15218
J. Phys. Chem. 1996, 100, 15218-15228
Knudsen Cell Studies of the Uptake of Gaseous HNO3 and Other Oxides of Nitrogen on Solid NaCl: The Role of Surface-Adsorbed Water Peter Beichert† and Barbara J. Finlayson-Pitts* Department of Chemistry, UniVersity of California, IrVine, IrVine, California 92697-2025 ReceiVed: March 27, 1996; In Final Form: June 18, 1996X
A newly designed and constructed Knudsen cell has been tested by measuring the reaction probability for gaseous N2O5 on 65% H2SO4/H2O at 210-230 K to be 0.075 ( 0.047 (2σ), which is in excellent agreement with the literature value. This cell has been applied to the study of gaseous HNO3 reactions with NaCl small crystals and ground powders at 298 K. A rapid initial uptake of HNO3 and production of gaseous HCl are observed when the crystals and powders are pumped but not heated prior to reaction. After this rapid initial reaction, a constant uptake of HNO3 and formation of HCl is observed from which a reaction probability of (1.4 ( 0.6) × 10-2 (2 σ) is calculated. When possible systematic errors (including uncertainties in the effective surface area available for reaction) are taken into account, the overall uncertainty is estimated to be about a factor of 2. The measured reaction probability is independent of the size or preparation of the salt crystals as well as the number of layers of salt in the sample holder. This reaction probability is in excellent agreement with results from the previous work of Rossi and co-workers21,23 and Leu and co-workers15 using powders but significantly larger than that measured by Laux et al.12 using single crystals and an ultrahigh vacuum system. Prior heating of the salts while pumping decreased the extent of the initial rapid reaction but did not affect the subsequent constant reaction uptake probability. Experiments on the reaction of HNO3 with NaCl crystals that had been previously exposed to D2O to replace any surface-adsorbed H2O and on the reactions of DNO3 with NaCl show that under all experimental conditions studied here, some water remains on the surface and plays a key role in the uptake of HNO3. We propose a new model for the reaction of HNO3 with NaCl powders in which HNO3 is taken up into strongly adsorbed water (SAW) on the salt. This SAW, for which there is prior evidence in the literature,34 appears likely to be held at defect sites on the powders. Acidification of this SAW leads to degassing of HCl due to dissolution of NaCl into the SAW from the underlying salt. As gaseous HNO3 continues to be taken up, HCl degasses and nitrate precipitates out as NaNO3 . This model represents a fundamental change in the description of the heterogeneous reactions of salt powders. The lower reaction probability for single crystals observed by Laux et al.12 is consistent with the lack of surface-adsorbed water on relatively defect-free single crystals. No uptake of the gases NO2, NO, HCl, ClNO, ClNO2, or H2O was observed on the finely ground NaCl powder from which an upper limit to the reaction probabilities for these gases with NaCl of ∼10-5 was derived. The atmospheric implications of this model are discussed.
Introduction Knowledge of the kinetics and mechanisms of heterogeneous reactions is now accepted to be critical to understanding tropospheric and stratospheric chemistry. For example, the uptake and subsquent reaction of gases such as HCl and ClONO2 on ice crystals that make up polar stratospheric clouds are known to be key components of the development of the Antarctic “ozone hole”.1,2 In the troposphere, there is increasing evidence that reactions of gases with airborne sea salt particles form photochemically active halogen compounds such as Cl2, which generate highly reactive chlorine atoms.3-6 There are a variety of techniques for studying the heterogeneous reactions of gases with solids.2 Golden, Spokes, and Benson7 pioneered the applications of Knudsen cells to such processes. Other approaches include transmission IR,8,9 diffuse reflectance infrared Fourier transform spectrometry (DRIFTS),10,11 X-ray photoelectron spectroscopy (XPS),12,13 flow systems,14-16 aerosol chambers,17,18 and transmission electron microscopy† Current address: Alfred Wegener Institute for Polar and Marine Research, Air Chemistry Division, Am Handelshafen 12, 27515 Bremerhaven, Germany. * To whom correspondence should be addressed. X Abstract published in AdVance ACS Abstracts, August 1, 1996.
S0022-3654(96)00925-2 CCC: $12.00
energy dispersive spectroscopy (TEM-EDS).19,20 However, results from different laboratories using different techniques to study what is ostensibly the same system are often not in good agreement. For example, the reaction probability for gaseous HNO3 on solid NaCl at room temperature has been reported to be (4 ( 2) × 10-4 using XPS and single crystals,12 (2.0 ( 1.0) × 10-2 using a Knudsen cell and finely ground powders,21 and (1.3 ( 0.4) × 10-2 using a fast flow system and commercially available small crystals.15 There are a number of possible reasons for such discrepancies in gas-solid kinetic measurements. First, how the nature of the surface on a molecular level controls the reactivity, e.g., the presence and number of surface defects, has not been explored in a systematic way. Second, the potential role played by small amounts of surface-adsorbed water is not known. Third, there is some controversy concerning the effective surface area available for reaction in experimental systems using powders where multiple layers of loosely packed solids are used. For example, in fast flow tube studies of the uptake of gases such as ClONO2, HNO3, and N2O5 on NaCl, Leu and coworkers14,15 assume that the spaces between the particles act like pores, and hence, some of the salt area forming the internal surface of these pores should also be taken into account in © 1996 American Chemical Society
Uptake of HNO3 and Nitrogen Oxides
J. Phys. Chem., Vol. 100, No. 37, 1996 15219
