Kinetics Study of Heterogeneous Bromine Release from the Reaction

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Kinetics Study of Heterogeneous Bromine Release from the Reaction Between Gaseous Ozone and Aqueous Bromide Solution Yosuke Sakamoto, Motoki Goda, and Jun Hirokawa J. Phys. Chem. A, Just Accepted Manuscript • DOI: 10.1021/acs.jpca.7b12819 • Publication Date (Web): 26 Feb 2018 Downloaded from http://pubs.acs.org on February 27, 2018

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The Journal of Physical Chemistry

Kinetics Study of Heterogeneous Bromine Release from the Reaction between Gaseous Ozone and Aqueous Bromide Solution Yosuke Sakamoto†,‡,ǁ**, Motoki Goda§ and Jun Hirokawa†,§,* †

Faculty of Environmental Earth Science, Hokkaido University, Sapporo 060-0610, Japan

‡

Research Fellow of the Japan Society for the Promotion of Science

§

Graduate School of Environmental Science, Hokkaido University, Sapporo 060-0610, Japan

*

To whom correspondence should be addressed.

E-mail: [email protected] Tel / Fax: +81-11-706-4528 **

To whom correspondence should be addressed.

ǁ

Present address:

· Graduate School of Global Environmental Studies, Kyoto University, Yoshida-Honmachi, Sakyo-ku, Kyoto 606-8501, Japan · Graduate School of Human and Environmental Studies, Yoshida-nihonmatsu-cho, Sakyo-ku, Kyoto 606-8501, Japan · Center for Regional Environmental Research, National Institute for Environmental Studies, Tsukuba City, Ibaraki 305-8506, Japan E-mail: [email protected]

1 Environment ACS Paragon Plus

The Journal of Physical Chemistry 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

ABSTRACT The heterogeneous release of molecular bromine, Br2, from the reaction between gaseous ozone and aqueous bromide ion in sea water ice and sea salt aerosols is considered to be an initial source of reactive bromine species in the troposphere. Recent studies have demonstrated that the uptake of ozone by aqueous bromide solution is promoted by reactions at the gas-liquid interface. The present work investigated the heterogeneous reaction between gaseous ozone and aqueous bromide solution at atmospheric pressure and room temperature using a wetted wall flow reactor combined with a chemical ionization mass spectrometer. The emission rate of Br2 was measured as a function of gaseous ozone concentration, aqueous bromide concentration, and pH. Also, we conducted a simple kinetics model simulation that included only bulk aqueous-phase reactions and compared the theoretical values with the experimentally determined values. The Br2 emission rates measured experimentally differ from the simulated rates at relatively high bromide concentration, as well as in the pH region of 6–9. These differences might be explained by different Br− concentration and/or de-protonation efficiency near the interface region and those in the bulk solution.


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INTRODUCTION Reactive bromine species are known to significantly impact tropospheric ozone depletion in the polar regions and in the marine boundary layer.1-3 Scheme 1 shows an example of gas-phase ozone depletion cycles in the remote marine boundary layer.

Scheme 1 An example of gas-phase ozone depletion cycles in the remote marine boundary layer by reactive bromine species. 2-3

These catalytic cycles are initiated by daytime photolysis of photolyzable bromine compounds such as molecular bromine, Br2. Heterogeneous reactions involving bromide ions in the aqueous phase are believed to enhance reactive bromine species through autocatalytic reactions4 and may also play a role in the initial release of photolyzable bromine compounds. Several groups have proposed that heterogeneous reactions of gaseous ozone with bromide ions in sea-salt particles and sea water ice provide a dark source of molecular bromine.5-7



The release of molecular bromine from the reaction between gaseous ozone and aqueous bromide ions broadly comprises three processes involving reactants and products: gas-phase diffusion to/from the liquid surface, mass transfer (accommodation) at the gas-liquid interface, and bulk aqueous-phase reactions. Following the uptake of ozone, the mechanism of formation of molecular bromine via bulk aqueous-phase reactions of ozone and bromide in the absence of chloride can be described by the following proposed reactions:8-11

O3 + Br− → BrOOO−

(R1)

BrOOO− → O3 + Br−

(R2)

BrOOO− + H+ → HOBr + O2

(R3)

BrOOO + H2O → HOBr + OH −

−

(R4)

HOBr + Br− + H+ → Br2 + H2O

(R5)

HOBr + Br− + H2O → Br2 + OH− + H2O

(R6)

Br2 + H2O → HOBr + Br + H

(R7)

−

+

Br2 + OH− + H2O → HOBr + Br− + H2O

(R8)

Br2 + Br− → Br3−

(R9)

