J. Phys. Chem. 1996, 100, 8015-8022
8015
Uptake of Gas-Phase Formaldehyde by Aqueous Acid Surfaces J. T. Jayne,* D. R. Worsnop, and C. E. Kolb Center for Chemical and EnVironmental Physics, Aerodyne Research, Inc., Billerica, Massachusetts 01821-3976
E. Swartz and P. Davidovits Department of Chemistry, Boston College, Chestnut Hill, Massachusetts 02167-3860 ReceiVed: October 30, 1995; In Final Form: February 29, 1996X
The uptake of gas-phase formaldehyde, CH2O, by aqueous sulfuric and nitric acid droplets has been measured as a function of temperature, acid concentration, and gas-droplet contact time. The results are interpreted in terms of the mass accommodation coefficient for formaldehyde on aqueous surfaces, the Henry’s law constant for formaldehyde in aqueous acid solutions, and the liquid-phase formation kinetics of methylenediol, CH2(OH)2, and protonated formaldehyde, CH3O+. Time-dependent uptake studies under mildly acid solutions where CH2(OH)2 and CH3O+ formation is slow reveal the existence of a chemisorbed surface complex for CH2O at the gas-liquid interface. Uptake studies on nitric acid and mixed sulfuric acid/nitric acid solutions show slightly enhanced uptake relative to sulfuric acid only solutions. This observation has been attributed to variation of formaldehyde solubility (expressed as Setchenow coefficients) and CH3O+ equilibrium constant in nitric and sulfuric acid solutions. The implications of these measurements for the aqueous acid chemistry of formaldehyde and its role in atmospheric chemistry are discussed.
Introduction Formaldehyde plays an important role in both industrial chemistry, where it is a key reactant in many industrial processes and a significant component of many manmade materials,1 and environmental chemistry, where it is a pollutant emitted by combustion exhaust sources2 as well as many building materials and furnishings3,4 and is also a key reactive intermediate in the photochemical oxidation of airborne hydrocarbons.5 Formaldehyde has also served as a molecular testbed for the development of stimulated emission pumping, an important tool for the study of molecular dynamics and highly excited state molecular spectroscopy, pioneered by Kinsey and co-workers.6-8 It is a pleasure to dedicate this study on the kinetic behavior of formaldehyde at the gas-liquid interface and in aqueous acid solutions to Professor James L. Kinsey in recognition of his seminal contributions to the fields of molecular dynamics and chemical kinetics and to his secure status as an inspiring mentor and a valued colleague. Current interest in the interaction of CH2O with aqueous and aqueous acid droplets arises from formaldehyde’s key role in atmospheric photochemistry. In the gas-phase CH2O serves as one of the most potent precursors of ozone pollution in urban and industrialized regions.9-11 Gas-liquid interactions with cloud droplets and aerosols remove atmospheric CH2O from the gas phase, resulting in significant levels of formaldehyde in rain, snow, and fog.12-15 In fog and cloud droplets CH2O reacts with SO2 to form hydroxymethanesulfonate, enhancing the level of dissolved S(IV) but inhibiting the aqueous-phase oxidation of SO2 to SO42- by H2O2 and other oxidants.16-21 CH2O also reacts with hydroxyl radicals in cloud and fog droplets to produce formic acid (HCCOH), a contributor to acid rain and deposition.20-23 Our previously published studies of the uptake of gas-phase CH2O by aqueous droplets provided the kinetic parameters needed to predict the rate of gaseous formaldehyde scavenging by tropospheric clouds.24 X
Abstract published in AdVance ACS Abstracts, April 1, 1996.