calculating reaction probabilities. The measured reaction probability for HNO3 on NaCl powders without correction for this increased available surface in the interstitial pores between the crystals is γ ) (6.0 ( 1.0) × 10-2, whereas the value corrected for the internal pore surface using a model developed previously by that group22 is (1.3 ( 0.4) × 10-2. On the other hand, Rossi and co-workers in their initial Knudsen cell studies23 of this reaction reported a value of γ ) (2.8 ( 0.3) × 10-2 calculated using only the top geometric surface area of the sample holder, i.e., uncorrected for the internal pore surface area. Leu et al.15 estimate that this would be lowered to ∼0.003 if corrections were made for the “pore diffusion”. In more recent studies, Rossi and co-workers21 have reexamined this reaction using NaCl prepared in a variety of ways and concluded that the measured uptake coefficient is (2.0 ( 1.0) × 10-2 and depends only on the geometric surface area. In short, the flow tube results corrected for diffusion into the interstitial spaces between the crystals appear to agree with those from Knudsen cell studies uncorrected for such pore diffusion. Surprisingly, Rossi and co-workers21 observe that unlike HNO3, the uptake of N2O5 on NaCl does increase with an increased number of layers of NaCl, i.e., with increased pore surface area. They propose that HNO3, being a “sticky” molecule, is rapidly physically adsorbed on the salt prior to reaction, with subsequent diffusion into the pores being slow, but that N2O5 does not show the same rapid initial adsorption. We report here the results of studies of the reaction of HNO3 with NaCl,
HNO3(g) + NaCl(s) f HCl(g) + NaNO3(s)
(1)
using a newly constructed Knudsen cell. To test this new design, the reaction probability for N2O5 on H2SO4/H2O mixtures was first measured at 210-230 K for comparison to the numerous studies of this reaction reported in the literature, which are in relatively good agreement.24 The Knudsen cell was then applied to the HNO3/NaCl system using different sources of NaCl: commercially available large powders sieved to obtain a narrow size distribution and finely ground powders made by grinding single-crystal cuttings. Some studies were also carried out using (100) and (111) single crystals. We show that the most important determinant of the reaction probability appears to be small amounts of surface adsorbed water. Studies of the interactions of NO2, NO, HCl, ClNO, ClNO2, and H2O on the finely ground powder were also carried out. The implications for the application of such measurements to atmospheric reactions are discussed. Experimental Section Knudsen cell studies are based on the competition between the escape of a gas from a low-pressure cell by “collision” with a hole leading to a detector and its removal by collisions with a surface, which remove the gas from the cell by adsorption and/or reaction. As discussed in detail elsewhere,7,23,25 the reaction probability, γ, is given by
γ)
()
Ah (N0 - Nr) As Nr
(2)
where Ah is the area of the aperture hole to the mass spectrometric (MS) detection system, As is the area of the reactive surface, and N0 and Nr are the MS signals in the absence and presence of the reactive surface, respectively. In systems where the uptake is due to reaction and there is no reevaporation of the gas from the surface, the reaction probability, γ, is the
Figure 1. Schematic diagram of the Knudsen cell apparatus.
same as the mass accommodation coefficient, R, where the latter is defined as the fraction of collisions of the gas with the surface that leads to uptake of the gas.25 The Knudsen cell used in this study is shown schematically in Figure 1. It differs from those used in other studies in that it has a single chamber with the reactive surface covered by a lid prior to reaction. This has the advantage over a dual chamber design in that lifting the lid to expose the sample to the gas does not significantly (99.5%) and retaining the fraction that passed through a no. 20 sieve but not a no. 40 sieve, giving particles in the size range 425-850 µm.27 However, the particles used were barely retained by the no. 40 sieve. We therefore estimate that a typical particle size is closer to ∼500 µm and shall refer to them as such in the following. Single crystals with (100) and (111) surface orientations were purchased from Solon Technologies. Before each experiment, the salt samples were introduced into the Knudsen cell and pumped on for several hours or, for some experiments, heated under vacuum at 353 K. The temperature was limited by melting of the halocarbon wax coating on the cell at higher temperatures. All uptake studies on NaCl were carried out at room temperature. For the studies of the uptake of HNO3 on NaNO3, finely ground NaNO3 was prepared by grinding small crystals (Aldrich, >99.0%) in the Wig-L-Bug for 5 min. Dry nitric acid was prepared by distillation of a mixture (1:2 ratio v/v) of concentrated nitric acid (Fisher, 69.0-71.0%) and concentrated sulfuric acid (Aldrich, 98%). N2O5 was produced by the reaction of NO (Matheson, 99.3%) with an excess of O3 produced using O2 (Air Liquide 99.999%) and a commercial ozonizer (Polymetrics Inc., Model T-816). The N2O5 was condensed in a dry ice-acetone trap. NO2 was prepared by the oxidation of NO with O2 (Air Liquide 99.999%) followed by trapping at 195 K. DNO3 was prepared by reacting D2SO4 (98% with D2O, >99.5 D atom %, Sigma) with NaNO3 (Aldrich, >99.0%) and trapping at 196 K. ClNO was synthesized as described in previous work28 by the reaction of NO with Cl2 (Mattheson, 99%). The crude ClNO was trapped in a dry ice-acetone bath, and the vapor was pumped off repeatedly until the vapor pressure reached an equilibrium of ∼5 Torr. ClNO2 was prepared as previously described29 by the reaction of HCl gas with a mixture of fuming nitric and sulfuric acids. Results Because many of the gases studied here do not have strong parent peaks in their electron impact mass spectra, for most experiments, one or more of the major fragment peaks in the mass spectra shown in Figure 2 were followed. To test our Knudsen cell, we first measured the reaction probability for N2O5 on sulfuric acid-water mixtures, since this reaction probability has been measured in the past using a number of different techniques.24 The results of past studies in other laboratories are in good agreement with values between 0.05 and 0.16 for 39-96 wt % H2SO4 at temperatures from 213 to 293 K. Since N2O5 does not show a parent peak in the mass spectrum (Figure 2), we followed the fragment peaks at
Beichert and Finlayson-Pitts
Figure 2. Typical mass spectra of gases of interest for this study.
