¯. To do this, we compared the total amount of N(III) species [N(III) = H2ONO+(aq) + HONO(aq) + NO2¯(aq)] formed in solutions at pH 6 with the amount formed at pH 2.34,35 If Φ2 is enhanced under acidic conditions, we expect to detect higher concentration of N(III) at pH 2 relative to pH 6. At pH 6, all of the generated N(III) is in the form of aqueous NO2¯. The amount of NO2¯ present in solution was determined by first converting it to a purple azo dye and measuring its absorbance with UV−visible spectroscopy.78 Using the method described in the Supporting Information, the NO2¯ concentration determined after irradiation of 0.125 M NaNO3 in the presence of ethylene glycol at pH 6 was found to be 74 μg L−1 (Supporting Information Figure S3). Determination of N(III) present when nitrate is photolyzed at pH 2 is complicated by the following: (1) The CEAS only measures gas phase HONO emitted from solution, which is only a fraction of the amount of N(III) present in solution; (2) aqueous NO2−, which is measured using the azo dye technique, is only a minor species at this pH and difficult to detect; and (3) the amount of HONO emitted depends on the water-to-air exchange kinetics of the system. This problem is simplified by measuring the amount of HONO evolved from solutions prepared from known amounts of NO2¯, under conditions that are identical to those used for the photochemical experiments described above. Aqueous solutions comprised of 0.01 M ethylene glycol and known concentrations of nitrite fixed at a pH of 2 were flowed through the reaction cell in the dark and the gas phase HONO was measured. The relationship between added [NO2¯(aq)] and [HONO(g)] allowed us to correlate the amount of HONO measured by the CEAS to the amount of N(III) actually present in the aqueous phase (Supporting Information Figure S4). Using this method, we determined that the 12.7 ppb of HONO measured over irradiated nitrate solutions at pH 2 (Figure 4) corresponds to an aqueous concentration of aqueous N(III) of 73 μg L−1. This is identical to the aqueous [NO2¯] measured at pH 6. The similarity in the amount of aqueous N(III) formed at pH 2 and 6 provides strong evidence against an enhancement in Φ2 under acidic conditions.

2NO2 + H 2O → HONO + HNO3

(10)

A recent molecular dynamics investigation at the DFT level also supports the above-mentioned mechanism, but also reveals the possibility of a water-assisted mechanism that does not involve the N2O4 dimer.90 In the latter case, the ion pair forms when a NO2 molecule complexes to the water end of an NO2(H2O)2 cluster.90 A related mechanism involves facilitation of NO2 hydrolysis by halides and HSO4¯.42,43,93 In this case, the rate-determining step is complexation and charge transfer between the first gas phase NO2 molecule and an anion, X¯. The resulting (X)(NO2•¯) complex reacts rapidly with a second NO2 molecule to generate HONO and HNO3.93 Additionally, trapping of NO2 by anions at the air−water interface may be enhanced by cationic surfactants that draw anions closer to the interface.43 To investigate whether NO2 hydrolysis plays a role during our experiments, we measured NO2 and HONO concentrations while 20.8 ppb of NO2 was bubbled through 25 mL of pure water or an aqueous solution of 0.125 M NaNO3 over a range of pH. As shown in Figure 5, the concentration of NO2

Figure 5. Concentration of NO2 and HONO measured after 20.8 ppb of NO2 was bubbled through 25 mL of pure water or aqueous NaNO3 (0.125 M) solutions at various pH. Error bars represent uncertainty of the least-squares fit of reference spectra to the measured spectra. The limit of detection for HONO and NO2 is ∼0.6 ppb.

dropped by ∼0.41 ppb at the most, while the amount of HONO present increased on average by 0.57 ppb, regardless of whether NO2 was bubbled through pure water or a solution of nitrate. Similar results were obtained when ppm levels of NO2 were bubbled through water. Clearly, NO2 did not dissolve to a significant extent, even under acidic conditions or in the presence of NO3¯ (and SO42− added during pH adjustments). These results are in agreement with previous studies that find 11995



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Table 1. Reactions and Rate Constants Used to Model NO2 and HONO Formation from Nitrate Photochemistry in Aqueous Solutions no.