Br3− → Br2 + Br−

(R10)

where all species are in the aqueous phase. As shown by these reactions, the formation of bromine in the aqueous phase requires a proton, i.e., the reaction depends on pH. Observations and theoretical arguments have suggested that reactions at the gas-liquid interface, which are promoted by the propensity of ozone to reside at the interface and by the surface enhanced concentration of bromide,8 significantly contribute to heterogeneous Br2 formation12 and release8, 13 in addition to the bulk aqueous-phase reactions. Hunt et al. reported that Br2 production by the heterogeneous reactions of gaseous ozone with aqueous bromide particles at pH ~6.6 cannot be explained only by bulk aqueous-phase chemistry, but can be


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reproduced with the following reactions (R11)–(R13) which are tentatively proposed as direct Br2 release by surface reactions:8, 14

O3 (g) + Br− (interface) → O3Br−(interface)

(R11)

O3Br (interface) →→ O3 + 1/2 Br2(g)

(R12)

O3 + H2O → OH + O2 + OH

(R13)

−

−

−

−

where O3(g) and Br2(g) are ozone and molecular bromine in the gas phase, and Br−(interface) and O3Br−(interface) represent bromide ion and an intermediate complex at the gas-liquid interface, respectively. Surface reactions (R11)–(R13) produce OH−, resulting in a pH increase. Clifford and Donaldson reported that the rate of pH increase at the surface of a NaBr solution exposed to gas-phase ozone depends on the “surface” ozone concentration controlled by octanol addition and on the “surface” bromide concentration controlled by Langmuir-Hinshelwood-type kinetics.12 Measurements of uptake rates have also suggested the importance of surface reactions. Using a horizontally-aligned tubular flow reactor, Oldridge and Abbatt13 measured the release rate of molecular bromine from the heterogeneous uptake of gaseous ozone by an aqueous bromide/chloride solution kept at 273 K in the pH range of 2−6, and found that the uptake coefficient shows a Langmuir-Hinshelwood-type dependence on the gaseous ozone concentration.13 On the other hand, Lee et al. measured the uptake coefficient from the rate of loss of gaseous ozone using a rectangular flow reactor at 289 K and found that the uptake of ozone by neat NaBr solution can be explained by bulk aqueous-phase reactions, despite the possibility that a surface reaction may be involved due to uncertainties associated with the kinetic parameters used.15



Very recently, Artiglia et al.16 used the same rectangular flow reactor as that used in Lee et al.15 at 273–278 K and reported that surface reactions can contribute to uptake rate. They found that the uptake coefficient is inversely dependent on the ozone concentration at ambient levels, consistent with the findings of Oldridge and Abbatt at 273 K.13 Artiglia et al. suggest that surface reactivity is due to the stability of a BrOOO− intermediate at the gas-liquid interface rather than to the surface preference of bromide ion. Thus, the contribution of surface processes to the heterogeneous uptake of ozone by bromide solution remains inconclusive. The discrepancies in results obtained to date may be due to differences in experimental conditions, including temperature. Furthermore, Lee et al. and Artiglia et al. determined uptake coefficients from the rate of loss of gaseous ozone15,16 while Oldridge and Abbatt measured the production rate of Br2 in the gas phase.13, 15-16 As shown in (R1)–(R10), the production of Br2 requires several steps following the uptake of ozone, and its production rate will show different kinetics that may be dependent on the concentrations of H+ and Br−. Surface-specific processes may also affect these subsequent reactions. From an atmospheric chemistry viewpoint, quantification of the emission rate of gaseous Br2 from the reaction between ozone and bromide solution is required to evaluate the impact of reactive bromine species in the troposphere. Here, we investigated the release of molecular bromine from the heterogeneous reaction between gaseous ozone and aqueous bromide solution at a wide pH range, from 1.7 to 11.2, at 295 K under atmospheric pressure using a vertically-aligned wetted wall flow reactor combined with a chemical ionization mass spectrometer for the detection of Br2 in the gas phase. We obtained the release rate of molecular bromine as a function of the gaseous ozone concentration, aqueous bromide concentration, and pH. We also conducted a simple kinetics model simulation


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that included only bulk aqueous-phase reactions, then compared the experimental and simulated results in order to explore the contribution of surface reactions as identified from differences between the experimental and simulated results.

EXPERIMENTAL Wetted Wall Flow Reactor Experiments were conducted using a wetted wall flow reactor consisting of a vertically aligned Pyrex glass tube with an internal radius, a0 , of 1.0 cm and a length of 80 cm, as described previously.17 The inner wall of the flow reactor was covered with a thin film of pH-adjusted aqueous NaBr solution flowing downward under the effect of gravity. The aqueous solution was prepared in a closed glass vessel, in which the solution was pressurized at ~1.2 atm with N2 and supplied to the flow reactor by the pressure difference. The liquid volume flow rate, F, was regulated by a needle valve before entering the flow reactor at 0.50 cm3 s−1, which gives a film thickness, δ, of 130 µm and a mean liquid flow velocity, 〈νliq〉, of 6.1 cm s−1 according to the following equations:18-20 1