m/e ) 46 and m/e ) 30. Figure 3 shows typical results for the uptake of N2O5 on 65% H2SO4 at 210 K. Although both the m/e ) 46 and the m/e ) 30 peaks decrease reproducibly on opening and closing the lid, showing no signs of surface saturation, the signal at m/e ) 46 drops to a slightly larger extent, indicating some contribution to the signals from species other than N2O5. As discussed in detail by Manion et al.,30 this is a common problem using EI-MS and increases the uncertainty associated with such measurements. However, the increased uncertainty is acceptable in our case, since these experiments were used simply as a probe for potential major systematic errors in our new apparatus. “Blank” runs were carried out to test for the uptake of N2O5 in the absence of sulfuric acid; no uptake was observed. Table 1 summarizes our measurements of the reaction probability for the N2O5/H2SO4/H2O experiments using the m/e ) 46 and m/e ) 30 peaks. The reaction probabilities calculated from the raw data are 0.073 ( 0.046 using the m/e ) 46 peak and 0.050 ( 0.040 (2σ) using m/e ) 30. Corrected reaction probabilities calculated as described in the Appendix, assuming either an NO2 interference or an HNO3 interference, are also given in Table 1. In a typical experiment, the initial ratio of m/e ) 46 to m/e ) 30 for the reactant N2O5 was ∼2:1. If NO2 is assumed to be present as a nonreactive, small impurity in the N2O5 and is not taken up by the liquid (for which there is experimental support),24 the calculated reaction probabilities using the two masses are consistent within experimental error, giving γ ) 0.075 ( 0.047 (2σ). If it is assumed that HNO3 is formed upon uptake of N2O5 and subsequently degasses, the corrected reaction probabilities are inconsistent and, indeed,
Uptake of HNO3 and Nitrogen Oxides
J. Phys. Chem., Vol. 100, No. 37, 1996 15221
Figure 3. Typical uptake of N2O5 measured using (a) m/e ) 46 and (b) m/e ) 30 for N2O5 on 65% H2SO4 at 210 K. [N2O5]0 ) 1 mTorr, total surface area is 18.9 cm2, and aperture diameter is 6 mm.
often negative or infinite. This is in contrast to the measurements of Manion et al.30 who report interference from degassing of HNO3. In their studies, stirring the acid solution to prevent the surface layer from becoming saturated increased the measured reaction probability from ∼0.02-0.03 to 0.06. In short, our results are consistent with the presence of small amounts of NO2 in the N2O5. The corrected value of the reaction probability is approximately 20% larger than the uncorrected value. Both, however, are within the range of values observed in earlier studies,24 indicating that there are no large unrecognized systematic errors in the design of our Knudsen cell. Figure 4 shows some typical data for the uptake of HNO3 by NaCl crystals prepared in two different ways and that had been pumped on but not heated. Figure 4a shows the uptake of HNO3 measured following the large m/e ) 46 peak on commercially available small NaCl crystals that had been sieved to provide ∼500 µm crystals. Figure 4b shows the uptake on finely ground
Figure 4. Typical data for the uptake of HNO3 (m/e ) 46) and the production of HCl (m/e ) 36) on (a) 500 µm and (b) 4 µm NaCl crystals at room temperature. Aperture diameter is 9 mm, sample holder surface area is 18.9 cm2, and [HNO3]0 = 0.1 mTorr. Salt samples were pumped for several hours but not heated.
powders with a typical 4 µm diameter. In both cases, there is a rapid initial uptake of the nitric acid followed by a slower, but constant, uptake. Corresponding to the uptake of HNO3 is the simultaneous (on our time scale) production of HCl (m/e ) 36). The large initial uptake is especially pronounced with the smallest (∼4 µm) ground NaCl crystals, where a steady state is reached after ∼30 min reaction time. This large initial uptake appears to be related to water adsorbed on the salt. Thus, if the salt samples were first heated for several hours under vacuum (Figure 5), the magnitude of the initial uptake was significantly smaller for both crystal sizes and preparations. In the case of the 500 µm crystals, the HNO3 uptake immediately assumed a constant value corresponding to that reached without heating (Figure 4a). However, as seen in Figure 5b, even after heating for 24 h under vacuum, the
TABLE 1: Summary of Studies of the Reaction Probability for N2O5 on 65% H2SO4/H2O at 210-230 K reaction probability aperture diameter [mm] 4 4 6 6 9 9
a
[N2O5] (1012 molecule cm-3)
raw data m/e ) 46 γ46uncorr
raw data m/e ) 30 γ30uncorr
corrected for NO2 m/e ) 46 γ46corr
corrected for NO2 m/e ) 30 γ30corr
90 9 7 50 3 60
0.041 0.061 0.103 0.065 0.073 0.097
0.027 0.032 0.069 0.039 0.058 0.074
0.044 0.071 0.112 0.071 0.076 0.102
0.039 0.059 0.102 0.057 0.070 0.093
mean (( 2 σ)
0.073 (0.046
0.050 (0.040
0.079 (0.048
0.070 (0.047
Calculated value is ∞. b Calculated value is negative.
corrected for HNO3 m/e ) 46 γ46corr
corrected for HNO3 m/e ) 30 γ30corr
a -0.095 -2.19 a 0.169 0.265
a b b a 0.169 0.260
15222 J. Phys. Chem., Vol. 100, No. 37, 1996
Beichert and Finlayson-Pitts TABLE 2: Summary of Reaction Probabilities for HNO3 on NaCl Small Crystals and Ground Powders at 298 Ka
mass of NaCl (g) 0.65 0.87 1.19 1.54g 1.79 1.93 2.08 2.44 2.54 3.54 n.d.g,h n.d.g,h 1.97 2.32 2.98g 3.93 5.25 8.31 4.18 6.85
Figure 5. Typical data for the uptake of HNO3 (m/e ) 46) and production of HCl (m/e ) 36) on (a) 500 µm and (b) 4 µm NaCl crystals at room temperature. Aperture diameter, 9 mm; sample holder geometric surface area, 18.9 cm2; [HNO3]o = 0.1 mTorr. Salt samples were heated at 353 K for 6 (a) and 24 (b) hours under vacuum.
