J. Phys. Chem. 1992,96,4979-4985
References and Notes (1) Erickson, R. E.; Yates, L. M.; Clark, R. L.; McEwen, D. Armos. Enuiron. . . 1977. 11. 813-817. (2) Penkett; S. A.; Jones, B. M. R.; Brice, K. A.; Eggleton, A. E. J. Armos. Emiron. 1979, 13, 123-137. (31 Calvert. J. G.: Lazrus. A.: Kok. G. L.; Heikes, B. G.; Waleaa, - J. G.; Lind,'J.; Cantrell, C. A. Nature 1985, 317, 27-38. (4) Chameides, W.L.; Davis, D. D. J . Geophys. Res. 1982,87,4863-4877. (5) Solomon, S . Reu. Geophys. 1988, 26, 131-148. (6) Schwartz, S . E. In Chemistry of Multiphase Atmospheric Systems; Jaeschke, W.,Ed.; NATO AS1 Series; Springer-Verlag: Berlin, 1986; Vol. G6. p 415. (7) Schwartz, S. E. Atmos. Enuiron. 1988, 22, 2491-2499. (8) DeMore, W. B.; Sander, S. P.; Golden, D. M.; Molina, M. J.; H a m p son, R. F.; Kurylo, M. J.; Howard, C. J.; Ravishankara, A. R. Chemical Kinetics and Photochemical Data for Use in Stratospheric Modeling, Evaluation 9 JPL Publication 90-1; California Institute of Technology: Pasadena, $A: 1990. (9) Hanson, D. R.; Ravishankara, A. R. J . Geophys. Res. 1991, 96, 5081-5090. (10) Gardner, J. A.; Watson, L. R.; Adewuyi, Y. G.; Davidovits, P.; Zahniser, M. S.; Worsnop, D. R.; Kolb, C. E. J . Geophys. Res. 1987, 92, 10887-10895. (1 1) Jayne, J. T.; Gardner, J. A.; Davidovits, P.; Worsnop, D. R.; Zahniser, M. S.; Kolb, C. E. J. Geophys. Res. 1990, 95, 20559-20563 and references cited therein.
4979
(12) Mozurkewich, M.; McMurry, P. H.; Gupta, A.; Calvert, J. G. J. Geophys. Res. 1987, 92, 4163-4170 and references cited therein. (13) Hoffmann, M. R. Armos. Ewiron. 1986, 20, 1145-1154. (14) Tang, I. N.; h, J. H. In The Chemistry of Acid Rain; Sources and Atmospheric Processes; Johnson, R. W., Gordon, G. E., Eds.; ACS Symposium Series 349; American Chemical Society: Washington, DC, 1987; pp 109-1 17. (15) Hanson, D. R.; Burkholder, J. B.; Howard, C. J.; Ravishankara, A. R. J . Phys. Chem., following paper in this issue. (16) Ridley, B. A. Armos. Ewiron. 1978, 9, 27-34. (17) Danckwerts, P. V. Gas-Liquid Reactions; McGraw-Hill: New York, 1970. (18) Handbook of Chemistry and Physics, 60th ed.; Weast, R. C., Ed.; CRC Press: Boca Raton, FL, 1979. (19) Geankoplii, C. J. Transport Processes and Unit Operations, 2nd Ed.; Allyn and Bacon: Boston, 1983. (20) Brown, R. L. J. Res. Natl. Bur. Stand. (US.)1978, 83, 1-8. (21) Howard, C. J. J. Phys. Chem. 1979,83, 3-9. (22) Mozurkewich, M. Private communication. (23) Goldberg, R. N.; Parker, V. R. J. Res. Narl. Bur. Srand. (US.)1985, 90, 341-358. (24) Kosak-Channing, L. F.; Helz, G. R. Ewiron. Sci. Technol. 1983, 17, 145-1 49. (25) Roth, J. A.; Sullivan, D. E. Ind. Eng. Chem. Fundam. 1981, 20, 137-140.
