Heterogeneous Interactions of OH and HO2 Radicals with Surfaces

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J. Phys. Chem. 1996, 100, 2249-2254

2249

Heterogeneous Interactions of OH and HO2 Radicals with Surfaces Characteristic of Atmospheric Particulate Matter P. L. Cooper and J. P. D. Abbatt* Department of the Geophysical Sciences, The UniVersity of Chicago, 5734 South Ellis AVenue, Chicago, Illinois 60637 ReceiVed: July 26, 1995; In Final Form: October 9, 1995X

The heterogeneous interactions of OH and HO2 radicals with a number of surfaces characteristic of atmospheric particulate matter have been studied by using a low-temperature flow tube coupled to a resonance fluorescence detector. In particular, the mass accommodation coefficients (R) of both OH and HO2 on supercooled sulfuric acid solutions have been measured: for OH, R > 0.2 for 45-96 wt % solutions from 220 to 295 K, and for HO2, R > 0.2 on 55 wt % solutions doped with 0.1 M CuSO4 at 223 K. Radical uptake coefficients (γ) were also measured on a variety of solid surfaces prevalent in the atmosphere: water ice, NH4HSO4 and (NH4)2SO4. On water ice, although it was found that the uptake coefficients of both OH and HO2 were relatively small on conditioned surfaces (γOH ) 0.03 ( 0.02 from 205 to 230 K, γHO2 ) 0.025 ( 0.005 at 223 K), it was observed that the OH uptake coefficient could be significantly increased by either adsorbing HNO3 to the surface or melting the ice surface by exposure to relatively high partial pressures of HCl. Similarly, on conditioned (NH4)2SO4 and NH4HSO4 surfaces at room temperature, the OH uptake coefficient is relatively small, γOH < 0.03. Nevertheless, the uptake coefficient can be significantly increased (γOH > 0.2) by exposing the surface to relatively low partial pressures of an organic species, 1-hexanol.

Introduction Motivated in large part by the depletion of stratospheric ozone, numerous studies of the heterogeneous reactivity of relatively long-lived atmospheric gas-phase species such as ClONO2, HCl, and N2O5 have been recently performed. At the same time, because the gas-phase lifetimes of many radical species are somewhat shorter than typical collision times with atmospheric aerosols, relatively little attention has been given to similar interactions which arise between radicals and surfaces characteristic of atmospheric aerosols. Nevertheless, a number of modeling studies have indicated that specific free-radical interactions may be atmospherically important processes under conditions of high aerosol surface area and facile heterogeneous loss. For example, in the troposphere, transfer of OH and HO2 from the gas-phase to cloud droplets, sea salt aerosols, and sulfate aerosols represents both a significant sink of gas-phase HOx radicals and a source of radicals to the aqueous phase.1-5 Similar conclusions have also been reached with respect to the NO3 radical.6 Also, it has yet to be determined whether freeradical heterogeneous chemistry plays a significant role in determining the photochemical steady state in high particulate surface area regions such as polluted urban environments and cirrus clouds. In addition to their atmospheric importance, the second motivation for studying radical-surface interactions arises from the potential role free-radical heterogeneous chemistry may play as an indirect probe of the composition and reactive character of surfaces prevalent in the atmosphere. For example, considering that OH is highly reactive in the gas-phase with molecules such as HCl, HNO3, and numerous organics, it is informative to inquire as to whether OH maintains this reactivity when these molecules are adsorbed to a surface. Using the low-temperature flow tube technique in this work, we report studies of the interactions of OH and HO2 radicals with a variety of surfaces typical of atmospheric aerosol X

Abstract published in AdVance ACS Abstracts, January 1, 1996.