 3µF  3  δ =   2πa0 ρg 

vliq =

ρgδ 2 3µ

(1)

(2)

Here, µ and ρ denote the viscosity and density of water (0.0096 g cm−1 s−1 and 0.998 g cm−3 at 295 K), respectively, and g denotes gravitational acceleration (980 cm s−2).21 The flow reactor



was rinsed with deionized water before and after each experiment. As shown in Supporting Information, the velocity at the gas-liquid interface, vint, was estimated to be 9.2 cm s−1. A flow of He carrier gas was introduced into the flow reactor after being humidified by passage through a water bubbler in order to prevent the evaporation of water from the liquid film. The gaseous reactant, O3, was produced by irradiating pure air with 185 nm light from a Hg lamp (Jelight Co. Inc., Model 610). A flow of ozone/pure air was introduced through a movable injector with an internal diameter of 0.6 cm inserted into the flow reactor in a coaxial configuration. The O3 thus introduced was allowed to interact with the flowing liquid surface. The reaction distance, z, was changed in the range of 15 to 40 cm by moving the injector position. The flow rates of the He carrier gas and the ozone/pure air reactant gas were controlled at 3400 and 300 STP (standard temperature and pressure) cm3 min−1, respectively, by calibrated thermal mass flow controllers (Kofloc Inc., Model 3650 and 3660). The mean gas flow velocity, 〈vg〉, is estimated to be 21 cm s−1 from the total flow rate. The reaction time, t, is obtained by dividing z by 〈vg〉.

t=

z vg

(3)

A part of the gas flow after the reaction region was extracted through a 0.25-mm-diameter orifice into a chemical ionization mass spectrometer (CIMS) for measurement of the Br2 concentration, as described below. Another part of the gas flow was introduced into an ozone monitor (Dylec, Model 1150), and the remainder was exhausted to a fume hood. The flow reactor was operated at around atmospheric pressure. Although the temperature of the flow


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reactor was not regulated, all experiments were conducted in a temperature-controlled room at 295 ± 3 K.

CIMS Details regarding the CIMS setup were described previously.22 The gas sampled from the wetted wall flow reactor was mixed with an argon flow containing the reagent ion, SO2Cl−, and then Br2 in the sampled gas was ionized through a Cl− transfer reaction from SO2Cl−.

Br2 + SO2Cl− → Br2Cl− + SO2

(R14)

After the chemical ionization region, ions (including the produced Br2Cl− and unreacted SO2Cl−) were introduced into a high vacuum region and mass-analyzed using a quadrupole mass filter. Figure 1 shows a typical mass spectrum obtained when a standard Br2/He/air mixture was sampled into the CIMS. The standard mixture was prepared by diluting gaseous Br2 with a He/air mixture. Gaseous Br2 was generated using a permeation tube at a permeation rate of 3310 ng min−1 at 323 K (Gastec Co., P-10-H) placed in a temperature-controlled permeator (Gastec Co., PD-13).



3.0

Ion counts / 104 s-1

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

SO235Cl-

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Br2Cl-

2.5 2.0 1.5

191

193

195

197

199

SO237Cl-

1.0 0.5 0.0 90

100

110

190

200

m/z Figure 1. A typical CI mass spectrum for a standard gas with [Br2] = 150 ppbv in a He and air mixture. Inset shows an enlargement of the Br2Cl− peaks.

In the mass spectrum shown in Figure 1, the reagent ion, SO2Cl−, was observed as two peaks at m/z = 99 and 101 due to the stable chlorine isotopes

35

Cl and

37

Cl, with an isotope ratio of

approximately 3:1. Br2Cl− produced from Br2 through (R14) was observed as four peaks at m/z = 193, 195, 197, and 199 with a relative intensity of approximately 3:7:5:1 due to the stable isotopes of chlorine as well as bromine (79Br and 81Br; 1:1). The most intense peak at m/z = 195 was monitored for measurements of the Br2 concentration. The signal intensity of Br2Cl− at m/z = 195 was normalized to that of SO2Cl− at m/z = 101 to cancel out time variations in the ion signal intensities. Figure S1 in Supporting Information shows the normalized signal intensity at m/z 195 as a function of Br2 concentration, [Br2], of the standard Br2/He/air mixture described above. The


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normalized signal intensity is proportional to [Br2]. The slope of the plot corresponds to the detection sensitivity of the CIMS to the Br2 concentrations. Using the sensitivity thus obtained, the concentration of Br2 produced from the reaction between gaseous O3 and aqueous Br− was determined from the signal intensities measured by CIMS. We verified that the presence of water vapor does not affect sensitivity to Br2 in our CIMS measurements. In this study, the Br2 formation and release to the gas phase was also simulated using a simple chemical kinetics box model that included only bulk aqueous-phase reactions, whose details are available in Supporting Information, and compared with experimental results.