S0022-3654(95)03196-0 CCC: $12.00
Recently, higher than expected ratios of NOx (NO + NO2) to HNO3 in the free troposphere have led to the speculation that CH2O reacts with nitrate ions in atmospheric aerosols to release volatile methyl dinitrate, hydroxymethyl nitrate, nitrous acid, or NOx back into the gas phase.25,26 If such processes are efficient in the predominantly acidic sulfate aerosols found in the free troposphere, they would have a significant impact on the regional and global production of tropospheric O3 by recycling unreactive aerosol NO3- back into reactive gaseous nitrogen oxide compounds. In addition, using Knudsen cell techniques, Tolbert and co-workers have demonstrated that CH2O is efficiently absorbed by cold, concentrated sulfuric acid solutions, suggesting that it will be absorbed into stratospheric sulfuric acid aerosols providing an additional potential sink for halogen and hydrogen oxide free radicals which catalytically destroy stratospheric ozone.27 The experimental studies presented below investigate the uptake kinetics of CH2O by cold, concentrated sulfuric acid, nitric acid, and mixed sulfuric acid/ nitric acid surfaces characteristic of the major components of upper tropospheric and stratospheric aerosols. In order to assess the relative importance of the heterogeneous chemistry of formaldehyde with sulfuric/nitric acid surfaces, one needs to know the accommodation coefficient R which governs the rate of mass transfer from the gas to liquid phase. This parameter is defined as
R)
no. of molecules absorbed by the surface no. of molecular collisions with the surface
(1)
The rate of trace gas transfer, which determines the maximum flux into the liquid, J, is given by
J ) ngjcR/4
(2)
where ng is the trace gas number density and jc is the trace gas average thermal velocity (cm s-1). In the atmosphere and in our experiments the actual rate of transfer is usually limited by process such as gas-phase diffusion and Henry’s law saturation © 1996 American Chemical Society
8016 J. Phys. Chem., Vol. 100, No. 19, 1996
Jayne et al. measured decrease in the trace gas signal (∆ng) resulting from an increase in the exposed droplet surface area corresponds to an uptake of the gas by the droplet surface. The uptake coefficient (γmeas) was obtained from the measured change in trace gas signal via eq 3:
γmeas ) 4Fg/cj∆A ln(ng/n′g)
Figure 1. Droplet train/flow reactor apparatus with tunable diode laser detection.
effects. For this reason in the following section we define a measured uptake coefficient, γmeas, which includes the above effects as well as liquid-phase reactivity when present. A detailed discussion of how these effects are accounted for in our experimental approach can be found in the work of Worsnop et al.28 Experimental Description The apparatus used in these experiments is shown schematically in Figure 1. The droplet train flow tube used in these experiments is similar to the one utilized in our previous CH2O uptake studies,24 but with modifications allowing the generation and manipulation of colder and more concentrated trains of acid droplets29. Also, in this study gas detection is performed using IR absorption rather than the mass spectrometry used in the earlier study. Here we summarize the experimental method and highlight the modifications required to extend the work to acid droplets at low temperatures. The gas uptake coefficient was measured by passing a fast-moving (1500-3000 cm/s), monodisperse (∼200 µm in diameter), and spatially collimated train of aqueous acid droplets through a 30 cm long longitudinal low pressure (4-10 Torr) flow tube which contained trace amounts of CH2O introduced through one of three loop injectors located along the flow tube and entrained in a flowing mixture of helium and water vapor. By selecting the gas inlet port (5.5, 15.5, or 29 cm) and the droplet velocity the gas-droplet interaction time could varied between 2 and 15 ms. The number density (ng) of the trace gas was measured downstream of the flow tube by infrared absorption. The surface area of the droplets is changed in a stepwise fashion by varying the driving frequency of a piezoelectric ceramic; a
(3)