0022-3654/96/20100-2249$12.00/0

composition. Specifically, we have measured the net uptake of OH by supercooled sulfuric acid solutions of 40-96 wt % composition and the uptake of HO2 by cold 55 wt % sulfuric acid solutions. To our knowledge, these are the first studies of HOx radical uptake by sulfuric acid solutions of stratospheric composition. We have also investigated the behavior of HOx radicals toward a variety of solid surfaces such as water ice, (NH4)2SO4, and NH4HSO4, all common components of tropospheric clouds and aerosols. To determine whether the presence of surface-adsorbed species enhances the reactivity of a relatively unreactive surface (such as water ice or dry (NH4)2SO4) toward gas-phase OH, we also report the interactions of OH with water ice surfaces exposed to HCl and HNO3, and with (NH4)2SO4 surfaces exposed to 1-hexanol. Throughout this paper we denote the net probability that a molecule will be lost from the gas-phase upon collision with a surface by an uptake coefficient, γ. The mass accommodation coefficient, R, refers to the probability that upon collision with a liquid surface, a gas-phase molecule will be taken into solution. Because liquid surfaces may readily saturate, leading to desorption of dissolved molecules and a reduction in the observed loss rate of the gas-phase species, the experimentally determined uptake coefficient is expected to be smaller than or equal to the mass accommodation coefficient, i.e., γ e R. A few experimental studies of free-radical heterogeneous chemistry have been performed previously. The loss of OH on sulfuric acid surfaces has been examined before in three studies: Baldwin and Golden7 report a low reactivity (γOH ) 5 × 10-4, where γOH is the net uptake coefficient of OH) on 96 wt % sulfuric acid at room temperature, whereas Hanson et al.8 and Gershenzon et al.9 report substantially larger uptake coefficients (γOH > 0.08 on 28 wt % acid at 249 K and γOH ≈ 1 on 96 wt % acid at room temperature, respectively). Gershenzon et al. have also reported a large loss coefficient for OH on ice surfaces (γOH g 0.4) at 253 K. In the case of HO2 radicals, Mozurkewich et al.10 have measured γHO2 to be larger than 0.2 on concentrated NH4HSO4 and LiNO3 solutions which © 1996 American Chemical Society

2250 J. Phys. Chem., Vol. 100, No. 6, 1996 have been doped with CuSO4, and Hanson et al.8 have measured γHO2 > 0.05 on 28 wt % sulfuric acid solutions at 249 K. Finally, using a room-temperature flow tube and resonance fluorescence detection of OH, Jech et al.11 studied hydroxyl radical interactions with a wide variety of solid substrates commonly found in tropospheric aerosols. A general conclusion from this work was that while OH was lost quite readily (γOH > 0.1) on transition metal-containing surfaces (e.g., Pb(NO3)2 and FeSO4‚nH2O) and with some organic surfaces (e.g., malonic acid), OH was quite unreactive (γOH < 0.01) toward more atmospherically-prevalent aerosol surfaces such as NH4NO3 and (NH4)2SO4. Experimental Section Uptake coefficients of gas-phase radicals were determined by monitoring their first-order loss within a low-temperature, low-pressure flow tube which has its inner walls coated with either a liquid or solid substrate. The first-order decay is linearly proportional to the net uptake coefficient for relatively inert surfaces, whereas for facile loss at the wall, the rate of gasphase diffusion to the wall significantly limits the overall radical loss rate and a correction is applied to the observed first-order rate constant.12 As an example, for γ’s of 0.1, the size of this correction is a factor of 3, whereas it is only 10% for γ’s of 0.01. In order to make these corrections, the following binary diffusion coefficients were used: 0.035T1.75 Torr cm2/s for OH in He (derived by applying a T1.75 dependence13 to the experimentally-determined diffusion coefficient of O atoms in He at room temperature14), 0.0063T1.75 Torr cm2/s for OH in H2O (derived by applying a T1.75 dependence to the 275 K diffusion coefficient calculated in reference 8) and 344 Torr cm2/s for HO2 in He at 223 K (the value for O2 in He at 223 K)13. To account for diffusion through both helium and water vapor, which together constitute the flow-tube carrier gas, the individual binary diffusion coefficients were combined into an overall diffusion coefficient through the buffer gas.8 Because only a few milliTorr of water vapor were prevalent in the flow tube for the majority of studies conducted here, the radical H2O contribution to the overall diffusion coefficient was less than 10%. Given the uncertainties in the diffusion coefficients ((15%) and the random errors associated with measuring fast, first-order decays, uptake coefficients greater than 0.2 cannot be measured accurately by this approach. The lowest uptake coefficients were 10-3-10-4, as measured on walls coated with halocarbon wax. Typical uncertainties (which take into account both random and systematic errors) for γ ≈ 0.01 were (15%, whereas the errors were somewhat larger for larger uptakes (e.g. (40% for γ ≈ 0.1). Radicals were generated within a 19-mm-o.d. axial injector which has a microwave-induced plasma mounted at the upstream end. To ensure that the injector output of radicals was constant as the injector was pulled out from a cold flow tube into the warm room environment, the injector was double jacketed with a room temperature ethylene glycol aqueous solution circulating through the inner jacket, thereby reducing the rate of radical wall loss within the injector. The outer jacket is evacuated and acts to keep the outermost walls of the injector cool so as to not significantly warm the reaction substrates. Typical readings from thermocouples mounted close to the reactor wall indicate 2 deg of warming induced by the injector at a position 2 cm downstream of the injector tip and 4 deg of warming when the injector is positioned over the thermocouple. Given the relatively low affinity shown by HF toward surfaces,15 the preferred method for the generation of OH radicals for the majority of experiments was via the titration