Measurement of OH and H02 Radical Uptake Coefficients on Water and Sulfuric Acid Surfaces David R. Hamon,*.+ James B. Burkholder,t Carleton J. Howard, and A. R. Ravishankaras NOAAIERL. Aeronomy Laboratory, RIElAL2, 325 Broadway, Boulder, Colorado 80303 (Received: June 10. 1991; In Final Form: February 17, 1992)
A wetted wall flow tube reactor was used to measure the uptake coefficients, y, of OH and H 0 2 on pure water at 275 K and 28% w/w sulfuric acid at 249 K. The uptake coefficients are lower limits to the mass accommodation coefficients, a, and the y were determined to be 0.0035 for OH and >0.01 for HOz on liquid water and >0.08 for OH and s0.05 for HOz on the sulfuric acid solution. In addition, the binary diffusion coefficients for H 0 2 and OH in water vapor were estimated to be 79 f 8 and 116 f 20 Torr cm2 8,respectively, at 275 K. The determination of uptake coefficients using wetted wall flow tube reactors is described including the effect upon y of gas-phase diffusion, solvation, and reaction within the liquid.
Introduction In the preceding paper in this issue, Utter et al.' explained the motivation for measuring the mass accommodation coefficients of trace atmosphericspecies onto liquids and solids. To summarize briefly, some reactions which are negligibly slow in the gas phase occur much more rapidly in/on condensed media. As the condensed-phase reactants are made in the gas phase and then transported into the liquid or solid phase, transport of gases into atmospheric particles can be the rate limiting step for these reactions. The mass accommodation coefficient, a,is the parameter that can limit transport at the gas-condensate interface. Utter et al.' described our experimental approach, where a wetted wall flow tube2 was used to measure the mass accommodation coefficient of ozone onto liquid surfaces. Using much the same apparatus, we report measurements of the uptake coefficients of OH and H 0 2 on liquid surfaces. These species could play important roles in the liquid-phase oxidation of acid precursors. Their transport into the condensed phase could account for a significant loss for these radicals from the gas phase and could lead to formation of peroxide^.^ No measurements of a To whom correspondence should be addressed. 'NOAA/NRC Post-Doctoral Research Associate. *Also affiliated with the Cooperative Institute for Research in Environmental Sciences, University of Colorado, Boulder, CO. 1 Also affiliated with the Department of Chemistry and Biochemistry, University of Colorado, Boulder, CO.
for OH and H 0 2radicals on liquid water have been reported, and such measurements would improve our understanding of the transport process. If a surface becomes partially saturated with a species, a measured loss rate may not be representative of the gross transport that occurs into the liquid. To avoid this, a chemical scavenger is dissolved in the liquid so that the species is rapidly removed and transport back from liquid to gas is minimized. The value obtained under these conditions represents the reaction probability or uptake coefficient, 7,for that species and is taken to be a lower limit to the mass accommodation coefficient. Although measurementsof the mass accommodation coefficient of OH on aqueous surfaces have not been previously reported, Baldwin and Golden4 measured y = 5 X lo4 for OH on -96% sulfuric acid at 298 K. Gershenzon et al.s report very high values for y of OH on ice at 253 K and on -96% sulfuric acid at 298 K, 0.4 and 1, respectively. The loss of H 0 2 on small NH4HS04and LiN03 aerosols (-0.1-pm diameters) has been investigated by Mozurkowich et a1.6 who report a 1 0.2. Their aerosols were concentrated solutions and doped with CuS04 in order to scavenge H 0 2 in the liquid. A critical process which can limit transport in our system, as well as in the atmosphere, is gas-phase diffusion. Therefore, knowledge of the diffusion coefficients of gases is important. The diffusion coefficients of polar gases, such as OH and HOz, can be used to test the calculated transport properties of systems that
0022-365419212096-4979$03.00/00 1992 American Chemical Society
4980 The Journal of Physical Chemistry, Vol. 96, No. 12, 1992
have strong dipole-dipole interaction^.^ Very few measurements have been reported for the diffusion coefficients of polar gases and polar gas mixtures,8 and the procedures for calculating the collision integrals for this process have not been thoroughly tested. We compare the binary diffusion coefficients inferred by our results to values calculated with a model which incorporates dipole-dipole interaction^.^ In this work we examine the details of the interpretation of data obtained using the flow tube method. Experiment