Preparation of Aqueous Solutions The aqueous solutions were prepared by dissolving solid NaBr (≥99.5%, Wako Chemicals) in deionized water. The pH of the solution was adjusted by adding dilute sulfuric acid (H2SO4) to pH = 1.7–3.0, potassium hydrogen phthalate buffer (KHP, KHC8O4H4) to pH = 4.2, potassium dihydrogenphosphate/disodium hydrogen phosphate buffer (PBS, KH2PO4 / Na2HPO4) to pH = 5.0–7.7, and sodium hydroxide (NaOH) to pH = 8.0–11.2. The pH of the solution without pH adjustment was 5.6–5.9, possibly due to impurities in the NaBr reagent. The pH of the prepared solution was measured by a pH meter (Horiba Ltd., Model F-53) before and after each uptake measurement. To remove dissolved gases, N2 was bubbled through the solution prior to each experiment.

Kinetics for data analysis



The removal of gaseous ozone, O3(g), via uptake is typically characterized by the uptake coefficient, γ, defined as:23

γ =

the net rate of O 3 (g) uptake the collision rate of O 3 (g) with surface

(4)

Using this uptake coefficient, the decay rate of O3(g) in the cylindrical reactor is described as:24

γω d[O3 ( g)] = − O3 [O3 (g)] dt 2a0

(5)

where ωΟ3 is the mean thermal velocity of O3(g) and is estimated to be 3.6 × 104 cm s−1 at 295 K. [O3(g)] is the O3(g) concentration averaged over the cross section of the flow reactor. As described, the reaction time was converted from the reaction length using eq. 3. In addition to γ, we defined the emission coefficient, ε, by the following equation in order to describe the release of gaseous bromine, Br2(g):

ε≡

the net rate of Br2 (g) release the collision rate of O 3 (g) with surface

(6)

Then, the emission rate of Br2(g) was described using the emission coefficient as below:


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d[ Br2 ( g)] εωO 3 = [O 3 (g)] dt 2a 0

(7)

Here, [Br2(g)] is the Br2(g) concentration averaged over the cross section of the flow reactor.

RESULTS AND DISCUSSION The filled squares in Figure 2 show an example of the ozone concentration measured as a function of reaction time.

5.0 4.0

[O3] / ppm

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Experiment Model simulation

3.0 2.0 1.0 0.0 0.0

0.3

0.6

0.9

1.2

1.5

1.8

Reaction time / s

Figure 2. An example of gaseous ozone concentration as a function of reaction time under the initial condition of [Br−] = 12.5 mM and [O3] = 2.6 ppm at pH = 2.1. The filled squares show experimental results and the dashed line shows model simulation results.



The ozone concentration remains essentially unchanged with reaction time, suggesting that the uptake coefficient of ozone is very small. Figure 2 also shows the model simulation results, with the initial ozone concentration adjusted to the experimental value. To allow comparison with the experimental results, the model results shown were averaged over the cross section of the flow reactor. Although the measured and simulated results are seen to agree well, the simulated results show a small decrease with reaction time. As shown by eq. 5, the slope of this decrease gives the uptake coefficient, γ. While the γ values could thus be obtained from the simulation, they were too small to be measured experimentally in the present study. In order to check the performance of our model simulation, we compared the γ values obtained from the simulation with those calculated based on the resistance model.23 When the interfacial process is ignored, the total resistance associated with ozone uptake can be represented by the sum of the partial resistances due to gas-phase diffusion, mass accommodation into the aqueous phase, and bulk aqueous-phase reactions. Because the uptake coefficient is the reciprocal of the total resistance, it is described in terms of the inverse of the resistance (the conductance) of each process, as follows:

1

γ

=

1 1 1 + + Γg α Γ b

(8)

where α is the mass accommodation coefficient, and Γg and Γb is the conductance for the gas-phase diffusion and bulk aqueous-phase reactions, respectively, both of which are


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normalized by the gas-surface collision rate. The value of α for O3 on the surface of aqueous halide solution was reported to be >0.1. 25 Γg can be approximated as the uptake coefficient in the extreme condition where the accommodation coefficient is unity and the bulk aqueous-phase reactions are very rapid, i.e., the ozone concentration is equal to zero at the surface. In this case, Γg is expressed as the following equation derived from eq. 5 and the analytical solution for the axial ozone profile in a tubular flow reactor, namely, the Gomeley-Kennedy (GK) solution:17, 26

Γg =

2 × 3.657 a 0ω O3 DO3,g

(9)

where DO3,g is the gas-phase diffusion coefficient of O3. Under the present experimental conditions, Γg is estimated to be 4 × 10−4. Γb is expressed by the following equation:18, 23

Γb =

4 H O3 RT DO3,aq k b II [Br − ]

(10)