Cooper and Abbatt reaction F + H2O f HF + OH. For this source, F atoms were formed by passing trace amounts of CF4 through the microwave plasma, and H2O was added in known amounts immediately downstream of the plasma by bubbling 20 f 50 sccm of He through a water trap. For a few experiments, the H + NO2 f OH + NO source of OH radicals was also used. In this case, H atoms were generated by passing H2 through the microwave discharge. HO2 radicals were generated by the F + H2O2 f HO2 + HF titration reaction, with H2O2 being added to the injector flow by bubbling a small flow of He (50 sccm) through a concentrated H2O2/H2O solution. The hydrogen peroxide solution was concentrated to greater than 80 wt % by bubbling dry argon through a 50 wt % solution for a couple of days. For the reactant concentrations used, the titration reactions were driven well to completion within the 10-ms residence time in the injector. Radical concentrations used in this work were less than 5 × 1010 molecules/cm3. Typical conditions in the flow tube were 1500 f 2200 cm/s flow velocity at a total He buffergas pressure of approximately 1 Torr. An OH resonance fluorescence system was used to detect both OH and HO2 radicals. Specifically, the fluorescence from OH radicals in the flow, excited by 309-nm radiation from an RF-induced He/H2O plasma, was imaged by a series of baffles and quartz lenses onto a photon-counting, bialkali photomultiplier tube (Hammamatsu R2801) sitting at right angles to the OH lamp. As determined by calibration with the H + NO2 f OH + NO titration reaction with H atoms in excess, the sensitivity of this system is approximately 1 × 10-8 (count/s)/ molecule/cm3). This sensitivity corresponds to a detection limit (S/N ) 1) of 3 × 108 molecules/cm3 for a 5-s integration time and 50 background counts/s. HO2 was detected by chemical conversion to OH by the addition of NO immediately upstream of the resonance fluorescence axis: HO2 + NO f OH + NO2. As part of a 15 sccm He sweep gas, NO was added through a 3-mm-o.d. tube which spans the flow tube and has five small injection ports. For the 5-ms reaction time between the NO injector position and the OH detection axis, and for typical NO concentrations of 3 × 1013 molecules/cm3, we calculate that 80% of the HO2 in the flow was converted to OH. To facilitate easy preparation and substitution of different surfaces within the flow tube, the studies were performed on surfaces coating the inner wall of a 2.62-cm-i.d. Pyrex reaction tube which fits snugly within the flow tube itself. To prepare a surface, the inner walls of the reaction tube were first washed with a dilute HF solution, then thoroughly rinsed with distilled water, and then totally wetted with a solution of sulfuric acid (for studies on sulfuric acid surfaces), water (for studies on ice), or a salt solution (for studies on dry salt surfaces). The reaction tube was then quickly inserted into the cold flow tube, the flow tube was sealed from the atmosphere, and the surface was allowed to thermally equilibrate with the flow tube walls. In the case of sulfuric acid, the high viscosity of the film maintains a liquid layer a few tenths of a millimeter thick over the entire tube surface during the course of the experiment (up to 30 mins). For ice, the water film quickly freezes forming a highly transparent, smooth ice surface. In order to prevent net evaporation of water from the films, gas-phase water was added at the upstream end of the flow tube so that the water partial pressure in the flow tube matched the water vapor pressure of the sulfuric acid16 and ice films. For the studies on dry salt surfaces (i.e., (NH4)2SO4, NH4HSO4), a relatively concentrated aqueous solution (≈20 wt %) was used to coat the reaction tube. The water from this solution was evaporated from the film by flowing dry helium over the room temperature surface for periods of a few hours. As determined from the overall pressure

Interactions of Radicals with Surfaces

Figure 1. First-order decay of the OH signal (in arbitrary units) on a 50 wt% sulfuric acid surface at 230 K. Flow tube conditions: pressure ) 1.00 Torr; velocity ) 1800 cm/s.