The experimental setup, except for the present radical production and detection schemes, has been described in the preceding paper' in this issue. The uptake coefficients were determined by measuring the loss rate of a species as a function of exposure to a liquid surface. The liquid was slowly flowing on the inner wall of a cylindrical flow tube, and exposure time was varied by changing the position of the injector through which the radicals were introduced into the main gas flow. All measurements reported here were performed over either deionized water (resistivity of 18 MQ cm) at 274.5 K or 28%w/w sulfuric acid at 249.5 K. A few HOZloss measurements were also conducted over water containing M CuS04. For a nominal liquid volume flow rate of 1 cm3 s-I, the thickness of the liquid on the flow tube wall was 0.2 mm for water and 0.3 mm for the sulfuric acid and the average speed of the fluid flow down the tube was 7.4 cm s-l for HzO and 4.4 cm s-l for the sulfuric acid. The Reynold's number for the liquid flow, NR,(fh),was 35 for the water and 8 for the sulfuric acid. Rippling is e ~ p e c t e dwhen , ~ no surfactant is present, for Reynold's numbers greater than 25 and was observed for pure water. The sulfuric acid flowed without noticeable rippling. The ripples appeared in the form of small waves in the water surface spaced at intervals of about 10 cm. Although they can have no significant effect on the wall surface area (on the order of a few percent), they can enhance both gas-phase and liquid-phase tran~port.~J~ For experimentswith the 28%acid solution, the heat exchanger and the salt-ice collection vessel were not used. We used a second glass bulb, cooled in a similar manner to the solution storage bulb, to collect the sulfuric acid as it flowed out the reactor. The OH or H 0 2 entered the flow tube through a double-walled glass injector (the outer tube, 9-mm o.d., the inner tube, 5-mm 0.d.). The OH was produced by the reaction of excess NOz with H atoms from a microwave discharge of a H2-He mixture at the top of the injector. The source reaction occurred at the bottom of the injector where the H atoms from the inner tube were mixed with NOz. HOz was produced by reacting F atoms with excess H202. The initial OH and HOz concentrations were between 0.5 X 10" and 3 X 10" molecules ~ m - ~[OH] . was estimated by reacting H atoms, [HI = 10"-10'2 molecules ~ m - with ~ , known amounts of NOz until the OH signal no longer increased. The reaction between H and NOz is fast; thus when a stoichiometric amount of NOz has been added, the OH signal no longer increases and [OH] N [N0210at the end point. This also provides a calibration for HOZsince it is converted to OH, as described below. The gas-phase second-order loss of OH and HOz are not siflicant for concentrations less than 10l2molecules for our pressure and temperature conditions." The small gas-phase loss for OH via reaction with NOz was estimated, using the recommended rate coefficient," to be generally less than 2%of the total loss rate. OH was detected using the pulsed laser-induced fluorescence method.I2 HOz was detected by converting it to OH just before the detection region via the reaction HOz NO NO2 OH. A pulsed dye laser, operating at 20 Hz and frequency-doubled to 282 nm, was used for excitation, and the fluorescence was detected at 308 f 4 nm with a photomultiplier tube (PMT), bandpass filter combination. The PMT signal was accumulated using a gated charge integrator, whose output was read by a computer. The computer also recorded the injector position and flow tube pressure. Carrier gas flow speeds used in this work were much larger than those used by Utter et al.; however the fractional pressure drop through the measurement region was small: less than 1% over
+
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Hanson et al. the water solution and less than 4%over the sulfuric acid solution assuming laminar flow. A heated water reservoir was connected to the system through a metering valve to humidify the helium carrier gas before it entered the flow tube. Otherwise, at high gas flow rates, the rapid evaporation of water in the flow tube entrance caused the liquid surface to freeze. This observation confirmed our expectation of rapid equilibrium between the flowing gas stream and the liquid surface. Uniform temperature and hence water pressure were maintained by rapid circulation of fluid through a temperature regulated bath and flow tube jacket, AT 5 0.5 OC. Complete saturation of the carrier gas with water vapor at the measured flow tube temperature was essential to accurately calculate the average flow speed in the measurement region. The mixing time, 'T-, for two gases in a cylinder of radius t i s roughly r2/5Dc.'3 This is about 5 ms using a diffusion coefficient 0, for HzO in He of 60 cm2s-l at 275 K and a pressure of 9 Torr.8 The distance covered in this amount of time is 20 cm at a flow speed of 4OOO cm s-I. The humidifkd camer gas was exposed to a length of at least 40 an of liquid before it entered the measurement region at the bottom