smallest crystals still showed some initial enhancement for the uptake of HNO3 and the formation of HCl. The reaction probability for HNO3 on NaCl was calculated using eq 2 for the reproducible plateaus (Figures 4 and 5) observed after the initial rapid uptake. Table 2 shows the results of these studies using the ∼4 and the ∼500 µm crystals, respectively, and assuming only the geometric surface area is available for reaction (this assumption is discussed in more detail below). It is seen that the results are reasonably consistent over a range of initial HNO3 concentrations, different apertures, estimated numbers of layers of salt crystals, and change in the particle size. In addition, as seen in Table 2, the reaction probabilities measured using the concentrated nitric acid (∼70%) without drying were identical, within experimental error, to those using dry HNO3 for both the 4 µm and the 500 µm crystals. This is in agreement with the observations of Fenter et al.23 and Leu et al.15 who report that their measured reaction probabilities are independent of baking of the salts or the addition of water vapor. HCl was the only gas-phase product observed. Calibration of the mass spectrometer for HNO3 and HCl under the conditions where most experiments were carried out gave relative sensitivities of ∼1.4:1, indicating from data such as those shown in Figures 4 and 5 that the yield of HCl was, within experimental error, 100%. To investigate the possibility that small amounts of surfaceadsorbed water were playing a role in the uptake, several experiments were carried out to replace surface-adsorbed water by D2O. The finely ground salt was exposed to 15 Torr D2O for about 15 min, the D2O pumped off, and exposure to D2O repeated a second time. The salt was then pumped on for approximately 5 h. Figure 6 shows the uptake of HNO3 and
typical crystal diameter (µm) b
4 4 4 4 4 4 4 4 4 4 4e 4f 500 500 500 500 500 500 500e 500f
estimated number of salt layers 40 53 73 94 110 118 127 149 156 217 >100 >100 1.0 1.1 1.5 1.9 2.6 4.1 2.0 3.4
[HNO3] (1012 molecules cm-3)c
γd (10-2)
3.0 5.2 4.7 35.1 17.2 8.7 5.6 14.3 6.5 5.4 23 13 4.5 7.5 13.5 6.2 28.6 4.9 17.8 11.1 average (( 2σ):
1.3 1.5 2.2 1.4 1.3 1.2 1.6 1.6 0.9 1.6 1.5 1.4 1.5 1.1 1.3 1.3 1.0 1.4 1.3 1.0 1.4 ( 0.6
a Unless otherwise stated, a 9 mm aperture was used. b The 4 µm powders were those prepared by grinding single crystals and were typically mixtures in the range 1-10 µm. The 500 µm were small commercially available crystals sieved to obtain a small size range from 420 to 840 µm. c Pressures g0.1 mTorr (3.2 × 1012 molecules cm-3) could be measured in the Knudsen cell. Smaller concentrations were estimated from the mass spectral signal intensities calibrated for HNO3 at P g 0.1 mTorr and assuming the observed linearity extended to smaller concentrations as well. d Calculated using geometric surface area; see text. e Using a 6 mm aperture. f Using a 4 mm aperture. g Using 70% HNO3 without distillation or drying. h n.d. ) not determined.
the simultaneous production of DCl (m/e ) 37) and HCl (m/e ) 36). Significant concentrations of both DCl and HCl are produced with relatively more DCl at the shortest reaction times. Thus, substantial amounts of surface-adsorbed D2O must have remained on the crystals after pumping, and the reaction was not the simple ion exchange reaction 1. A similar experiment in which the salt was heated while pumping for approximately 14 h after the D2O treatment gave similar results except that the initial HCl and DCl signals were approximately of equal intensity. Two additional experiments were carried out in which DNO3 was reacted with the finely ground powder, which had been pumped on for approximately 6 h without heating, or with the 500 µm small crystals, which had been heated at 353 K for 12 h. Mass spectral analysis of the DNO3 reactant in the Knudsen cell was carried out at concentrations higher than those used in the kinetic studies in order to accurately measure the parent peaks. The ratio of the parent peaks at m/e ) 64 (DNO3) and 63 (HNO3) was measured to be 5:1. In the reaction of the finely ground powder (Figure 7), DNO3 is taken up and both DCl and HCl are formed. Indeed, the initial pulse produces predominantly HCl, with DCl appearing subsequently. Figure 8 shows the reaction of DNO3 with the 500 µm crystals that had been extensively heated and pumped on prior to the reaction. As expected from the experiments with HNO3 (Figure 5), no initial rapid uptake was observed. However, during the reproducible uptake at longer reaction times, not only is DCl produced but smaller amounts of HCl are as well, in a ratio DCl/HCl = 1.8:1. Despite the extensive heating and pumping, it is apparent that some water must have remained on the surface.
Uptake of HNO3 and Nitrogen Oxides
Figure 6. Uptake of HNO3 (m/e ) 46) and production of HCl (m/e ) 36) and DCl (m/e ) 37) when HNO3 reacted with finely ground (110 µm) NaCl powders that had been previously exposed to D2O vapor to replace surface-adsorbed water (see text).
We also studied the uptake of HNO3 on single crystals with two different crystal faces, the (100) using 1 and 4 mm apertures and the (111) surface using a 1 mm aperture, and with HNO3 pressures of 1.2 and 0.1 mTorr. The top crystal surface areas in the sample dish were 13.0 cm2 for the (111) surface and 10.8 cm2 for (100). If the sides of the crystals are taken into account, the total surface areas become 20.8 and 24.8 cm2, respectively. As seen in Figure 9, there was no measurable uptake of HNO3 or production of HCl observed in either case. Studies were also carried out to measure the uptake of HCl, NO2, NO, ClNO, ClNO2, and H2O, respectively, on the smallest ground powders (1-10 µm). Under the conditions of the present experiments, no uptake of any of the gases was observed on NaCl, from which we estimate an upper limit to the uptake coefficients of ∼10-5. Finally, the uptake of HNO3 on NaNO3 was studied, since a nitrate film will be formed during the HNO3 reaction with NaCl and this may take up HNO3 at a different rate than the unreacted NaCl. Again, no uptake of HNO3 was observed using the smallest ground powders. Discussion For the uptake of HNO3 on NaCl, the shapes of the uptake curves in Figures 4 and 5 demonstrate that there is an initial rapid uptake of HNO3 and production of HCl followed by a smaller, but constant uptake of HNO3 and production of HCl. We believe that in both cases strongly bound surface water plays
J. Phys. Chem., Vol. 100, No. 37, 1996 15223
Figure 7. Uptake of DNO3 (m/e ) 46) and production of HCl (m/e ) 36) and DCl (m/e ) 37) when DNO3 reacted with finely ground (110 µm) NaCl powders (see text). The lid was closed again from ∼5700-6100 s.