ω O3

Here, HO3 and DO3,aq denote Henry's law constant and the aqueous-phase diffusion coefficient of O3, respectively, R is the gas constant, and kbII is the apparent second-order rate constant defined by the following equation.

k  k k1  3 [ H + ] + 4  k k2  d[O 3 ] II − =  2 [Br − ][O 3 ] = k b [ Br − ][O 3 ] k k dt 1 + 3 [H + ] + 4 k2 k2

(11)



where k1-k4 denote rate coefficients for (R1)-(R4). The diffuso-reactive length, l, is defined as the depth where the concentration of the dissolved gas decreases to 1/e of the concentration at the gas-liquid interface,23 and can be estimated using the following equation.

l=

DO3,aq

(12)

kb II [Br− ]

In our experimental conditions, the value of l is below ~40 µm which is sufficiently shorter than the solution thickness of 130 µm. Therefore, the eq. 10 can be applied in all conditions. Figure S2 of the Supporting Information shows the results of the comparison between the model simulation and the resistance model. The model simulation shows good agreement with the resistance model, with a small overestimation of about 4% due to the coarseness of the grid division along the radial direction. Both results show that the uptake coefficient is on the order of 10−7, indicating that the terms 1/α and 1/Γg are negligible compared to 1/Γb in eq. 8 and that the bulk aqueous-phase reactions are rate limiting, i.e., γ ≈ Γb. The filled squares in Figure 3 show an example of [Br2(g)] measured during the uptake experiments as a function of the reaction time.


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20 16

[Br2] / ppb

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


12

Experiment Linear fit to experiment Simulation Linear fit to simulation (> 0.72 s)

8 4 0 0.0

0.3

0.6

0.9

1.2

1.5

1.8

Reaction time / s

Figure 3. An example of gaseous Br2 concentration as a function of the reaction time for the initial condition [Br−] = 12.5 mM and [O3] = 2.6 ppm at pH = 2.1. Filled squares show the experimental results and the solid line shows the model simulation results. Dashed and dotted lines show linear fits to the experimental and model values obtained at reaction times greater than 0.72 s, which corresponds to a reaction distance of 15 cm.

The measured [Br2(g)] increases almost linearly with time, with a positive intercept with the horizontal axis. In Figure 3, the simulated [Br2(g)] under the same conditions as the experiment is shown as the solid line. The simulated line starts from 0 at t = 0 and first increases nonlinearly, but then increases linearly with increasing reaction time. Because the uptake rate is not limited by gas-phase diffusion, as described above, the concentration of gaseous ozone exiting from the movable injector will be mixed uniformly along the radial direction after a certain induction time, t0. Our model simulation shows that a reaction distance of 15 cm, which corresponds to a reaction time of 0.72 s, is required for uniform mixing, as shown in Figure S3(a) of the Supporting Information. During this induction period, the ozone concentration in the aqueous phase reaches a steady state, as shown in Figure S3(b) of the Supporting Information.



Because [O3(g)] remains essentially unchanged with time and, at t ≥ t0, O3(g) is uniformly distributed, [O3(g)] in the right-hand side of eq. 7 can be replaced by the initial ozone concentration, [O3(g)]0, and the following relation is obtained after integration:

[ Br2 (g)] =

εωO3 [O3 (g)]0 (t − t0 ) 2a0

(13)

Therefore, the slope of a linear fit to the measured [Br2(g)], as shown by the dashed line in Figure 3, experimentally provides the emission coefficient, ε. Similarly, the simulated [Br2(g)] at t ≥ t0 can be fitted by a linear regression line, as shown by the dotted line in Figure 3, to give the modeled ε for these conditions. In the examples shown in Figure 3, the regression lines for the experimental and simulated results appear to intersect with the horizontal axis at the same point, but in several experiments, they intersect at different points. This may be because the time required for gas-phase mixing in the experiments deviated from the time required in the simulation. Regardless, an experimental ε value could be obtained from the slope of the fitted line, and thus such discrepancy between the intercepts would not affect the results. As shown in Figure 3, the measured Br2(g) concentrations gave a steeper slope than the simulated concentrations, leading to a larger ε value in the experiment than in the simulation. Below, we compare the experimental and modeled ε values in the context of their dependence on [O3(g)], [Br−], and pH.


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Figure 4 shows the experimental and modeled ε values as a function of [O3(g)] (panel (a)) and [Br−] (panel (b)) at pH 2.1. There is uncertainty in the experimentally obtained ε values due to the linearity of the measured [Br2(g)] against the reaction time and the systematic errors associated with the determination of [Br2(g)], [O3(g)], and 〈vg〉. The overall uncertainty was estimated to be at most 10%.

3.0 a)

Uptake and emission coefficient / 10-7

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


2.0 ε (exp.) ε (sim.) γ (sim.)