of the reactor. To test whether saturation was complete, the rate at which water vapor was added to the helium gas before it entered the flow tube was varied. Two observations were noted as water vapor was added to the carrier gas: the temperature of the cooling fluid exiting the top of the flow tube increased 1-2 OC, while the total pressure-measured at the bottom of the flow tube-remained constant, and there was no detectable change in the measured loss rate coefficients. The temperature of the liquid increased at the point where the added water vapor condensed; however this temperature increase did not extend into the measurement region as the measured loss rate coefficients and the total pressure were not affected. Thus we conclude that the helium flow was fully saturated with water vapor at the flow tube temperature, as given by the thermocouple at the bottom of the flow tube. Temperature measurements at the top and bottom of the flow tube agreed to within f0.5 OC at low camer gas flow rates, i.e., when there was very little evaporation or condensation at the liquid surface. The HzO partial pressures were taken to be the vapor pressures over the liquids: 5.0 f 0.2 Torr over pure water at 274.5 f 0.5 KI4 and 0.55 f 0.05 Torr over the 28%w/w sulfuric acid solution at 249.5 f 1 K. The helium partial pressure was between 1 and 4 Torr. Vapor pressures over sulfuric acid solutions have not been investigated in this temperature region. The HzO vapor pressure over the liquid solution will, however, be very close to the H 2 0 pressure over the solid that freezes out of this solution, as the freezing point of the H$04 mixture, about 247 K,lS was very close to our operating temperature. This solid is a very dilute HzS04-in-icesolid solution, and the HzOpressure over the solid solution is very close to that over pure ice. Since we did not operate at exactly the freezing point of the solution, the HzO pressure of our liquid mixture will be slightly lower than the ice pressure at the operating temperature. It can be shown, using the Clausius-clapeyron equation and the enthalpies of vaporization of water from 30% sulfuric acidI6 and from ice,14 that the HzO partial pressure over the solution at 249.5 K is within 3.5% of the ice pressure at this temperature. Therefore, we take the HzO pressure over the sulfuric acid to be the H20ice pressure14at 249.5 K which is 0.55 Torr. The average gas flow speed in the measurement region ranged for the from 400 to 5000 an s-l. The Reynold's number, NRC,Io average helium-water vapor flow was between 40 and 400 for the 275 K measurements. Turbulence does not set in for NRcless than 2100, but the distance required for the development of laminar flow, given by -O.lrNRe,lowhere r is the radius of the tube, becomes an appreciable fraction of the length of the tube for N R ~ on the order of 200. Laminar flow is perturbed by the presence of the movable injector, and the distance to re-establish laminar flow in the measurement region below the injector also increases as NRcincreases. The analysis used to correct for gas-phase diffusion is based upon the assumption of laminar flow. This condition is met as NRe 0. The gas flows at 249 K had smaller
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OH and H 0 2 Radical Uptake on H 2 0 and H2S04Surfaces
The Journal of Physical Chemistry, Vol. 96, No. 12, I992 4981
I
40 1 0 )
40
6 4
-> -E 5
2t
>
1
-10 C 200 rtL,
0
IO
20
30
40
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4b
5L
Figure 2. Same as in Figure 1 for HO2.
250 OH
200
1000
"0
2000 3300 4000 Flow S p e e d h s-0
5000
1
6000
300 L 250
0
t 50
o 0t '
"
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I
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'
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2000 3000 4000 Flow Speed (cm s ' ' )
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Figure 3. First-order loss rate coefficients determined for OH (a) and H02 (b) plotted as functions of the average flow speed of the carrier gas. The squares are data from measurements over sulfuric acid at 249 K while the circles are measurements over water at 274.5 K. The filled circles in (b) are HOz loss measurements taken over a lo-' M CuSO,
solution. exposure time, were fitted to a straight line using a weighted least squares procedure to calculate the first-order loss rate coefficient. The weights were scaled to the statistical uncertainty of each data point. A systematic variation of the measured first-order rate coefficients over liquid water with carrier gas flow speed was observed. In Figure 3 are shown the measured first-order loss rate coefficients, k,, vs average gas flow speed for OH (a) and HOz (b) over water and sulfuric acid. Measurements over both pure water and M CuS04solution are included in Figure 3b because rippling
4982 The Journal of Physical Chemistry, Vol. 96, No. 12, 1992
Hanson et al.