where Fg is the carrier gas volume rate of flow (cm3 s-1) through the system, ∆A is the change in the total droplet surface in contact with the trace gas, and ng and n′g are the trace gas densities at the inlet and outlet of the flow tube respectively (i.e., ng ) n′g + ∆ng). The droplet stream was produced by forcing the acid solution through a 70 µm diameter platinum electron microscope aperture surrounded by a donut-shaped piezoelectric ceramic. Before flowing through the nozzle aperture the acid solution is precooled initially using an external temperature-controlled circulating bath. A chromel-alumel thermocouple in a stainless steel sheath was fixed in place just above the aperture and provided a continuous measure of the stream temperature. As the droplets traverse the flow tube they equilibrate with the ambient water vapor. In order to minimize changes in droplet temperature and concentration from evaporation, a calibrated flow of H2O in helium was introduced into the droplet generation region of the flow tube. In practice, because the droplets do evaporate somewhat, it is never possible to add precisely the correct amount of water vapor to be in equilibrium with the droplets at the initial concentration and nozzle temperature. In other recent experiments with this apparatus Robinson et al.29 have utilized direct measurements of water vapor, with the tunable diode laser system, to demonstrate that droplet surface acid concentrations in the gas/droplet interaction region are within 0.5 wt % of their initial value. As noted above, the uptake coefficients, γmeas, were obtained by changing the droplet frequency in a stepwise fashion thus varying the droplet surface area. In most of these experiments, the low frequency was exactly 8 times smaller than the high frequency resulting in a factor of 2 increase in total exposed surface area. Since the liquid flow rate remains constant as the frequency is changed, the diameter of the higher frequency droplets are smaller than the lower frequency droplets by a factor of 2. The increased surface-to-volume ratio for the higher frequency droplets in the presence of room temperature flow tube walls results in the high-frequency droplets being 1-2 K warmer than the low-frequency droplets. The droplet temperatures reported below here are the averages of the high- and low-frequency droplet temperatures. Pressure differences between the flow tube and the droplet collection chamber of the apparatus were monitored and balanced. The pressure balance was checked further by monitoring the concentration of a reference gas, in this case methane, which is added to the flow. Because methane is effectively insoluble in the droplets, changes in methane concentration resulting from small pressure imbalances associated with droplet frequency switching were subtracted from the observed changes in the CH2O concentration. Such changes in methane concentration were usually less than 0.5%. The vapor pressure of formaldehyde at room temperature is about 6 atm; however, formaldehyde gas readily polymerizes into paraformaldehyde which is a white solid and can be heated to obtain gas-phase formaldehyde. In our experiment, approximately 2 g of paraformaldehyde is loosely packed into a 1/4 × 6 in. glass column with glass wool end plugs. The column is heated to about 60 °C. One of the ends of the column is connected to the sample loop injector mounted inside the
Uptake of Gas-Phase CH2O by Aqueous Acid Surfaces
J. Phys. Chem., Vol. 100, No. 19, 1996 8017
TABLE 1: Experimental Uptake and Model Values wt % H2SO4
wt % HNO3
exp no.
temp (K)
aH+ (mol L-1)
aH2O
Dl (cm2 s-1)
slope (σ) (×100 s-1/2)
γ8 ms (σ)
10 20 40 55 55 55 55 60 70 70 70 70 80 85 8 10 20 30 36 36 36 40 40 40 40 40 0 0
0 0 0 0 0 0 0 0 0 0 0 0 0 0 12 38 24 11 18 18 18 10 15 15 20 20 20 54
16 15 14 8 9 7 6 17 13 12 11 10 27 28 26 24 22 23 1 2 3 18 19 20 4 5 25 21
267 258 263 273 261 253 241 263 273 263 253 245 300 300 263 263 262 263 234 244 274 263 263 247 254 261 265 263
1.14E+00a 3.68E+00 4.77E+01 4.85E+02 6.06E+02 7.52E+02 9.60E+02 1.59E+03 8.35E+03 1.19E+04 1.72E+04 2.29E+04 1.70E+04 3.35E+04 3.58E+00 4.32E+01 4.30E+01 4.39E+01 3.48E+02 3.02E+02 1.81E+02 1.67E+02 2.98E+02 4.00E+02 6.18E+02 2.06E+02 3.50E+00 4.29E+01
0.955 0.874 0.522 0.219 0.207 0.194 0.180 0.127 0.034 0.030 0.026 0.024 0.006 0.001 0.854 0.391 0.461 0.505 0.221 0.233 0.271 0.321 0.233 0.215 0.148 0.254 0.842 0.305
4.96E-06 2.23E-06 1.11E-06 6.08E-07 3.79E-07 2.65E-07 1.41E-07 2.74E-07 1.63E-07 1.03E-07 6.01E-08 3.63E-08 1.57E-07 8.44E-08 5.56E-06 3.38E-06 2.39E-06 1.72E-06 2.62E-07 4.60E-07 1.58E-06 9.66E-07 8.76E-07 4.34E-07 5.39E-07 7.25E-07 8.79E-06 4.29E-06
49.5 (16.5) 34.9 (7.6) 21.6 (4.6) 8.57 (5.30) 8.40 (1.68) 10.5 (3.07) 25.2 (0.20) 3.83 (0.49) 1.34 (0.75) 1.81 (0.96) 1.73 (0.67) 0.86 (0.89) 10.1 (1.91) 11.4 (2.1) 34.0 (17.2) 6.25 (0.88) 8.90 (0.34) 16.4 (3.8) 5.91 (2.24) 4.79 (3.78) 11.6 (0.6) 5.77 (1.65) 3.95 (1.63) 3.25 (0.44) 5.82 (0.21) 3.37 (0.91) 20.3 (3.05) 8.10 (1.22)
0.0027 (0.0004) 0.0046 (0.0006) 0.0064 (0.0008) 0.0062 (0.0008) 0.0109 (0.0012) 0.0078 (0.0015) 0.0067 (0.0009) 0.0130 (0.0011) 0.0244 (0.0022) 0.0244 (0.0018) 0.0274 (0.0021) 0.0254 (0.0035) 0.0130 (0.0013) 0.0114 (0.0009) 0.0040 (0.0008) 0.0126 (0.0009) 0.0075 (0.0005) 0.0060 (0.0009) 0.0110 (0.0011) 0.0108 (0.0014) 0.0089 (0.0010) 0.0101 (0.0008) 0.0144 (0.0013) 0.0162 (0.0013) 0.0090 (0.0019) 0.0166 (0.0011) 0.0050 (0.0007) 0.0121 (0.0009)
a
E(n ≡ ×10(n.