1.0 0.0 0.0

1.0

2.0

3.0

4.0

[O3(g)] / ppmv 3.0 b) 2.0 ε (exp.) ε (sim.) γ (sim.) 20

1.0 0.0

0

10

[Br-] / mM Figure 4. Br2 emission coefficient as a function of gaseous ozone concentration at pH 2.1 with the initial condition of [Br−] = 15 mM (a) and as a function of [Br−] at pH 2.1 with the initial



concentration of [O3] = 3.8 ppmv (b). Solid and dashed lines show the simulated emission and uptake coefficient, ε and γ, respectively.

As shown in Figure 4, both the measured and modeled ε values showed no dependence on [O3] but showed nonlinear dependence on [Br–]. These behaviors of ε are primarily affected by the behavior of the uptake coefficient, γ. As described above, under the present experimental conditions, the overall uptake process is not limited by gas-phase diffusion and mass accommodation, but rather is governed by the bulk aqueous-phase reactions if reactions at the liquid surface are negligible. In this case, γ is approximated to the normalized conductance for the bulk aqueous-phase reactions, Γb, which is independent of [O3] but is dependent on [Br−]0.5, as shown in eq. 10. In Figure 4, γ values obtained in the simulation in which the surface reactions were not taken into account are shown as a dashed line. In agreement with eq. 10, the modeled γ values are independent of [O3] and increase with [Br−]0.5. From the above considerations, the measured ε values shown in Figure 4(a), which are independent of [O3], can be attributed to the [O3]-independent behavior of γ, which is governed by the bulk aqueous-phase reactions under our experimental conditions. This finding appears inconsistent with those of previous studies. Oldridge and Abbatt observed that the uptake coefficient is inversely dependent on [O3(g)] and suggested that the uptake rate of ozone by bromide solution is affected by the surface reactions in parallel with the bulk aqueous-phase reactions: namely, the overall uptake coefficient is represented as the sum of the conductance for the surface reactions, Γs, and that for the bulk aqueous-phase reactions, Γb.13


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γ = Γs + Γb

(14)

Very recently, Artiglia et al.16 observed a similar inverse dependence on [O3(g)], supporting the findings of Oldridge and Abbatt. Their arguments are based on the suggestion that the conductance

for

the

surface

reactions,

Γs,

can

be

described

according

to

the

Langmuir-Hinshelwood mechanism as follows: 13, 15

4 Kks II [Br − (s)] 4 Kks I Γs = = ω O3σ O3 (1 + K [O 3 (g)]) ωO3σ O3 (1 + K [O 3 (g)])

(15)

where K is the Langmuir adsorption equilibrium constant, ksII is the second-order rate coefficient for the reaction between O3(g) and Br− on the surface, [Br−(s)] is the concentration of Br− on the surface, and σO3 is the surface area occupied by one adsorbed O3 molecule. ksI on the far right-hand side of eq. 15 is the first-order rate coefficient, ksI = ksII [Br−(s)]. Equation 15 demonstrates that Γs depends inversely on [O3(g)] at K[O3(g)] >> 1, which is consistent with the dependence of γ observed by previous studies.13, 16 One difference between our studies and previous studies that reported the contribution of the surface reactions is temperature. Our studies were conducted at 295 K, which is much higher than the temperatures used for the studies by Oldridge and Abbatt (273 K) and those by Artiglia et al. (273–278 K). It should be noted that Lee et al.15 investigated the uptake of O3(g) by NaBr solution at 289 K and could not obtain clear evidence for the surface reactions.



Experimental conditions other than temperature may also contribute to the differences between the present and previous investigations. Equation 15 shows that Γs will approach zero at extremely high concentrations of O3(g). Artiglia et al. proposed temperature-dependent expressions for K and ksI to reproduce their observations at different temperatures, from 273 to 278 K. In addition, they estimated σO3 to be ~10−12 cm2 molecule−1 from their liquid-jet XPS measurements.16 (Note that they actually estimated the maximum surface coverage, Nmax, at ~1012 molecules cm−2, which is the reciprocal of σO3.) Using their parameters, Γs values at 298 K under our experimental conditions are calculated to be 5.3 × 10−7 and 1.0 × 10−7 at [O3(g)] = 0.7 and 3.8 ppmv, respectively, which are not negligible compared with the modeled γ (≈ Γb) value of 2.8 × 10−7 from the bulk aqueous-phase reactions only. However, Artiglia et al. assumed that ksI is independent of the bulk aqueous-phase Br− concentrations (ranging from 0.12–0.24 M in their studies), which are approximately one order of magnitude higher than those in our studies.16 If [Br−(s)] depended linearly on the bulk aqueous-phase bromide concentration, then Γs in the present study would be on the order of 10−8, which would be negligible compared to Γb. Thus, we can infer that we would not observe the inverse [O3(g)] dependence of γ due to the surface reactions because of the low bromide concentrations. In Figure 4, the experimental and modeled ε values are both smaller than the modeled uptake coefficient γ. This indicates that not all O3(g) molecules taken up by the bromide solution necessarily lead to the release of Br2(g). Our simulation demonstrates that a part of the Br2 formed in the bulk aqueous-phase reactions diffuses deeper and remains in the aqueous phase (see Figure S4 of the Supporting Information). Comparison of the experimental and modeled ε values shows that the experimental values are slightly higher than the modeled values, even after