reaction

rate coefficient

ref

R1 R2 R3 R4 R5 R6 R7 R8 R9 R10 R11 R12 R13 R14 R15 R16 R17 R18 R19 R20

NO3¯ + hν (+H+) → NO2 + OH NO3¯ + hν → NO2−+ O(3P) 2 NO2 (+H2O) → 2H+ + NO2− + NO3− NO2 + OH → NO3− + H+ NO2 + hν → NO + O(3P) NO2¯ + OH → NO2 + OH− NO2¯ + hν → NO + OH OH + C2H6O2 → HO2 + C2H4 + H2O HONO (+C2H6O2) → products (e.g., RONO) NO2 + HO2 → HOONO2 HOONO2 → H+ + NO2− + O2 HOONO2 → HO2 + NO2 HO2 + OH → H2O + O2 2HO2 → H2O2 + O2 NO2¯ + H+ ↔ HONO HONO → HONO(g) NO2 → NO2(g) HONO + hν → NO + OH HONO(g) + hν → NO(g) + OH(g) NO2(g) + hν → NO(g) + O(g)

8.5 × 10−7 s−1 8.5 × 10−8 s−1 4.0 × 107 M−1 s−1 4.5 × 109 M−1 s−1 1.1 × 10−2 s−1 1.0 × 1010 M−1 s−1 1.0 × 10−4 s−1 1.7 × 109 M−1 s−1 7.0 × 10−2 s−1 1.0 × 108 M−1 s−1 7.0 × 10−4 s−1 8.6 × 10−3 s−1 1.0 × 1010 M−1 s−1 1.0 × 106 M−1 s−1 k15/k−15 = 103 M−1 s−1 7.2 × 10−7 to 2.5 × 10−5 s−1 between pH 6 and 2 8.0 × 10−4 s−1 1.0 × 10−3 s−1 1.0 × 10−3 s−1 1.1 × 10−2 s−1

this work this work 79, 101 74 this work 111 this work 112 this work 54, 74 74 74 113 114 35 this work this work this work this work this work

on a modification of the reaction schemes reported elsewhere for nitrate photochemistry54,74,98 and are listed with their room temperature rate constants in Table 1. The Kintecus compiler99 was used to model the system as a continuously stirred tank reactor comprised of aqueous and gas phases, where a constant supply of aqueous NO3¯, H+, and OH-scavenger (when present) is provided by an external reservoir and flowed through the reactor at a constant rate. The model is constrained by the dimensions of the photochemical reactor, the gas and aqueous phase residence times, measured gas phase concentrations of NO2 and HONO, aqueous [N(III)] for experiments utilizing the OH-scavenger, estimated photolysis rate constants, and previously determined reaction rate coefficients for reactions of HOx (HOx ≡ OH + HO2) and NOx in solution. Transfer of HONO and NO2 between the aqueous and gas phases was treated as a first order process with rate constants derived using the two-film model described by Schwarzenbach et al.;100 see Supporting Information. In the absence of an OH radical scavenger, measured HONO and NO2 concentrations are reproduced by a mechanism that includes reactions R1−R7 and R15−R20 (Table 1). In this case, optimization was achieved by allowing the rate coefficient for R3 to vary within the bounds of reported literature values, (1−7) × 107 M−1 s−1.79,101 Initial values of photolysis rate coefficients were estimated from the measured lamp intensity, absorption spectra, and reported quantum yields photolysis rate coefficients for NO3¯, NO2¯, and HONO; these were allowed to vary by ±20% during optimization. As shown in Figure 3, the good agreement between measured HONO and NO 2 concentrations and model simulations suggests that a mechanism involving six chemical reactions and the pHdependent speciation of N(III) and partitioning of HONO into the gas phase is sufficient to explain the high HONO/NO2 ratios in the presence of OH at pH < 6. Addition of ethylene glycol (C2H6O2) required the addition of reactions R8−R14 (Table 1) to account for the reaction of OH with C2H6O2 and subsequent reactions of HOx and nitrogen oxide species in solution. It was assumed that C2H6O2