droplet flow tube. The other end of the glass column is connected to a controlled flow of helium used to entrain the vaporized formaldehyde and carry it through the loop injectors into the flow tube. Formaldehyde is detected using a diode laser absorption measurement. The gases exiting the flow tube are pumped through a multipass absorption cell (White cell) with an effective path length of about 750 cm. The absorption line for formaldehyde used in these experiments is located at a frequency of 2916.8183 cm-1 which has a line strength of 3.6 × 10-20 cm2 molecule-1 cm-1. Gas-phase densities of about 5 × 1012 cm-3 will yield a fractional absorbance of 5% which was typical for the experiments. The effect of significant variation of the formaldehyde density on the measured uptake was not investigated in these studies. Acid solutions were prepared by diluting known volumes of reagent grade sulfuric acid (95-98 wt %) and/or nitric acid (70 wt %) with known volumes of distilled water. The solution densities were further checked by weighing 25 mL aliquots. The uncertainty in the solution concentrations is ∼(1.0 wt %. Results Gas uptake measurements for formaldehyde were performed as function of gas-droplet interaction time, acid concentration, and droplet temperature. Initial experiments were conducted using sulfuric acid mixtures ranging from 10 to 70 wt %. To test for the effect of nitrate on the formaldehyde uptake mechanism experiments were also conducted on droplets of 20 and 54 wt % nitric acid as well as mixtures containing both sulfuric and nitric acid. The full range of droplet compositions and temperatures is presented in Table 1. In all experiments the uptake of formaldehyde exhibited a time-dependent behavior. As an example, we show in Figure 2a-h results which plot γmeas vs gas droplet interaction time for H2SO4 at 263 K. Two trends are evident: (1) the uptake decreases with increasing interaction time and (2) the uptake increases with increasing acid strength up to 70 wt %. These trends are qualitatively
consistent with a mechanism involving limited solubility and acid-catalyzed reactivity. The dotted lines in the figures are plots of the formaldehyde uptake model presented in Jayne et al.24 which is built on the following equilibria: H1
K2
CH2O(g) T CH2O(aq) + H2O T CH2(OH)2
(4)
Here, H1 is the physical Henry’s law constant for formaldehyde and K2 is the equilibrium constant for the hydration of formaldehyde to its more soluble gem-diol form:
K2 )
[CH2(OH)2(aq)] [CH2O(aq)][H2O]
(5)
The formation of the diol greatly increases effective formaldehyde solubility in water since K2 is ∼2000 at 298 K.30,31 The increased solubility can be expressed in terms of an effective Henry’s law constant, H*,
H* )
[CH2O(aq)] + [CH2(OH)2(aq)] [CH2O(g)]
) H1(1 + K2aH2O) (6)
which has value of 2.97 × 103 M atm-1 at 298 K.31 Here, aH2O is the activity coefficient for water. The formation rate of the gem-diol is both acid and base catalyzed. The hydrolysis rate constant has been measured by Schecker and Schulz;32 under acid conditions their hydrolysis rate constant is in reasonable agreement with the gem-diol dehydration rate constant for the reverse reaction published by Bell and Evans33 and the Bell equilibrium constant.30 As shown in Jayne et al.,24 the rate-limiting step in formaldehyde uptake is governed by its rate of hydration. However, it is immediately evident from Figure 2a-h that the aqueous kinetics (displayed as dotted lines from modeling discussed below) used in Jayne et al.24 fail to account for the observed uptake across the entire range of acidities. Similar results were observed in the work of Tolbert
8018 J. Phys. Chem., Vol. 100, No. 19, 1996
Jayne et al.