TABLE I: Summary of the Experimental Results, Diffusion Coefficients That Were Used,and Calculated Uptake Coefficients experimental conditions 5 Torr of H 2 0 with 4 Torr of He: water at 275 K 0.55 Torr of H 2 0 with 1 Torr of He: 28% H2S04at 249 K
species i OH H02
OH
H02
k,/s-'
yplusb
31"
0.0013c 0.0017'' 0.011 0.010
30" 264 183
-%n2o/ (Torr cm2 s-I) 116 86 94 70
ai-ncl
(Torr cm2 s-I) 658 485 536 409
DCl
(cm2 s-') 20 15 130 97
ycOn 0.0035 >0.01
>0.08 >0.05
" k,
values are from Figure 3 and obtained from the fitted line extrapolated to zero flow speed for the measurements over pure water at 275 K. byplug is the observed uptake coefficient calculated assuming no radial concentration gradients exist, Le., using the plug-flow approximation. C A larger value of yplug= 0.004 was observed for OH; see the text for details. d A larger value of yplug= 0.0033 was observed for HO,; see the text.
was present in both liquids, and, as demonstrated below, gas-phase transport of H 0 2 is identical for both cases. As mentioned before, nonlaminar flow could have been present due to the ripples on the liquid surface and to the high N R in~ the water measurements, increasing the rate of transport to the walk beyond what gas-phase diffusion would permit. This explanation is supported by the measurements at 249 K which show much less flow speed dependence: the Reynold's number and entrance length for the measurements at 249 K were 5-10 times lower than at 274.5 K, and there was no rippling of the liquid surface. The first-order loss rate coefficients were fitted as a linear function of the flow speed using a least squares procedure. Assuming the observed speed dependence is due to nonlaminar flow, the unperturbed wall loss rate coefficients were then taken to be the values extrapolated to zero flow speed, i.e., zero NRe. The slopes of the lines were statistically insignificant for the OH and H 0 2 measurements over the sulfuric acid, so we averaged the data to obtain the reported first-order loss rate coefficients. The results of the OH and H 0 2 loss measurements are summarized in Table I. The y we determined, and thus the lower limits to a,are 0.0035 and 20.08 for OH over pure water and 28% sulfuric acid, respectively. Those for HOZare 20.01 over liquid water and 10.05 over sulfuric acid. We believe that the largest source of error in evaluating the measured loss rate coefficients is due to the uncertainty in the gas flow rate. This is because of the uncertainty in the partial pressure of water (*lo%) and the resultant error in the carrier gas volume flow rate. For low values of the uptake coefficient, Le., y < the overall uncertainty in y is approximately the same as the uncertainty in the first-order loss rate coefficient. When y > diffusion begins to limit transport in our experiments and the largest uncertainty in calculating the uptake coefficient comm from the estimation of the diffusion coefficients, which are probably accurate to f20%. Under certain situations, this 20% error propagates into 1 order of magnitude uncertainty in the calculated uptake coefficients,as will be demonstrated below, and therefore only a lower limit for y can be determined. For the measurements over liquid water, where we believe that the effects of nonlaminar flow were present, the systematic error could be substantial in the calculation of y from gas-phase diffusion considerations. As explained above, we believe we have minimized this effect by extrapolating to zero flow velocity. The rationale for this is that the Reynold's number is a measure of gas turbulence (and nonlaminar flow) and NRe is a linear function of the flow velocity. Since the extrapolation leads to lower values of y than those obtained from our raw data, this analysis is consistent with our interpretation of the reported values of y as lower limits to the true mass accommodation coefficient. The largest values of y that can be derived from our raw data with the plug flow model are y O H = 0.004 and yHOl = 0.0033. These correspond to the steepest portions of the decay plots shown in Figures 1 and 2. They are annotated in the footnotes of Table I. The y values in Table I were calculated from the k, values listed there. Discussion The experimental loss rate coefficient, k,, is a measure of the change in the average concentration of the species, n,, with exposure time to the surface. The uptake coefficients are calculated by equating the gas-kinetic equation for the number of molecules striking the wall per unit time per unit area with the comparable measured quantity. The former is given as ywnJ4 where w is