a critical role. First, a comparison of Figures 4 and 5 shows that even relatively moderate heating of the salt samples while pumping decreased the extent of this initial rapid uptake. Indeed, in the case of the 500 µm crystals, the more rapid initial uptake was not observed after heating. Second, the observation of both DCl and HCl from the reaction of HNO3 with NaCl in which adsorbed surface water had first been replaced with D2O (Figure 6) establishes that it is not a simple ion exchange reaction between HNO3 and NaCl as depicted by eq 1. Third, the reaction of DNO3 with NaCl gives an initial “pulse” of HCl (Figure 7); only at longer reaction times is DCl observed. Finally, one can calculate, using gas kinetic molecular theory, the time required to completely react the surface chloride, i.e., to “saturate” the surface with a thin film of the product nitrate. Such a surface saturation has been observed in earlier XPS studies of the reaction of HNO3 with single crystals.12 Assuming the surface area available for reaction is the geometric surface area determined by the sample holder (18.9 cm2), the time to completely react the surface chloride is approximately 2 s at an HNO3 concentration of 3 × 1012 molecules cm-3 (0.1 mTorr). However, Figures 4 and 5 demonstrate that after the initial rapid uptake, a reproducible uptake of HNO3 and production of HCl subsequently continue for the time scale of these experiments, i.e., minutes to hours. One possibility is the formation of a surface nitrate layer followed by the uptake of HNO3 on the nitrate film. However, in separate experiments using finely ground NaNO3 powders,
15224 J. Phys. Chem., Vol. 100, No. 37, 1996
Beichert and Finlayson-Pitts
Figure 9. HNO3 uptake signal on NaCl (a) (100) and (b) (111) single crystals.
Figure 8. Uptake of DNO3 (m/e ) 46) and production of HCl (m/e ) 36) and DCl (m/e ) 37) when DNO3 reacted with 500 µm NaCl crystals that had been heated and pumped on for 12 h prior to reaction.
we did not observe a net uptake of HNO3. More important, the uptake of HNO3 on the NaCl was accompanied by HCl production (Figures 4 and 5), indicating that the uptake was not simply due to physical adsorption. Surface-Adsorbed Water. These results are consistent with the existence of strongly bound surface-adsorbed water, persistent even after heating and pumping. The existence of surface-bound water under these conditions is confirmed by the data in Figure 8 where the reaction of DNO3 with NaCl, which was sufficiently “dry” that it did not exhibit the initial rapid uptake, formed DCl and HCl in a ratio of ∼1.8:1. The production of HCl is thus enhanced relative to DCl in terms of the initial isotope distribution of D/H ) 5:1 measured in the reactant gases using the parent peaks at m/e ) 64 and 63. Figure 10 is a schematic of the surface with this surfaceadsorbed water (referred to hereafter as SAW) and the chemistry occurring when HNO3 is introduced into the gas phase. The SAW is expected to be near neutral in pH initially and to contain both chloride and sodium ions dissolved from the underlying salt. Under these conditions, HNO3 is taken up readily, leading to decreasing pH values as more HNO3 is absorbed. It is this process that leads to the initial rapid uptake of HNO3. Once this SAW becomes sufficiently acidified, HCl degasses. The uptake of HNO3 is then essentially an uptake into a concentrated solution containing H+, NO3-, Na+, and Cl-. As the nitrate concentration in the SAW increases, it may reach the solubility of NaNO3, at which point it is removed from the SAW by “precipitation” onto the salt surface.
Figure 10. Schematic of uptake of HNO3 in proposed SAW on NaCl surface and degassing of product HCl.
Brimblecombe and Clegg,31,32 and Tang and co-workers33 have treated the thermodynamics of HNO3 and HCl being taken up by concentrated aqueous salt solutions representative of those believed to exist in the marine boundary layer. For the dissolution of gaseous HCl into bulk aqueous solution,
HCl(g) T H+(aq) + Cl-(aq)
(3,-3)
Brimblecombe and Clegg31 report an equilibrium constant KHCl ) (mH+)(mCl-)γ(2/PHCl at 298 K of 2.04 × 106 mol2 kg-2 atm-1, where “m” represents concentration in molal units and γ( represents the mean activity coefficient. By use of the extended Debye-Hu¨ckel treatment discussed in detail by Brimblecombe and Clegg,31 γ( ) 0.59 at mCl- ) 6.3 m, which is the solubility limit of NaCl at room temperature, so γ(2 ) 0.35. As a first approach, the activity coefficients and equilibrium constants recommended by Brimblecombe and Clegg31,32 are applied to the SAW proposed here. It is important to recognize that this is a crude approximation, since the nature of the SAW is not known. It is expected (e.g., see refs 34-40) that water adsorbed on the surface under these conditions is associated
Uptake of HNO3 and Nitrogen Oxides
J. Phys. Chem., Vol. 100, No. 37, 1996 15225
TABLE 3: Comparison of Reaction Probability for the Reaction of HNO3 on NaCl in These Studies to Literature Values
a
NaCl source/size
reaction probability at 298 K
ref
method
single (100) crystals ground powders small crystals single crystals powders, small crystals, spray- coated sample, polished and depolished window faces ground powders, small crystals
(4.0 ( 2.0) × 10-4 (2.8 ( 0.3) × 10-2 (1.3 ( 0.4) × 10-2 g(2.4 ( 0.6) × 10-3 (2.0 ( 1.0) × 10-2
Laux et al.12 Fenter et al.23 Leu et al.15
XPS/UHV Knudsen cell/EI-MSa flow system/CI-MSb
Fenter et al.21
Knudsen cell/EI-MS
this work
Knudsen cell/EI-MS
(1.4 ( 0.6) x 10-2
EI-MS ) electron impact mass spectrometry. CI-MS ) chemical ionization mass spectrometry. b
with surface defects, steps, or edges. For example, Barraclough and Hall,35 using infrared spectroscopy, showed that it was necessary to heat NaCl crystals to ∼623 K to completely remove water from the surface. In addition, about 25% of the NaCl surface sites were active for water adsorption. This SAW is not expected to behave like a bulk solution. However, this approach at least provides an initial framework to examine the processes occurring in the SAW. We first assume that the concentration of Cl- in the SAW is determined by the solubility of NaCl in water, i.e., that the SAW is saturated with NaCl. At room temperature, this corresponds to a concentration of 6.3 molal (m). For a typical experiment such as those shown in Figure 4, one can calculate the pH of the SAW. During the later, constant uptake of HNO3 and evolution of HCl, the HCl concentration corresponding to Figure 4b is ∼8.2 × 1011 molecules cm-3, or 3.3 × 10-8 atm. By use of the value of KHCl and the activity coefficients reported by Brimblecombe and Clegg,31 mH+ is calculated to be ∼3.1 × 10-2 m and the activity aH+ ) 1.8 × 10-2, corresponding to a pH of 1.7, at the saturation concentration of Cl- of 6.3 m and the measured pressure of HCl in the Knudsen cell. For HNO3, the uptake and dissociation in the SAW is represented by eq 4,-4:
HNO3(g) T H+(aq) + NO3-(aq)
(4,-4)
The equilibrium constant31-33 is given by KHNO3 ) (mH+)(mNO3-)γ(2/PHNO3 ) 2.66 × 106 mol2 kg-2 atm-1. The mean molality, m( ) (m+m-)1/2 ) [(mH+)(mNO3-)]1/2 can be estimated using (mH+) ) 3.1 × 10-2 m (based on the calculation for HCl above) and (mNO3-) ) 17.0 m. The latter is the nitrate concentration in the SAW as determined by its equilibrium reaction with Na+ to form NaNO3. The concentration of Na+, however, is determined by the solubility of NaCl to be 6.28 mol L-1. Hence, [NO3-] ) KspNaNO3/[Na+], where the solubility product constant for NaNO3, KspNaNO3 ) 107 mol2 L-2, is calculated from the room temperature solubility of NaNO3 of 10.4 mol L-1. By use of these concentrations of H+ and NO3-, the mean molality is calculated to be m( ) 0.725 m. The value of γ(2 at this mean molality can then be calculated to be 0.516 using the extended Debye-Hu¨ckel equation discussed in detail by Brimblecombe and Clegg31,32 and Tang et al.33 With a 3.1 × 10-2 m concentration of H+, a typical (Figure 4b) equilibrium pressure of gaseous HNO3 of 8.9 × 10-8 atm in the Knudsen cell, and the equilibrium constant KHNO3 and activity coefficient cited above, a nitrate concentration of 14 m is calculated. Given the approximate nature of these calculations, this is within experimental error of the nitrate concentration determined by the solubility of NaNO3. Thus, as HNO3 is taken up into the SAW, it can precipitate out on the salt surface as NaNO3 or possibly reevaporate as HNO3. This model suggests that there should be an induction period for the production of HCl, i.e., it should not degas until the SAW has become sufficiently acidic. At the concentrations used in our experiments, we did not observe such an induction period.
However, Rossi and co-workers23 do report observing an induction period at HNO3 concentrations approximately 50 times smaller than those we were able to use in the present study. The shape of the uptake curves and production of HCl, with a rapid initial uptake, followed by a slower, constant uptake (e.g., Figure 4) suggests there are two kinds of sites holding surface water. The first, responsible for the initial large cusp in the uptake, is apparently destroyed upon reaction with HNO3, leaving the second type of site, which continues to react at a constant rate at longer reaction times. The nature of these two sites is not clear. One possibility is that the first involves surface OH- formed by dissociation of water on defect sites as reported, for example, by Ewing and co-workers.40 Although the number of such defect sites initially is expected to be small,40 Laux et al.13 have shown recently that reaction of NaCl with HNO3 actually generates surface defects that enhance the dissociative adsorption of water, forming surface OH-. Water would be expected to cluster at such a site owing to both the polarity and to the formation of hydrogen bonds. The first molecules of nitric acid taken up into the water cluster would neutralize the hydroxide. The DNO3 or HNO3 taken up subsequently would acidify the cluster, leading to HCl and/or DCl outgassing. Since there are more H atoms than D in the water/OH- cluster, HCl is produced initially, consistent with Figure 7. With the removal of OH- as the central species in this water cluster, the binding of the water to the salt surface is weakened and the water originally associated with the OH- is pumped off. This leaves only the second type of adsorbed water, which is sufficiently strongly adsorbed that it takes up HNO3 and degasses HCl without desorbing. The observation that the same behavior and uptake coefficient were observed using 70% HNO3/H2O in place of dry HNO3 suggests that these two types of sites are, in effect, saturated with water so that no net uptake from the gas phase occurs. It is also consistent with our observation that even the finely ground powder does not show a measurable net uptake of H2O. Reaction Probability. Table 3 compares the reaction probability obtained in these studies as determined from the data after the initial rapid uptake (assuming only the geometric surface area is available for reaction) to those reported in previous studies. Our value using powders and small crystals is in good agreement with the Knudsen cell studies of Fenter et al.21,23 and the flow tube studies of Leu et al.15 This reaction probability may in fact reflect a combination of uptake into the SAW, reevaporation of HNO3 into the gas phase, and reaction to form HCl. As discussed by Quinlan et al.,25 if the reactant gas can reevaporate from the surface, the measured reaction probability (γ) reflects the fraction of the gas taken up by the surface (R), which reacts rather than reevaporates, i.e., γ ) R [krxn/(krxn + kevap)], where the k’s are the rates of reaction and evaporation, respectively, and R, the mass accommodation coefficient, is the fraction of gas-surface collisions leading to uptake of the gas. For the SAW model proposed here, γ ) (1.4 ( 0.6) × 10-2. Van Doren and coworkers41 measured the mass accommodation coefficient for
15226 J. Phys. Chem., Vol. 100, No. 37, 1996
Figure 11. Schematic of salt crystals and HNO3-surface collisions under Knudsen cell conditions. +0.021 HNO3 on water droplets to be R ) 0.193-0.015 at 293 K. If this also applies to the uptake of HNO3 on the high ionic strength SAW, then krxn/(krxn + kevap) ) 7.3 × 10-2, i.e., 7.3% of the HNO3 initially taken up into the SAW reacts, with the rest in effect evaporating. The calculation of reaction probabilities from studies of the uptake of gases by solid powders has been somewhat controversial owing to the uncertainty in the amount of surface available for reaction.15,21,23 In our current studies, the measured reaction probability for HNO3 depends only on the geometric surface area and is independent of the particle size or number of layers of salt particles, in agreement with the observations of Rossi and co-workers21,23 but not those of Leu et al.15 in a flow system. One explanation is that the structure of the solid surface composed of individual small crystals of NaCl is such that it does not have deep crevices or “pores” between the crystals. This is reasonable, given that it is highly unlikely that salt will stack in