hydrolysis of gaseous NO2 to be insignificant at low NO2 concentrations due to aqueous phase mass transport limitations.94−97 There is reason to believe that the kinetics of NO2 hydrolysis at the air−water interface is not representative of what happens following primary photodissociation of dissolved NO3¯. During the experiments depicted in Figure 5, there is a significant thermodynamic barrier to dissolution as the slightly hydrophobic NO2 molecule must diffuse across the air−water interface of a bubble to become solvated in bulk water; i.e., mass transfer is the rate limiting step.94,95 In contrast, during the photolysis experiments, NO2 is formed in situ from solvated NO3¯ photochemistry. Since NO2 is already solvated upon formation, mass transfer limitations are nonexistent and subsequent hydrolysis ensues.94,95 In the absence of mass transfer limitations the amount of HONO formed is controlled rather by the fraction of NO2 molecules that are able to escape geminal recombination in the solvent cage and dimerize to form N2O4. However, when NO3¯ photolyzes in atmospheric particles or on boundary layer surfaces, it is possible that an abundance of organic molecules present in these systems will scavenge OH and prevent nongeminal recombination of OH and NO2. More work is needed to ascertain the influence this has on the effective quantum yield of NO2 formation. It should be mentioned that although theoretical studies on reactions in small water clusters have been used in the context of understanding reactions of NO2 at the air−water interface (reaction 10), they are also applicable for modeling for reactions of NO2 in bulk water.89,92 Indeed, quantum mechanical models of reactions in small water clusters are often the only computationally feasible way to model reactions in solution. Therefore, a wealth of theoretical work exists supporting the thermodynamic feasibility of NO2 → HONO conversion in our experiments.87−92 Modeling HONO Release from Aqueous NO3¯ Photochemistry. We constructed a steady-state box model to further our understanding of the reactions leading to HONO and NO2 formation in our experiments. The reactions included are based 11996



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surfaces and in aerosols and surface water. For example, there is evidence that reactive nitrogen emissions are enhanced when NO 3 ¯ is irradiated in the presence of organic molecules.45,47−49,106 Those studies suggest that photoexcited polycyclic aromatic hydrocarbons or chromophores present in humic acid may be able to reduce NO3¯ to HONO, although these mechanisms are difficult to quantify and remain uncertain. In our OH scavenger experiments, enhanced HONO formation was observed in the presence of organic compounds that did not absorb light from our excitation source. While organic photosensitizers may be important under certain conditions, our modeling results highlight the role of organic compounds as radical scavengers that eliminate OH before it can react with HONO/NO2¯. In addition, HO2 formed during oxidation of organic matter will both react with NO2 to produce NO2¯ and limit the amount of HONO formed by removing NO2 from the system. The chemistry described here represents a potentially important source of HONO in urban areas where atmospheric deposition of HNO3 contributes NO3¯ to the thin films of water coating boundary layer surfaces, and leads to acidification of these surfaces. The abundance of NO3¯ in urban areas is apparent from the high loads of NO3¯ measured in urban stormwater runoff.107−109 For example, in a modeling study of nitrogen deposition in the Ballona Creek watershed, which includes downtown Los Angeles, it was estimated that one rain event washed off 25 tonnes of NO3¯ into Santa Monica Bay.107 This was estimated to be only 10% of the total NO3¯ present in the watershed, where the estimated surface concentration was 0.73 g m−2.107 If the 25 tonnes of NO3¯ were photolyzed rather than washed off, it would have increased the reactive nitrogen level in air over this region by 33 ppb (assuming an airshed volume of 300 km3). This can be compared to estimates of reactive nitrogen emitted by aerosol nitrate. During the CalNex 2010 field campaign, NO3¯ concentrations averaged 2.4 μg m−3 in Pasadena, CA in airborne particles.110 If all this nitrate was photolyzed it would yield a ∼1 ppb increase in reactive nitrogen in the atmosphere if we assume these particle nitrate levels are typical of the entire airshed. This simple comparison suggests that surface NO3¯ in urban areas may be a significant reservoir of photochemical HONO and NOx that is not considered in air quality models. Unfortunately, previous field studies aimed at understanding daytime HONO in urban areas have not included measurements of NO3¯ surface concentrations, so we do not know whether they correlate to daytime HONO. Also, rapid photochemistry of NO3¯ on real surfaces47 may mean that NO3¯ is only present in low levels, making it difficult to measure reproducibly and easy to underestimate the importance of nitrate photochemistry to air quality.24,47,48 Clearly, more work needs to be done to establish the impact of NO3¯ on photochemical air pollution in urban areas.