reading, the partial pressure due to H2O evaporating from the film was less than a milliTorr at this point. We made no attempt to further quantify the water partial pressure of the film because the OH source used for the majority of these experiments, F + H2O f HF + OH, delivered a few milliTorr of water vapor to the flow tube itself. For the most part, the dry films totally coated the reaction tube as a somewhat transparent film. Commercially available gases were used without further purification: He (99.999%); CF4 (99.99%); NO2 (99.5%); NO (99.0%); HCl (99.0%); H2 (99.99%). Dilute reservoirs of HNO3 were prepared by mixing HNO3 vapor taken from 3:1 H2SO4: HNO3 solutions with helium. 1-Hexanol (Kodak) was added to the flow tube by bubbling small flows of He through a bubbler containing 1-hexanol. NH4HSO4 solutions were prepared by mixing (NH4)2SO4 (Aldrich) and H2SO4 solutions in a 1:1 molar ratio. Results and Discussion 1. Supercooled Sulfuric Acid Surfaces. For all experiments conducted on sulfuric acid solutions, the OH uptake coefficients were as large as can be measured by this technique, i.e., γOH > 0.2. The loss of OH was always irreversible and first order in nature as demonstrated in Figure 1, which shows a typical decay of OH on a supercooled sulfuric acid solution, in this case over a 50 wt % solution at 230 K. These experiments were conducted both on solutions typical of stratospheric composition (45-65 wt % at 220 to 230 K) using the F + H2O f HF + OH source and on sulfuric acid stock solutions (≈96 wt % at both 230 and 298 K) using the H + NO2 f OH + NO source. The H/NO2 source was used for the experiments on 96 wt % solutions in order to keep the water content of the film as small as possible. As has been suggested previously,8 it is expected that the irreversible loss of OH from the gas-phase is due to reaction between OH and bisulfate ion in solution, OH + HSO4- f H2O + SO4-, a reaction known to proceed quite rapidly in aqueous solutions at room temperature (kII ) 1 × 106 M-1 s-1).17 Although it is known that the sulfate radical reacts slowly with water to reform OH, SO4- + H2O f OH + HSO4-,18 for the OH concentrations used in this work it is likely that the ultimate fate of the SO4- radical in solution is the sulfate radical self-reaction, SO4- + SO4- f S2O82-, which proceeds in dilute aqueous solutions at the diffusion limit.19 Finally, in solutions more concentrated that 95 wt%, it has been estimated that the concentration of H2SO4 exceeds that of HSO4-.20 Consequently, we cannot rule out that a direct reaction between OH and H2SO4 in such solutions is leading to OH loss.

J. Phys. Chem., Vol. 100, No. 6, 1996 2251 These results are in accord with the flow tube measurements of Hanson et al.,8 who placed a lower limit of γOH > 0.08 in an experiment conducted over 28 wt% acid at 249 K, and Gershenzon et al.,9 who reported unity loss on highly concentrated sulfuric acid solutions at room temperature. Because our experiments were conducted with relatively low partial pressures of water (PH2O < 3 × 10-2 Torr), the gas-phase diffusion limitation for OH loss was dominated by diffusion through helium. As a result, we are able to put a higher lower limit on γOH in this work than were Hanson et al.,8 who had significantly higher water partial pressures which more greatly limited the rate of OH gas-phase diffusion to the wall. It appears as though our results are at odds with those of Baldwin and Golden,7 who reported a low uptake coefficient (γOH ) 5 × 10-4) on 96 wt% acid at room temperature. However, in order to reduce the water partial pressure in the Knudsen cell, the samples in these experiments were pumped at 10-7 Torr, which may have changed their composition from that of a stock, 96 wt% solution to one somewhat more concentrated. Because solutions close to 100 wt % H2SO4 may have HSO4- concentrations orders of magnitude smaller than those in 96 wt% solutions,20 their reactivity toward OH may be correspondingly lower. The first-order loss of HO2 radicals on 55 wt% sulfuric acid solutions was studied at 223 K. For pure, undoped solutions, an average value of γHO2 was determined from four experiments to be 0.055 ( 0.020, a value consistent with the lower limit of Hanson et al.8 (γHO2 > 0.05) measured on 28 wt% solutions at 249 K. Although it is possible that the irreversible loss of HO2 is due to the self-reaction of HO2 in solution (kII ) 3 × 106 M-1 s-1 for aqueous solutions at room temperature),17 firstorder kinetics would not be expected for such a mechanism. Alternatively, a potential loss process for HO2 is reaction with HSO4-: HO2 + HSO4- f SO5- + H2O. Although no kinetic information are available for the reaction, it may proceed quite readily given the high HSO4- concentration in sulfuric acid solutions. The liquid-phase reaction can be shown to be thermoneutral at 298 K by using literature standard heats of formation for SO5-(aq),21 H2O(aq), and HSO4-(aq) and by estimating the heat of formation of HO2(aq) to be -9 kcal/mol at 298 K by using thermodynamic values from ref 4 and by taking the entropy of solvation of HO2 to be -25 cal/(mol K). To test whether surface saturation and subsequent desorption of HO2 from solution was limiting the net HO2 uptake coefficient, we followed the approach of Mozurkewich et al.10 by doping our sulfuric acid solutions with CuSO4. Specifically, Mozurkewich et al. have measured large HO2 uptake coefficients, γHO2 ) 0.40 ( 0.21, on concentrated NH4HSO4 solutions doped with approximately 0.06 m Cu(II), whereas only minor uptake was observed on pure solutions. The enhancement in the uptake coefficient is thought to be due to the reaction sequence HO2 + Cu2+ f H+ + O2 + Cu+ followed by HO2 + Cu+ + H+ f H2O2 + Cu2+ proceeding in solution and leading to an irreversible loss of HO2. In our case, we observed a similar enhancement in the uptake coefficient for doped solutions: γHO2 > 0.2 for 55 wt% solutions with 0.1 M CuSO4 at 223 K. We conclude that the HO2 mass accommodation coefficient into sulfuric acid solutions of stratospheric composition is at least this large. 2. Water Ice Surfaces. OH radicals were lost irreversibly on ice surfaces with an uptake coefficient which is dependent upon the exposure time to OH. That is, while uptake coefficients on fresh ice surfaces were relatively large (γOH ≈ 0.1), it was observed that the surfaces become somewhat deactivated with time and that the uptake coefficients approached relatively low, steady values of 0.03 ( 0.02 after 10-20 minutes of