the mean radical speed, and n, is the gas-phase concentration of the reactant at the liquid surface. Motz and Wisezoand others21-22 have shown that the gas-kinetic equation must be modified in the presence of a wall with a large loss coefficient. Under our experimental conditions this modification arises only when the gas-phase diffusion process introduces a much larger correction factor. The measurements give us the loss rate coefficient k, defined by the equation dn,/dt = -kwn,. Multiplying this by the volume-to-surface ratio of the cylindrical reactor, r/2, gives the measured loss of molecules per unit time per unit area. This results in the equation (r/2)kWn, = ywn,/4. Under the plug-flow approximation, n, is taken equal to n, and y = 2 k , r / ~ . ~When ~ gas-phase transport becomes diffusion limited, there is a substantial radical concentration gradient, n, 0.1 leads to a k, that is within 10%of the diffusion limited loss rate coefficient, and this is consistent with our observed k, for OH that was near the diffusion limit. Using Z)OH.HO 7 &ELF.HZO = 94 Torr cmz s-I, the value for y for OH on sul!uric acid is 0.08. If we assume that the observed k, is at the diffusion limit, a value about 10% lower for Z)0H.H20 is implied. .ooolo ' ' I ' 10 20 30 40 50 (ii) Second-Order Loss in the Liquid. We have identified no k,(s-l) significant first-order losses for H 0 2 in the sulfuric acid solution or for OH in pure water. Therefore we consider loss in the liquid Figure 4. Uptake coefficient shown as a function of measured loss rate due to second-order recombination of the radicals. The self-recoefficient corrected for gas-phase diffusion limitation. (a, top) Two action of OH to make H202,a process which is second order in curves are shown for OH (solid lines) and H02 (dotted lines) at two representative pressures and temperatures: (1) 5 Torr of water vapor OH, is very fast in the kII = 5 X lo9 M-'s-I. Similarly, with 4 Torr of helium at 275 K and (2) 0.55 Torr of H20with 1 Torr the HOz loss process in low pH solutions is the self-reaction to of He at 249 K. The measured first-order loss rate coefficients for OH make H 2 0 2and Ozwith a second-order rate constant of 8 X lo5 and H02 are indicated by solid and dashed arrows, respectively. (b, M-'S - I ? ~ We investigated the degree of saturation of the surface bottom) An expanded portion of (a) for the OH at 275 K experimental under these conditions, using a simplified diffusion equation with conditions. In this part, the dashed line is a second diffusion corrected a second-order loss. Ignoring liquid transport along the direction curve where a 20%smaller diffusion coefficient was used. The uptake of the flow and using planar geometry for the liquid, the coefficient calculated from the plug flow approximation is also shown. steady-state equation is Dp(aZC(x)/dxz)= kII[C(x)]2 where C(x) saturation effects and our analysis procedures. is the concentration in the liquid, x is the distance from the (2) Saturationof the Liquid Surface and Aqueous-Phase Loss gas-liquid interface, and kIIis the mnd-order loss rate coefficient. Processes. A characteristic time for the saturation of a liquid This can be solved numerically with the boundary condition at surface, T,~,, by a gas-phase species with a Henry's law coefficient x = 0 of Dp(dC(O)/dx) = -(Js - J H J ,where J,, the flux into the H is isat, = Dc(4HRT/aw)2:4-26 where Dp is the liquid-phase liquid from the gas phase, is equal to a w 4 4 and the Henry's law cm2s-l), H i s in M atm-l, diffusion coefficient (we used Dp = flux out of the liquid, JHL,is equal to awC(O)/4HRT. The unique, and R is the gas constant (0.082 L atm mol-' K-l). This is the well-behaved solution gives a value for C(O), the concentration amount of time needed for a gas-phase species introduced over of the species in the liquid at the surface, and hence, J H L . The an initially pure, quiescent liquid surface to come within 60% of quantity ymax = [1 - JHL/J,], the maximum measurable uptake Henry's law saturation. Although this equation is not strictly coefficient (assuming a = l), is shown in Figure 5 as a function applicable to our situation, it provides a rough estimate for satof gas-phase number density at the gas-liquid interface for three uration effects. Assuming a = 1, rat,for OH is about lod s, and different Henry's law coefficients for OH at 275 K and an HHs for HOz it is about 0.1 s. The Henry's law coefficients used in = lo5 M atm-' for HOz at 249 K. The data for O H in Figure this equation are discussed in Appendix 11. The time for the liquid 5 can be