the sample holder in a uniform manner to give deep pores. A random stacking of interlaced particles such as that shown schematically in Figure 11 seems much more likely. The porosity of the surface, i.e., the fraction of the geometric surface that is open, can be calculated from the volume of the sample holder (7.5 cm3), the mass of the salt required to just fill it (10.7 g for the 500 µm particles), and the bulk density of the salt (2.165 g cm-3). For the 500 µm particles, the porosity is calculated to be 0.35. Thus, on average, for every two salt particles on the surface, there is a void area equivalent to that of one salt particle, i.e., relative to the size of the salt particles themselves, there are large openings on the salt surface, as shown schematically in Figure 11, rather than narrow, deep pores that can effectively trap the gas. Take the cases in Figure 11 where molecules enter an indentation in the surface and collide with salt crystals in the second layer. The opening is sufficiently wide that collisions with one or possibly two sides of the well may occur before the molecule is scattered back out. Under Knudsen cell conditions at a pressure of 0.1 mTorr, the mean free path is ∼20 cm, much larger than the size of the well. Thus, the scattering out of the well occurs without collisions with other gas molecules. The total available surface area for reaction for the molecule theoretically includes the four sides of the well and also the bottom. However, in practice, it is expected that elastic collisions will generally occur with one or perhaps two sides of the well before the molecule bounces back out into the volume of the reactor. This will give a maximum increase of 33-67% in the effective surface area and a corresponding decrease in the calculated value of the reaction probability for this portion (about a third) of the surface, which is “open”. Averaging over the entire surface, the reaction probability would be lowered to ∼77-83% of that for a flat, nonporous surface. Because of this and other possible systematic errors, we estimate the uncertainty in our reaction probability to be a factor of 2. In their flow tube studies, Leu and co-workers15 apply a correction to the geometric surface area using a pore diffusion
Beichert and Finlayson-Pitts model developed earlier22 for application to ice surfaces. Their reaction probability for the HNO3 reaction, (1.3 ( 0.4) × 10-2, has been corrected by a factor of ∼5 for increased available surface area in the “pores” between salt crystals. In a flow system, corrections for increased effective surface area of the salt compared to the geometric surface area are reasonable because the direction of pumping in a flow system will tend to pull the gas through the salt from the upstream to the downstream end, exposing the gas to a total surface area larger than the geometric surface area. In addition, since the flow tube studies are carried out at 0.3-0.5 Torr in He, the mean free path is much smaller, on the order of the indentations in the salt surface. This is in contrast to Knudsen cell conditions where the gas is not pumped across the salt and the mean free path is very large compared to the dimensions of salt crystals and surface indentations. An alternate explanation for the lack of dependence on the number of layers or particle size is that the reaction is fast relative to diffusion into the crevices between the particles.42 In this case, collisions near the top of the crevice or pore lead to the uptake of the gas so that the reactant gas in effect never sees the subsurface area in the crevices. Rossi and co-workers,21 for example, propose that HNO3 is taken up by the salt on every collision to form an adsorbed species that subsequently reacts. This uptake could be into the SAW proposed here. However, if this hypothesis is correct, then the same should apply to the studies of Leu and co-workers,15 and their reaction probability, uncorrected for internal surfaces, would be 6 × 10-2, a factor of 3-4 larger than those measured in Knudsen cell studies. Rossi et al. also report21 that in contrast to the HNO3 reaction, the reaction of N2O5 with NaCl does depend on the number of salt layers, with a rapid falloff in the uptake coefficient in going from several salt layers to one layer. Given the evidence reported here for the importance of surface-adsorbed water, there is, however, an alternative explanation for the N2O5 behavior. First, it is difficult to heat salt evenly because of its poor thermal conductivity. The bottom layer would be expected to be heated most effectively and, hence, contain less adsorbed surface water, leading to a less efficient uptake of the gas. In the case of multiple salt layers, if there is a thermal gradient from the bottom to the top of the salt, water could desorb from the bottom onto the top layers where reaction occurs. The decrease in the uptake coefficient at one layer may therefore reflect a decrease due to decreased surface water compared to the case with multiple layers. Second, with only one layer on average, the surface may be a combination of some multilayers and some with no particles. We did not observe uptake of HNO3 on single crystals, for which there are two possible explanations. The most likely explanation is that the relatively defect-free NaCl surface does not readily hold adsorbed water, and in the absence of SAW or reactive surface defects, the uptake of HNO3 is very slow. An alternative, but less likely, explanation is that the surface reacted so rapidly that the initial uptake leading to saturation was not observed. In this case, the reaction probability without the SAW must be quite high to saturate the surface so quickly; SAW cannot be present on these single crystals, since continuing uptake of HNO3 and production of HCl similar to that seen for the powders were not observed. The absence of SAW on single crystals under these experimental conditions is in agreement with studies reported in the literature on the lack of uptake of water onto single crystals at room temperature.36,38 At the present time, we cannot distinguish between these two possibilities. The first explanation is, however, reasonable based on earlier XPS studies where an uptake coefficient of (4 ( 2)