2252 J. Phys. Chem., Vol. 100, No. 6, 1996 conditioning with OH. Similar behavior was observed at temperatures from 205 to 230 K. By analogy with previous flow tube studies (e.g., refs 2225) of atom wall loss where atom-atom recombination is known to occur readily in a first-order manner, it is likely that the irreversible loss of OH is either due to recombination (OH + OH f H2O2) or to self-reaction (OH + OH f H2O + O) proceeding on the ice surface. Note that the first-order kinetics observed in this work are not in disagreement with such a mechanism if the rate-determining step for OH loss is transfer of OH from the gas-phase to the surface. For example, consider an Eley-Rideal mechanism26 where OH radicals readily bind to surface active sites (S) to form adducts (S-OH). If all the active sites are occupied by OH (i.e., the surface concentration of S-OH species is not dependent upon the gas-phase concentration of OH), the kinetics will be first order if the ratedetermining step is reaction of a gas-phase OH radical upon collision with an S-OH adduct. Noting for such a mechanism that the rate of OH loss is directly related to the number of active sites available to bond with OH, it is possible that the slow deactivation of the ice surfaces toward OH is due to a reduction in the number of active sites as the film ages. A number of experiments, all with null results, were performed which are consistent with OH loss not arising from reaction with impurities (such as HF or organic species) present on the ice film. Specifically, similar OH uptake coefficients were observed from two different radical sources (i.e., H/NO2 and F/H2O) and from three different water sources (tap water, laboratory deionized/distilled water, and commercially-available HPLC-grade water). Also, purification of the helium buffer gas by passing it through a trap at liquid nitrogen temperatures had no effect on the OH loss rate. Although the ice films may contain dissolved CO2 and O2, as a result of rapidly freezing liquid water, it is unlikely that such species will react with OH. By contrast to OH, the loss of HO2 on an ice surface does not show the same conditioning behavior, exhibiting uptake coefficients of γHO2 ) 0.025 ( 0.005 on ice surfaces at 223 K. As with OH, HO2 was lost irreversibly in a first-order manner on these surfaces. Using an ice surface which had been conditioned to OH, a number of experiments were performed to examine whether the loss of OH could be enhanced by the presence of adsorbed species on the ice surface. Specifically, both HCl and HNO3 were added in the gas phase to the flow tube in amounts at which significant adsorption to an ice surface is known to occur, leading in some cases to chemical transformations of the surface. The concentrations of HCl and HNO3 were sufficiently low ( 0.2). It is known for acidic aqueous solutions that OH reacts at the diffusion limit with chloride ions, OH + Cl- f Cl + OH-.17 While it is probable that this reaction is leading to the large OH uptake coefficients in the HCl/H2O liquid regime, it is also possible that the liquid surface is simply promoting the rate of the OH self-reaction. Nevertheless, it can be said generally, both from these observations and from the sulfuric acid experiments, that the mass accommodation coefficient for OH into concentrated acid solutions is very large, i.e., ROH > 0.2. For the HCl/H2O system, these experiments illustrate the contrast in surface structure for HCl partial pressures above and below the solid/liquid coexistence curve. In particular, the very large uptake of OH for high partial pressures is indirect confirmation that the surface has indeed melted under such conditions, whereas for low partial pressures the surface is clearly distinct from a liquid surface. Although uptake studies have indicated a substantial fraction of a monolayer uptake of HCl at low partial pressures, which by itself suggests an HCl bonding interaction stronger than simple physical adsorption,28 the low reactivity of OH toward such a surface indicates that HCl adsorption does not necessarily lead to behavior characteristic of a true liquid surface. Experiments were also performed using HNO3 as a reactive

Interactions of Radicals with Surfaces

J. Phys. Chem., Vol. 100, No. 6, 1996 2253

Figure 6. Ratio of OH signal to its original value as a function of time over dry (NH4)2SO4 (solid line) and NH4HSO4 (dashed line) surfaces at 296 K (see text for explanation). Figure 4. HNO3 vapor pressure as a function of temperature for the HNO3/H2O binary system (redrawn from ref 31). Dashed line represents the HNO3 vapor pressures scanned by the data in Figure 5.