Figure 3. Resistor model diagram for formaldehyde uptake.
wt % acid). In this range, according to our previous model, the uptake should be limited only by the rate of acid-catalyzed diol formation and should be time independent on the millisecond time scale of our experiment. Observation of uptake decreasing at longer gas-droplet interaction times, with γmeas inversely proportional to time, is indicative of chemisorbed surface species. Such time-dependent behavior has been observed previously in our uptake studies of acetaldehyde24 and sulfur dioxide.35 The existence of a surface chemisorbed SO2 complex has recently been spectroscopically confirmed by Donaldson et al.36 CH2O Uptake Model
Figure 2. Uptake coefficients (γmeas) measured for CH2O(g) exposed to H2SO4 solutions as a function of gas-droplet contact time. Solid line is plot of new formaldehyde uptake model which includes CH3O+ solubility γ(CH3O+), gem-diol hydration and solubility (γ(CH2(OH)2)) and surface complex (γsurf). The dotted line shows γ(CH2(OH)2), and the dashed line is the sum of γ(CH2(OH)2) + γsurf. The contribution of γ(CH3O+), indicated by the difference of the solid and dashed lines, increases with H2SO4 concentration.
et al.27 where enhanced uptake was measured on surfaces with g70 wt % H2SO4. For all the acid uptake results presented in Table 1 the measured uptake is larger than predicted by gemdiol formation kinetics alone. In fact, in our original work24 the uptake in the pH range of 0-2 was also larger than the model prediction. These results imply that at least one additional mechanism is responsible for uptake in this acidity range. We propose two additional uptake mechanisms: first, following the work of McTigue and Sime34 and Tolbert et al.,27 the equilibrium formation of a protonated form of formaldehyde, CH3O+ K3
CH2O + H+ T CH3O+
(7)
accounts for the observed uptake at high acidities and, second, in the mildly acidic range (20 wt %. As with γ(CH2O), no γrxn term is included in the γ(CH3O+) formulation, since there is no indication of any observed rate limitation due to the protonation reaction. In other words
CH2O + H+ f CH3O+
(17)
Figure 5. Uptake coefficient, γ8 ms, plotted for constant acidity as the fraction of H2SO4 and HNO3 is varied. Solid lines are model predictions for the experimental conditions at 263 K, showing increase in H*CH3O+ in HNO3 solution.
is fast. This implies that the rate-limiting step for acid-catalyzed CH2(OH)2 formation is the reaction
CH3O+ + H2O f CH2(OH)2 + H+
(18)
This is consistent with the formulation of the aH+ term in kH2O,H+,OH- for γ(CH2(OH)2) in Table 2; i.e., kH+ (as reported from previous experiments) is the product of K3 (assuming equilibration of [CH3O+] via eq 17) and the reaction rate of eq 18. In other words, as discussed above for γ(CH2O), it is appropriate to include γ(CH2(OH)2) and γ(CH3O+) as parallel uptake channels as shown in Figure 3. It should be noted that the value of R ∼ 0.04, which is only weakly constrained by the nonlinear fit to the data in concentrated acid solution, might reflect a surface-specific reaction for CH2O protonation (eq 17), which is indistinguishable from R (see eq 15) for the maximum γmeas into 70 wt % H2SO4 solution. We have seen such surface reactivity in our previous work which measured uptake of Br2 and Cl2 onto aqueous droplets containing Br- and I- ions.40 As is clear from Figure 2c-h, γmeas shows a time dependence that is consistent with t-1/2 predicted by eq 12. As shown in Table 2, H*CH3O+ has been modeled by expressing K3 for CH2O protonation as a temperature-dependent equilibrium constant. As shown in Figures 2 and 4b, the model fits the observed time dependence for uptake in H2SO4 solutions well, including predicting the turnover in observed uptake for H2SO4 concentration >70 wt % caused by rapidly decreasing aH2O. In Figure 5 we show time-dependent uptake results at 263 K for mixed H2SO4/HNO3 acid solutions prepared at constant acidity (experiments 14, 21, 22, 23, and 24). The relative uptake of CH2O in H2SO4 and HNO3 solution is accounted for by differing CH2O solubility as expressed by the effective Setchenow coefficients ΓH2SO4 and ΓHNO3 in Table 2. The use of these separate ΓH2SO4 and ΓHNO3 terms implicitly includes any variation of the activity coefficient of CH3O+ species in H2SO4 and HNO3 solution. However, lack of knowledge of CH3O+ activity precluded its explicit parametrization in the model. In this context, it should be noted that the overall model parametrization depends on the product of the terms ΓCH2O, Dl and K3. Changes in the values of any of these terms will inversely affect the magnitude of the others. Thus,
Uptake of Gas-Phase CH2O by Aqueous Acid Surfaces
J. Phys. Chem., Vol. 100, No. 19, 1996 8021
Figure 6. Temperature dependence of the uptake coefficient, γ8 ms, for 55 and 70 wt % H2SO4 solutions. Lines are model predictions for the experimental conditions. Overall (weak) temperature dependence reflects combination of HCH2O, Dl and K3.