Uptake of HNO3 and Nitrogen Oxides × 10-4 was measured for dry HNO3 uptake by single crystal (100) NaCl, approximately 2 orders of magnitude less than that on powders. Similarly, the reaction probability for N2O4 on finely ground NaCl powders,10 (1.3 ( 0.6) × 10-4, is also about 2 orders of magnitude larger than that measured recently on single (100) crystals.9 It is especially noteworthy that Peters and Ewing9 report a dramatic increase (again about 2 orders of magnitude) in the reaction kinetics in the presence of 9.5 mbar water vapor, which is sufficiently high to generate adsorbed water even on single crystals. Leu et al.15 also report an enhancement in the reaction of N2O5 with NaCl if there is water on the salt surface. Uptake of Other Gases on NaCl. Finally, no uptake of NO2, NO, ClNO, ClNO2, H2O, or HCl was observed using the finely ground powders from which a maximum uptake coefficient of ∼1 × 10-5 can be calculated. The lack of uptake of NO2 at the smaller concentrations used in this study is consistent with the earlier kinetic studies9,10 at higher concentrations discussed above. The lack of uptake of H2O is in agreement with the results of Rossi and co-workers23 who report an upper limit of 2 × 10-4 on finely ground powders. This lack of measurable net uptake of water vapor may reflect equal rates of uptake and evaporation from the sites holding adsorbed water on the surface rather than no uptake at all. Both Leu and co-workers15 and Rossi and co-workers23 report an initial uptake of lower concentrations of HNO3 by NaNO3 powders followed by a rapid “saturation” of the surface. Rossi and co-workers23 also report efficient uptake of HCl, with an uptake coefficient of 3 × 10-2, and show that HCl is displaced by subsequent exposure to gaseous HNO3. These observations are consistent with our observation of no measurable net uptake of HNO3 on NaNO3 or HCl on NaCl. If a SAW is involved, saturation would occur on a time scale too short for us to observe with the higher concentrations used in these studies. In the earlier studies15,23 at lower concentrations, these gases would be taken up into the SAW until it is sufficiently acidified that equilibrium between uptake and degassing is reached. In the case of HCl, subsequent uptake of HNO3 would result in degassing of the HCl as shown by the model in Figure 10 and the calculations of Brimblecombe and Clegg.31,32 Atmospheric Implications. The results presented here suggest that even small amounts of surface-adsorbed water play a key role in the reactions of NaCl with gases such as HNO3. Under atmospheric conditions, even at low relative humidities where airborne sea salt particles are solids with adsorbed water rather than concentrated aqueous solutions, it appears that the uptake of HNO3 and other gases can be thought of as dissolution into a concentrated salt solution, causing acidification, followed by outgassing of HCl. Surface NaCl continuously dissolves to maintain this concentrated chloride solution, and at large uptakes of HNO3, precipitation of NaNO3 occurs to maintain the uptake of gas-phase HNO3. Studies are currently underway to test the applicability of this model to other heterogeneous reactions of NaCl. It is stressed that this SAW is not believed to be a layer of water covering the entire salt surface but rather water that is strongly bound to certain defect sites. Conclusions These studies of the reaction of HNO3 with NaCl dramatically illustrate the effects that a very small amount of surface-adsorbed water has on the heterogeneous reaction. The initial rapid reaction, the subsequent constant uptake long after the expected “saturation” of the surface, and the results of experiments using deuterated species all support the existence of strongly bound
J. Phys. Chem., Vol. 100, No. 37, 1996 15227 surface-adsorbed water on small crystals and finely ground powders, even after prolonged heating and pumping. This alters our fundamental understanding of the nature of these surface processes and suggests that except for relatively defect-free large (100) single crystals, they are better regarded as a concentrated salt solution on the surface into which HNO3 is absorbed. This causes acidification of the layer, until HCl degasses when a sufficiently low pH is reached. Subsequently, a balance is reached between dissolution of NaCl into the SAW, absorption of HNO3 from the gas phase, degassing of the HCl from the SAW, and precipitation of the nitrate as NaNO3. Although the nature of the SAW is not known, based on previous studies in the literature, it is likely to be bound at surface defects, steps, and edges. Work is underway to try to define the nature of the SAW. Acknowledgment. We gratefully acknowledge the financial support of the National Science Foundation (Grant No. ATM9302475), the Department of Energy (Grant No.94ER61899), and the Joan Irvine Smith and Athalie R. Clarke Foundation. We are also grateful to Professor M. Tolbert for comments on the design and testing of the Knudsen cell, to Professor John Hemminger, Dr. J. N. Pitts Jr., and Dr. M.-T. Leu for helpful discussions and comments on the manuscript, Professor George Ewing for helpful discussions and providing reprints prior to publication, and Dr. Sarka Langer for assistance with some experiments. Appendix The corrections to the reaction probability for the N2O5 on H2SO4/H2O were calculated for possible interferences due to NO2 or HNO3 in the following manner. In the case of NO2, it is assumed that the contributions of NO2 to the signal intensities at m/e ) 46 and 30 are relatively small prior to reaction. However, after the lid has been opened and N2O5 is removed to a large extent by reaction with H2SO4/H2O, the contribution of NO2, which is not rapidly taken up,24 becomes significant. Thus, the ratio of the signals at 46 and 30 is no longer ∼2:1 as expected for N2O5, and the relative intensity of the m/e ) 30 peak has increased. The measured signal intensity at m/e ) 46 after reaction, Ir46, is then a weighted combination of signals due to N2O5 and NO2 given by
Ir46 ) 100([N2O5]fN2O5) + 40([NO2]fNO2)
(a)
) 100C + 40D where C ) [N2O5]fN2O5 and D ) [NO2]fNO2 and the “f” factors represent the intrinsic MS sensitivity for N2O5 and NO2, respectively. The multipliers of 100 and 40 represent the relative signal intensities from scans such as those in Figure 2. Similarly, for m/e ) 30,
Ir30 ) 50([N2O5]fN2O5) + 100([NO2]fNO2)
(b)
) 50C + 100D Since the two signal intensities, Ir46 and Ir30, are measured, these simultaneous equations can be solved for C and D. The corrected signals at m/e ) 46 and m/e )30 due to N2O5 alone are then 100C and 50C, respectively. The contribution of NO2 to the signals at m/e ) 46 and 30, respectively, after reaction can then be compared to the total signals observed before reaction to assess whether the assumption that they are small is valid. In all cases, the contribution at 46 is