Figure 5. OH uptake coefficients as a function of HNO3 concentration over an ice surface at 228 K. The transition from the water ice to the NAT thermodynamic regime occurs at [HNO3](g) ) 8 × 1011 molecules/ cm3. Uptake coefficients with a value of 0.2 are lower limits.

species known to adsorb to ice surfaces.29,30 These experiments were performed at a temperature of 228 K where scanning the HNO3 partial pressure allowed the transition from an ice regime to a regime where nitric acid trihydrate (NAT) is thermodynamically stable to be observed (see Figure 4). As the data in Figure 5 illustrate, the system behaves somewhat differently than the HCl system by exhibiting significant reactivity toward OH for partial pressures within the ice regime. That is, for HNO3 partial pressures below the coexistence curve between ice and NAT, there is a significant enhancement in γOH above the value measured for a clean ice surface. It has been shown for similar conditions that HNO3 adsorbs to ice surfaces,29,30 and we conclude that a reactive process is being promoted between OH and HNO3 adsorbed on the surface. Whether the ice surface is acting as an efficient third body and promoting the termolecular component of the OH + HNO3 f H2O + NO3 gas-phase reaction (which is thought to proceed through a HNO3-OH adduct)32 can only be suggested at present. Although the OH uptake coefficients are greater than 0.2 within the NAT stability regime at 228 K, experiments were also performed at a temperature of 200 K in order to better simulate stratospheric polar vortex conditions. In order to ensure that the concentration of HNO3 was in excess of that of OH, these experiments were performed at HNO3 concentrations greater than 3 × 1011 molecules/cm3, considerably higher than those prevalent in the stratosphere. For such conditions, exposure of ice to HNO3 raised γOH to the lower limit measurable in this experiment, i.e., γOH > 0.2. Considering that in prior experiments NAT was formed by exposing ice films

to nitric acid partial pressures comparable to those used in this work,29,33 we believe that NAT is present on the surface, probably in coexistence with ice. Nevertheless, mass spectrometric characterization of the NAT film and experiments at lower nitric acid partial pressures will be necessary to confirm that OH is highly reactive on a NAT surface under stratospheric conditions. 3. Dry Salt Surfaces. The interactions of OH with dry ammonium sulfate ((NH4)2SO4) and ammonium bisulfate (NH4HSO4) surfaces were studied. As described in the Experimental Section, the solid films were formed by drying the reaction tube, the inner walls of which had been totally wetted by an aqueous salt solution. The films formed in this manner are likely to be porous to some extent, and, thus, uptake measurements which use surface collision rates calculated by using the geometric surface area of the reaction tube would be expected to yield larger apparent γ’s than would experiments conducted on smooth films.34 Because we do not know the porosity and internal surface areas of these films, the goal of these experiments was not so much to quantify γOH as to qualitatively examine the OH uptake behavior on these surfaces as a function of time after initial exposure to OH. Figure 6 illustrates typical time-dependent behavior for two salt surfaces, (NH4)2SO4 and NH4HSO4, at 296 K. In the case of NH4HSO4, the steady-state OH signal at times less than 40 s is representative of the injector pushed in past the dry film. At 40 s, the injector is withdrawn and the OH signal drops after exposing 41 cm2 of the film to OH. With increasing time the OH signal does not significantly recover until the injector is pushed back to its original position at 675 s, at which point the OH signal returns to its original value. By assuming the film is smooth, we can calculate an upper limit for the OH uptake coefficient of γOH < 0.03 on this surface. It is interesting to contrast the behavior of a NH4HSO4 surface with that of a dry (NH4)2SO4 surface shown as the solid line in Figure 6. In this case, the film exhibits a rapid conditioning process so that by the time the injector is pushed back to its original position, the film has become highly unreactive, exhibiting a low value of γOH of 0.002 (again calculated assuming a smooth film). This result is in qualitative agreement with the experiments of Jech et al.,11 which also showed low reactivity of OH toward dry (NH4)2SO4 surfaces (γOH ) 0.0085). Although a conditioned (NH4)2SO4 surface is highly unreactive toward OH, its reactivity can be greatly increased by exposure to an organic species with a polar functional group. For example, a number of experiments were performed at room temperature where a deactivated (NH4)2SO4 surface was exposed to 1-hexanol in the gas-phase at concentrations which would not lead to a significant gas-phase reaction with OH. However, for concentrations greater than 2 × 1011 molecules/cm3, the OH uptake increases to γOH > 0.2 from a value of γOH < 0.01 on