generates HO2 in unity yield since OH + alcohol reactions are known to produce HO2 as a major product.102 In addition, it was necessary to include reaction R9 to account for the loss of HONO(aq) to the formation of glycol-derived alkyl nitrites.103,104 As shown in Figure 4, the model including all the reactions in Table 1 is able to simulate the measured HONO and NO2 data reasonably well over the indicated pH range. The model results provide insights into the role of HO2 during NO3¯ photolysis in the presence of a scavenger. Our initial supposition was that the reaction of HO2 with NO2 will be an important source of NO2¯ (via HOONO2 decomposition, reaction R11).105 Indeed, when removing reaction R11, the amount of gas phase HONO predicted by the model was ∼20% lower than the model that included this reaction. It also appears that the HOONO2 intermediate serves an important function to moderate the amount of NO2 and HONO formed in the system, since a second decomposition channel recycles HOONO2 back to NO2 via reactions R10 and R12. The modeled HONO concentration was most sensitive to the nitrate photolysis rate and amount of NO2 formed. Any reaction that removes NO2 from the system acts to limit the amount of HONO that can be formed by NO2 hydrolysis (reaction R3). Reactions of OH radical with dissolved nitrogen oxides will limit the amount of HONO formed in the absence of a scavenger. However, in the presence of an organic OH trap, the reaction with HO2 is an important sink for NO2. For example, when reaction R10 is selectively removed from the mechanism in Table 1, the model overestimates gas phase HONO and NO2 levels by a factor of 2.4 relative to experimental observations.

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ATMOSPHERIC IMPLICATIONS The presented work demonstrates that aqueous solutions of NO3¯ irradiated with UV−visible light in the actinic range (λ > 290 nm) are capable of producing levels of HONO/NO2¯ in excess of what is expected based on known quantum efficiencies of reactions 1 and 2. The evidence accumulated here shows that HONO/NO2¯ production is not due to enhancements of NO3¯ absorption coefficients or quantum yields. Instead, experiments designed to study the effect of OH radical scavengers on the amount of product formed are consistent with a mechanism in which significant amounts of HONO are derived from secondary reactions of solvated NO2. A particularly important insight is that the kinetics of NO2 hydrolysis following in situ production from NO3− photoodissociation is much more efficient than for the case when gas phase NO2 reacts at the air−water interface. Mass transfer limitations ensure that hydrolysis of gaseous NO2 on surfaces is slow under atmospherically relevant levels of NO2,94,95 but this is not the case for solvated NO2 generated from aqueous NO3¯ photodissociation. This insight has important implications for our understanding of other atmospherically relevant reactions involving NO2 as an oxidant. Previous studies have suggested that the reaction between gaseous NO2 and S(IV) species is too slow to be important. However, NO2(aq) generated in situ from the photolysis of nitrate may be an important oxidant of aerosol S(IV), where rate constants of between 105 and 107 M−1 s−1 have been inferred.79 The results of this work have consequences for understanding how HONO emissions from NO3¯ photolysis can be enhanced in the presence of ubiquitous OH scavengers (e.g., organic matter, HCO3¯, and CO32−) present on boundary layer

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ASSOCIATED CONTENT

* Supporting Information S

Details of procedures used to determine aqueous N(III) concentrations, derivation of gas phase NO2 and HONO concentration, and a description of the two-film model. This material is available free of charge via the Internet at http:// pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

* E-mail: JDRaff@indiana.edu; Phone: +1 (812) 855-6525. 11997



Article

Present Address

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†

A.E.B.: Department of Chemistry, Smith College, Northampton, Massachusetts 01063, United States. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was funded by Indiana University and the National Science Foundation (CAREER Award, AGS-1352375 to J.D.R.). N.K.S. was supported by an EPA STAR Graduate Fellowship. The authors thank Prof. Benny Gerber and Dr. Mychel Varner at the University of California−Irvine for many stimulating discussions regarding this work. We are also grateful to the instrument makers of Indiana University’s Edward J. Bair Mechanical Instrument Services, John Poehlman for electronics expertise, and Don Garvin for glassblowing services.

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