the relative CH2O solubility (or CH3O+ activity) in H2SO4 and HNO3 solution implied by the fitted ΓH2SO4 and ΓHNO3 values could as easily be applied to the K3 or Dl parametrizations. This formulation uncertainty also applies to the temperature dependence included in the K3 term (see Table 2), which likewise might apply to ΓCH2O or Dl. As can be seen in Table 1, the observed temperature dependence of γmeas is small. Figure 6 displays the temperature dependence of γ8 ms for 55 and 70 wt %, which is representative of the T-dependent results in Table 1. The model predictions (solid and dashed lines, Figure 6) include the weak T dependence included in K3 (Table 2). The overall lack of any significant temperature dependence is attributed to opposing dependencies of H°CH2O and Dl (i.e., viscosity) which largely cancel one another. These uncertainties are most apparent in the 80-85 wt % results as plotted in Figure 4. There γmeas (both γ8 ms and the slope b) critically depend on steep variations in HCH2O, Dl, and aH+ with H2SO4 concentration. Any errors in the formulation (e.g., nonlinearities in ΓCH2O, the functional dependence of C on wt %, or the parametrization of aH+37) would significantly affect the predicted turnover in uptake for >70 wt % H2SO4 solution, where all those parametrizations are at the extreme of their validity. This includes temperature dependencies, since the 80-85 wt % H2SO4 experiments could only be performed at room temperature because of viscosity limitations in producing droplets. We can use our model results to predict the thermodynamic fractionation of the three formaldehyde forms, CH2O(aq), CH2(OH)2, and CH3O+, for a given temperature and composition. In Figure 7 we plot the model results for the Henry’s law solubility constant for each of the three species as a function of aH+ activity of H2SO4 for three atmospherically relevant temperatures: 298, 263, and 230 K. Such a plot is helpful in illustrating which species controls the overall uptake. Discussion We have demonstrated that two parameters, K3 regulating the formation of CH3O+ and γsurf describing the formation of a chemisorbed formaldehyde complex, describe CH2O uptake over a wide range of acid composition. The formation of a chemisorbed surface complex is not unreasonable, since our previous studies of aldehyde uptake by aqueous surfaces clearly demonstrated the existence of such a state for acetaldehyde;24 however, as discussed in ref 24 the acetalaldehyde surface
Figure 7. Henry’s law solubilities calculated by model for CH2O(aq) (solid line), CH2(OH)2 (dotted line), and CH3O+ (dashed line) plotted as a function of aH+ activity for H2SO4 solutions at three temperatures, (a) 298 K, (b) 263 K, and (c) 230 K. Numbers on the top axis of each graph are the corresponding wt % sulfuric acid content.
complex is observable at much higher pH levels because that species has the ability to isomerize to the enolate form in interactions with OH-. The action of CH2O as a base, forming CH3O+, is also not surprising,27,34 but the value of K3 indicated by our data is surprising since it requires CH2O to act as a much stronger base in concentrated acid solutions than is indicated by the previously measured value of its basic pKa.34 Our work does confirm earlier conclusions by Tolbert et al.27 that CH2O uptake on sulfuric acid aerosols will be efficient, suggesting that liquid-phase CH2O reactions may play an important role in upper tropospheric and stratospheric chemistry. Those experiments, which used stirred H2SO4 solutions in a Knudsen flow reactor,27 observed somewhat higher CH2O
8022 J. Phys. Chem., Vol. 100, No. 19, 1996 uptake rates (especially for H2SO4 >60 wt %) that exhibited significant negative temperature dependence. Those results are qualitatively consistent with H*CH3O+ solubility in the current model with additional uptake enhanced by stirring of the liquid. The current work also shows that mixed sulfuric acid/nitric acid aerosols will also efficiently absorb formaldehyde. The uptake model parameters displayed in Table 2 can be used by atmospheric modelers to predict the gas/aerosol CH2O partitioning for realistic aerosol compositions. Our CH2O uptake model attributes the differences between sulfuric and nitric acid surfaces in terms of differing Setchenow coefficients. However, some of the observed differences in γmeas may be due to reaction between CH2O and NO3-. On the time scale of the experiments reported here (e15 ms) it is not clear how much, if any, of the uptake of CH2O in droplets containing HNO3 is due to its reaction. Low-temperature (