2254 J. Phys. Chem., Vol. 100, No. 6, 1996 the deactivated (NH4)2SO4 surface. It is reasonable to suggest that the enhanced OH loss is due to reaction with surfaceadsorbed 1-hexanol. Since it has been suggested that organic surfactant coatings exist on tropospheric aerosols,35 further experiments to quantify the extent to which organic species are adsorbed on the surfaces of tropospheric aerosols and the effects that these layers may have on the rates of gas-to-solid/liquid transfer are currently under way in this laboratory. It is likely that the qualitative difference in reactivity between dry NH4HSO4 and (NH4)2SO4 surfaces after long OH conditioning times is due to the presence of the HSO4- ion in the NH4HSO4 film. Whether a gas-solid ionic reaction is proceeding or whether OH is reacting with a small amount of liquid remaining in the film after NH4HSO4 has crystallized out is unknown at present. Conclusions and Atmospheric Implications As mentioned in the Introduction, it has been shown previously by Mozurkewich et al.10 and by Hanson et al.8 that the mass accommodation coefficients of HOx radicals are quite large in relatively dilute aqueous solutions. Similarly, it is apparent from this work that the mass accommodation coefficients of OH and HO2 are comparably large in concentrated acid solutions, characteristic of acidic sulfate aerosols in the stratosphere and upper troposphere. While it is has been modeled2-5 for tropospheric cloud droplets many microns in size that gasphase diffusion will control the overall rate of mass transfer from the gas-phase to the particles for mass accommodation coefficients of this size, for the case of submicron sulfate aerosols, the overall rates of heterogeneous chemistry are expected to be more sensitive to the specific value of the mass accommodation coefficient. From a reactive standpoint, Hanson et al.36 have identified reactions with NO3- and Cl- as likely routes for loss of the SO4- radical formed from the OH + HSO4- f H2O + SO4reaction that occurs when a gas-phase OH radical is taken up by a stratospheric sulfate aerosol. In the troposphere, given that most reactions of SO4- with organic species proceed at close to the diffusion-limited rate in dilute aqueous solutions,19 it is also possible that the SO4- radical will react with organic species either dissolved in or on the surface of the aerosol. This reaction would then provide a source of organic free-radicals to solution. For the case of HO2 loss on sulfuric acid solutions, studies where the composition of the condensed phase is monitored may help to determine whether HO2 is in fact directly reacting with HSO4- to form SO5- or whether HO2 is selfreacting to form H2O2. In contrast to aqueous solutions, the net uptake coefficients measured for HOx radicals on conditioned water-ice, (NH4)2SO4, and NH4HSO4 surfaces are relatively low (γ < 0.05). Nevertheless, the studies involving reactive species adsorbed to ice and to dry (NH4)2SO4 surfaces have demonstrated the importance of defining the surface composition in order to accurately assess the rate at which these radicals will be lost on atmospheric particulate surfaces. As an example, the enhancement in the OH loss rate due to adsorbed 1-hexanol on dry (NH4)2SO4 indicates the importance which adsorbed organic layers may have on the gas-to-condensed phase mass transfer rates of a number of key tropospheric species such as OH, HO3, and N2O5. While a species such as OH, which readily abstracts H-atoms from organic molecules, would be more easily lost on an organic-rich surface, a molecule such as N2O5 would be hydrolyzed less easily if the water content of the surface is low.