duplicated very accurately using eq 1 with a = 1 and surface to traverse the measurement region in our experiment is k p = klIIOH], where [OH],, is the aqueous concentration of OH on the order of 10 s, indicating that, with no chemical loss in the that would be in equilibrium with the gas-phase OH, i.e., [OH],, liquid, we would measure only small loss rate coefficients. = HOHPOH. (i) PseudO-FiRt-Order Loas in the Liquid. For a first-order loss The calculations presented in Figure 5 show that for our op in the liquid, DanckwertsZSpresented an exact solution of the erating conditions the second-order loss for H 0 2 is sufficiently steady-stateconcentration of the species at a planar liquid surface. fast to allow uptake coefficients of >0.1 to be measured over After suitable substitution, eq 5.4 in ref 25 becomes sulfuric acid. This value is sufficient to give first-order loss rates close to the diffusion limit, which is what we found in the case 1 1 0 _(1) of H 0 2 on sulfuric acid. f f 4HRT[kcDc]1/2 This brings us to the special case of our OH uptake measurements over water. As discussed in Appendix 11,a low solubility where His the Henry's law coefficient, R is the gas constant (0.082 of OH in water, HoH 100 M am-', is derived from the available L atm mol-' K-I), kp is the first-order loss rate coefficient in the thermodynamic data. This limits the buildup of OH in solution and Dp is the liquid-phase diffusion coefficient. liquid phase (d), and hence the effectiveness of the second-order reaction to scavenge This equation has been duplicated by others using an approach based on characteristic times for diffusion and r e a c t i ~ n . ~ ~ - ~ ~OH. Thus surface saturation may affect our measurements. From
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4984 The Journal of Physical Chemistry, Vol. 96, No. 12, 1992
ficients for each species in helium and in water vapor:30
I
l/Dc =
[b2/al2)+
b3/al3)1pT
(All
where y j are the mole fractions, ajjare the binary diffusion coefficients (Torr cm2 s-'), and PT is the total pressure (Torr). The diffusion coefficient for OH in helium was taken to be that of 0 atoms in helium, which is 658 Torr cm2 s-' at 274.5 K and 536 Torr cm2 s-I at 249 K using a temperature dependence.I3v3' H 0 2 in helium was taken to have the same value as O2in helium, which is 485 and 409 Torr cm2s-' at 274.5 and 249 K, respectively." In principle, the pressure-independent binary diffusion coefficients for OH and HO2 in water vapor can be estimated from7 109
IO6
ns (
molecule ~
IO"
IOD
m
-
~
Figure 5. Maximum measureable uptake coefficient plotted as a function of gas-phase number density at the surface for a second-order liquidphase loss process. These were obtained from a steady-state numerical solution of the diffusion equation with a second-order loss process. The results depend upon the physical Henry's law coefficients, and three different curves are given for OH: 40, 100, and 250 M atm-' (dashed lines). The results for H 0 2 at 249 K are presented as a solid line.
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the analysis shown in Figure 5 , we predict OH uptake coefficients over pure water for our OH concenof no greater than trations. This is consistent with our observation of a first-order loss rate coefficient that was significantly below the gas-phase diffusion limit, shown in Figure 4b,and a value for y of -0.003. If we assume that the observed uptake was limited by the recombination of OH in solution and not by the mass accommodation coefficient, which seems reasonable in light of the lower limit to a we found for OH on 28% H2S04, a value for HoH I 100 M atm-' at 275 K is consistent with our measured loss of OH on water. (3) Comparisoo with Preriow Studies. Our results for OH on sulfuric acid support the assertion that stratospheric sulfuric acid aerosols will efficiently scavenge OH radicals as postulated by Gershenzon et al.' The low uptake coefficient for OH on sulfuric acid observed by Baldwin and Golden4may be explained by the fact that they attempted to remove water from the -96% sulfuric acid by pumping on it and thus it may have had a small [HSO;]. In addition, their experiment may not have been able to accurately measure high uptake coefficients. They report only lower limits of for the uptake coefficients of NH3 and H 2 0 on concentrated sulfuric acid, whereas these y are expected to be high. The lower limits to the mass accommodation coefficient of H 0 2 that we determined are consistent with those measured by Mozurkewich et a1.6 for H 0 2over concentrated solutions of NH4HS04 and LiN03.