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taking into account the uncertainties associated with the experimental and modeled ε values. In Figure 4(b), the difference between the experimental and modeled ε values appears to increase with [Br−] systematically. Thus, ∆ε (= the experimental ε − modeled ε) was plotted as a function of [Br−] in Figure 5.

6

4 ∆ε / 10−8

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


2

0

0

5

10

15

20

25

[Br−] / mM

Figure 5. A plot of the difference between the experimental ε value from the modeled ε value as a function of [Br−] at pH 2.1 with an initial concentration of [O3] = 3.8 ppmv.

As seen in this figure, a discernible ∆ε appears at a certain [Br−], increases almost linearly with [Br−], and levels off at high [Br−]. As described above, γ was not affected by the surface reactions under the present experimental conditions. According to Artiglia et al., the high preference of the intermediate BrOOO− species for the interface results in surface reactivity which could enhance γ .16 On the other hand, molecular dynamics (MD) simulations have



demonstrated the propensity of Br− for the gas-liquid interface,8, 27 which may affect the rates for the reactions succeeding the uptake of O3(g) because the subsequent reactions can also depend on the bromide concentration, as shown by reactions R5 and R6. The ∆ε values observed here may reflect this surface propensity of Br−. Although it is important to explore ∆ε in more detail, further analysis is difficult because of uncertainties associated with both the experimental and modeled ε values. Because H+ can be involved in the formation of Br2, we investigated the pH dependence of ε. Figure 6 shows the experimental and modeled ε values at [Br−] = 20 mM and [O3] = 3.8 ppmv as a function of pH.

4.0

Emission coefficient, ε / 10-7

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

w/ sulfuric acid w/ KHP w/ PBS w/o buffer w/ NaOH ε (sim.)

3.0

ε (modified sim.) γ (sim.)

2.0

1.0

0.0 2

4

6

8

10

pH

Figure 6. The Br2 emission coefficient as a function of pH at [Br−] = 20 mM and [O3] = 3.8 ppmv. Filled squares, an open square, filled circles, filled triangles and open circles show experimental results with sulfuric acid, KHP, PBS, NaOH, and without pH adjusters, respectively. Solid and dotted lines show modeled ε values and those modified with a triple k8/k6, respectively. Dashed line shows the simulated uptake coefficient, γ.


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As shown in Figure 6, the experimental ε values decrease gradually with an increase in pH in the pH range from 3 to 6, while rapid decreases with pH are seen at pH < 3 and 6 < pH < 9. The modeled ε values shown as a solid line show similar pH dependence, but the second rapid decrease appears at higher pH compared to the experimental data. Figure 6 also shows the modeled uptake coefficients as a dashed line which decrease with pH at pH < 3 but are almost unchanged at pH ≥ 3. This behavior of γ reflects the pH dependence of kbII, as shown in eq. 11, and thus the rapid decrease in ε at pH < 3 is also attributable to the change in kbII. At pH 6–9, on the other hand, ε decreases with pH while γ remains constant, which suggests that the pH dependence of ε in this region may be attributed to the effects of H+ on the reactions subsequent to the uptake of O3(g). The uptake of ozone leads to the formation of HOBr, Br2, and Br3−, whose partitioning is governed by following equations as shown in the Supporting Information.

[Brtot ] [HOBr] [Br3 − ] =1+ + [Br2 ] [Br2 ] [Br2 ]

(16)

[HOBr] k 7 + k8 [OH − ] = [Br2 ] k 5 [Br − ][H + ] + k 6 [Br − ]

(17)

[Br3 − ] k9 [Br − ] = [Br2 ] k10

(18)

where k5–k10 denote rate coefficients for (R5) – (R10).

In particular, the pH dependence in the

pH range 6–9 is largely determined by the magnitudes of the rate coefficients for (R5)–(R8).