Cooper and Abbatt Similarly, radical heterogeneous reactions on ice surfaces present in cirrus or polar stratospheric clouds could be affected by adsorbed species such as HNO3. Acknowledgment. Acknowledgement is made to the donors of the Petroleum Research Fund, administered by the American Chemical Society, and to NASA (Atmospheric Effects of Aviation/Subsonic Assessment Program) for support of this research. We also thank an anonymous reviewer and M. Mozurkewich for comments on some of this work. References and Notes (1) Mozurkewich, M. J. Geophys. Res. 1995, 100, 14199. (2) Lelieveld, J.; Crutzen, P. J. J. Atmos. Chem. 1991, 12, 229. (3) Jacob, D. J. J. Geophys. Res. 1986, 91, 9807. (4) Schwartz, S. E. J. Geophys. Res. 1984, 89, 11589. (5) Chameides, W. L.; Davis, D. D. J. Geophys. Res. 1982, 87, 4863. (6) Heikes, B. G.; Thompson, A. M. J. Geophys. Res. 1983, 88. 10883. (7) Baldwin, A. C.; Golden, D. M. J. Geophys. Res. 1980, 85, 2888. Baldwin, A. C. In Heterogeneous Atmospheric Chemistry; Schryer, D. R., Ed.; AGU Press: Washington, DC, 1982, pp 99-102. (8) Hanson, D. R.; Burkholder, J. B.; Howard, C. J.; Ravishankara, A. R. J. Phys. Chem. 1992, 96, 4979. (9) Gershenzon, Yu. M.; Ivanov, A. V.; Kucheryavyi, S. I.; Rozenshtein, V. B. Kinet. Katal. 1986, 27, 923. (10) Mozurkewich, M.; McMurry, P. H.; Gupta, A.; Calvert, J. G. J. Geophys. Res. 1987, 92, 4163. (11) Jech, D. D.; Easley, P. G.; Krieger, B. B. In Heterogeneous Atmospheric Chemistry; Schryer, D. R., Ed.; AGU Press: Washington, DC, 1982; pp 107-121. (12) Brown, R. L. J. Res. Natl. Bur. Stand. (U.S.) 1978m 83, 1. (13) Marrero, T. R.; Mason, E. A. J. Phys. Chem. Ref. Data 1972, 1, 3. (14) Plumb, I. C.; Ryan, K. R.; Barton, N. G. Int. J. Chem. Kin. 1983, 15, 1081. (15) Hanson, D. R.; Ravishankara, A. R. J. Phys. Chem. 1992, 96, 9441. (16) Zhang, R.; Wooldridge, P. J.; Abbatt, J. P. D.; Molina, M. J. J. Phys. Chem. 1993, 97, 7351. (17) Farhatziz; Ross, A. B. Selected specific rates of reactions of transients from water in aqueous solutions. Hydroxyl radical and perhydroxyl radical and their radical ions. NBS National Standards Reference Data Series NSRDS-NBS 59; National Bureau of Standards: Washington, DC, 1977. (18) Tang, Y.; Thorn, R. P.; Mauldin, R. L., III; Wine, P. H. J. Photochem. Photobiol. A: Chem. 1988, 44, 243. (19) Neta, P.; Huie, R. E. J. Phys. Chem. Ref. Data 1988, 17, 1027. (20) Cox, R. A. J. Am. Chem. Soc. 1974, 96, 1059. (21) Stanbury, D. M. Reduction Potentials involving inorganic free radicals in aqueous solution. In AdVances in Inorganic Chemistry; Sykes, A. G., Ed.; Academic Press: San Diego, 1989; Vol. 33. (22) Wood, B.; Wise, H. J. Phys. Chem. 1962, 66, 1049. (23) Mannella, G.; Harteck, P. J. Chem. Phys. 1961, 34, 2177. (24) Hacker, D. S.; Marshall, S. A.; Steinberg, M. J. Chem. Phys. 1961, 35, 1788. (25) Smith, W. V. J. Chem. Phys. 1943, 11, 110. (26) Steinfeld, J. I.; Francisco, J. S.; Hase, W. L. Chemical Kinetics and Dynamics; Prentice Hall: Englewood Hills, 1989. (27) DeMore, W. B.; Sander, S. P.; Golden, D. M.; Hampson, R. F.; Kurylo, M. J.; Howard, C. J.; Ravishankara, A. R.; Kolb, C. E.; Molina, M. J. Chemical kinetics and photochemical data for use in stratospheric modeling. JPL Publication 92-20; NASA Jet Propulsion Laboratory: Pasadena, 1992. (28) Abbatt, J. P. D.; Beyer, K. D.; Fucaloro, A. F.; McMahon, J. R.; Wooldridge, P. J.; Zhang, R.; Molina, M. J. J. Geophys. Res. 1992, 97, 15819. (29) Hanson, D. R. Geophys. Res. Lett. 1992, 19, 2063. (30) Diehl, K.; Mitra, S. K.; Pruppacher, H. R. Atmos. EnViron. 1995, 29, 975. (31) Hanson, D.; Mauersberger, K. J. Phys. Chem. 1988, 92, 6167. (32) Mozurkewich, M.; Benson, S. W. J. Phys. Chem. 1984, 88, 6441. (33) Abbatt, J. P. D.; Molina, M. J. Geophys. Res. Lett. 1992, 19, 461. (34) Keyser, L. T.; Leu, M-T.; Moore, S. B. J. Phys. Chem. 1993, 97, 2800. (35) Gill, P. S.; Graedel, T. E.; Weschler, C. J. ReV. Geophys. Space Sci. 1983, 21, 903. (36) Hanson, D. R.; Ravishankara, A. R.; Solomon, S. J. Geophys. Res. 1994, 99, 3615.

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