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While the values of k5 and k7 have been obtained experimentally, the values of k6 and k8 were calculated from Brønsted-Pederson relationship for general-acid/base-assisted reactions.8,

10

Therefore, errors in the estimation of k6 and k8 would cause the observed discrepancy between the pH dependence of the experimental and modeled ε values. If the ratio of k8 to k6 is tripled, from the original value of 2.2 × 105 to 6.6 × 105, then modeled ε values that are closer to the experimental ε would be obtained, as shown by the dotted line in Figure 6. In order to investigate pH dependence, this study used a variety of pH adjusters, including sulfuric acid solution, potassium hydrogen phthalate buffer (KHP), potassium dihydrogenphosphate/disodium hydrogen phosphate buffer (PBS), and sodium hydroxide. Such pH adjusters might act as surfactants, preventing the uptake of ozone or pushing bromide ion away from the surface. However, fairly good agreements were found between the results obtained without the use of pH adjusters and those obtained with PBS around pH 6, and between the results obtained with PBS and those with NaOH around pH 8, showing that such surfactant effects, if present, are negligible. Another possible explanation for the discrepancy in pH dependence between the experimental and modeled ε values is the involvement of surface processes. Mishra et al. reported that hexanoic acid is deprotonated in much lower pH than its pKa of 4.8 at the air/water interface, implying OH− availability is more effective at the interface than in the bulk.28 This interfacial OH− availability may partly contribute to the discrepancy shown in Figure 6. Our simulation suggests that Br2 released into the gas phase is produced within ~20 µm depth from the surface into the solution (see Figure S4 of the Supporting Information). Therefore, because of the relatively effective OH− availability (and, consequently, the relatively high de-protonation


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efficiency) near the interfacial region, the experimental ε values may start decreasing at a slightly lower pH value and show rather more moderate pH dependence with an increase in pH compared to the modeled ε values.

ATOMOSPHERIC IMPLICATIONS Finally, we evaluated the importance of Br2(g) release from the heterogeneous reaction between O3(g) and aqueous bromide solution into the atmosphere. As shown above, most studies, including ours, suggest that surface reactions may contribute to the uptake of ozone and the release of molecular bromine at high Br− concentration and low O3(g) concentration. However, here we used a model that does not include surface reactions in the evaluation because the model formulates the bromine release rate and can thus be applied to different conditions. There are two principal types of surfaces where heterogeneous reactions between ozone and aqueous bromide can occur in the remote marine boundary layer: the ocean surface, and deliquesced sea salt aerosol surfaces. First, we estimated the emission coefficient under ocean surface conditions. As shown in Figure 6, the emission coefficient depends on pH and decreases dramatically in the pH 6–9 region. At a typical bromide concentration of ~1 mM at pH 8.2 in seawater,29 the equilibrium between HOBr and Br2 is inclined significantly towards HOBr. In this case, our model simulation estimates that ~0.3% of the O3 uptake leads to Br2 emission when the uptake and emission coefficients are estimated to be 7 × 10−8 and 2 × 10−10, respectively. Note that in this condition, HOBr emission might not be negligible compared with Br2 emission.



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Next, the emission coefficient for deliquesced sea salt aerosols was estimated. In this case, pH decreases to ~230 and the ion concentration increases by up to a factor of 103 (~ several M)8, 31 The volume distribution of typical atmospheric sea salt aerosols have a maximum at a radius of ~5 µm. Hanson et al. showed that when the particle radius is the same as or smaller than the diffuso-reactive length, which is defined as the depth where the concentration of the dissolved gas decreases to 1/e of the concentration at the gas-liquid interface, the reaction occurs throughout the particle volume, leading to reduction of the uptake coefficient.

32

They proposed

the following equation for correcting the uptake coefficient on small particles:32

 rp l  Γb,aerosol = Γb  coth −  l rp  

(19)

where rp is the particle radius and l is the diffuso-reactive length, which in this case is provided by the eq.12.23Assuming that the ions in the particles are ten-fold concentrated compared to sea water and have the same composition ratio, l is estimated to be 25 µm at 10 mM Br−. The correction factor can then be calculated and is 0.03 for a particle with a radius of 2.5 µm, which is representative of the surface area distribution of atmospheric particles. Since there are no analytical equations for emission coefficients from aerosols, we used our model simulation to estimate the ratio between emission and uptake coefficient. Although our model simulation was designed for an aqueous film on the inner surface of a tubular flow reactor, simulation with a film several µm thick qualitatively showed the effect of film thickness on Br2 emission from sea salt aerosols.


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Figure 7 shows the ratio of the Br2 emission coefficient, ε, to the O3 uptake coefficient, γ, calculated by the model simulation as a function of film thickness, δ, for different [Br−] values at pH 2.1 and [O3] = 3.8 ppmv.

Ratio of emission to uptake

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


1.0 0.8 0.6 0.4 0.2

1 mM 5 mM 10 mM 20 mM

0.0 0.000 0.002 0.004 0.006 0.008 0.010 0.012

Film thickness, δ / cm

Figure 7. Ratio of the Br2 emission coefficient to the O3 uptake coefficient calculated by model simulation as a function of film thickness, δ, at pH 2.1 and [O3] = 3.8 ppmv.

As shown in Figure 7, the fraction of O3 uptake leading to Br2 emission depends on the film thickness and is almost unity for films

Kinetics Study of Heterogeneous Bromine Release from the Reaction

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