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Conclusions Measured first-order wall loss rate coefficients for H 0 2 and OH on a sulfuric acid solution and for H 0 2 on water were found to be close to the diffusion limit. The lower limits to the mass accommodation coefficients of OH and H 0 2 on 28% w/w sulfuric acid are 0.08 and 0.05, respectively, and for H 0 2 on water at 275 K it is 0.01. Since our measured values are strongly influenced by the value of the diffusion coefficients, the mass accommodation coefficients could be as high as unity. The measured uptake coefficient for OH of 0.0035 on liquid water at 275 K was probably influenced by surface saturation; thus a for OH on pure water may also be very high. Because diffusion limits gas-phase transport in our experiment and in the atmosphere for similar values of a, results from our experiments indicate that a does not limit the mass transport into aqueous droplets in the troposphere. Acknowledgment. This work was supported by NOAA as a part of the National Acid Precipitation Assessment Program. Appendix I. Diffusion Coefficients
The diffusion coefficients for OH and H 0 2in the helium-water vapor mixture can be estimated from the binary diffusion coef-
ajj
= P[Dij] =
0.002628 (2Pjj)' P U 2 ( Q('.')*)
(atm cm2 s-I)
(A2)
where pv is the reduced mass (amu), Tis temperature,p is pressure (atm), u is the combined collision parameter (A), and (Q('J)*) is the reduced collision integral for diffusion calculated from dipole-dipole interaction^.^,^ Collision parameters and force constants have been estimated for H 2 0 from viscosity measurements of steam,' but none are available for OH or HOz. The dipole moments" (D) are as follows: OH, 1.6; H 0 2 , 2.1; H 2 0 , 1.85; H202,2.26. The selfdiffusion coefficient for H20, B s ~ from (A2) is 116 Torr cm2 s-I at 275 K, and it has a T 2 . I temperature dependence in this range. The binary diffusion coefficients for OH in water vapor, ~ O H . H ~ Owould , be expected to be close to 3 3 s a F . H ~because the d i p o l d p o l e interaction dominates and OH and H 2 0 have similar dipole moments and masses. Similarly, 3 3 H O r ~ , 0should be less than B~ELF.H~O because of a larger reduced mass for the H02-H20 interaction,and the collision parameter for H 0 2 is probably greater than that for OH. Monchick and Mason7point out that the transport properties of highly polar substances like water can show substantial discrepancies from their calculations. The collision integral for diffusion is particularly sensitive to the calculation procedure at low temperatures. Danon and A m d ~ calculated r~~ the collision integral for H20-H20 based on a different treatment of the dipoltdipole interaction. They obtained collision integrals for viscosity that were 10-2095 larger than those presented in ref 7. Therefore, the accuracy of the calculated binary diffusion coefficients involving molecules with a large dipole interaction is on the order of h2W0. The binary diffusion coefficient for H 2 0 2in H20, aH orHp was estimated from viscosity measurements of H202-H2bmixtures to be 231 Torr cm2 s-' at 443 K.34 It should be noted that this estimation procedure was based on calculationssimilar to those discussed in the previous paragraph. Using a P.07dependence,' this gives a value of 86 Torr an28' for DH*HP at 275 K. Thw two values, the self-diffusion coefficient for H 2 0 and the H202-H20 binary diffusion coefficient, serve as reference points for the OH-H20 and H02-H20 binary diffusion coefficients, respectively.
Appendix II. Henry's Law Coefficients The aqueous-phase loss of a radical due to its self-reaction is strongly coupled to the Henry's law coefficient for that gas. Also, from eq 1, the uptake coefficients for a pseudo-first-order loss depends on H. The Henry's law coefficients for H 0 2 and OH have not been measured. However, the free energies of solvation AGwlhave been estimated3sZ7for both species by evaluating the standard free energy changes in thermochemical cycles from known standard enthalpies and entropies. The value for OH, -2.0 f 0.4 kcal mol-',27 gives H = 30 M atm-' at 298 K using H (m atm-') = exp(-AGw1/RT)
('43)
At low concentrations,as is usually encountered in the atmosphere and laboratory, concentrations given in molar and molal are in. terchangeable for water at a density of 1 g ~ m - ~Schwartzj estimated AGd(H02) to be -4.2 kcal mol-', while Golden et ala2' calculated a value of -4.9 f 0.5 kcal mol-'. The Henry's law coefficient for H 0 2 we use was calculated using -4.9 kcal mol-'
~
.
~
J. Phys. Chem. 1992,96,4985-4990
for AGmI(H02)which yields HHQ= 4 X lo3M atm-' at 298 K. To obtain estimates of the Henry's law coefficients at lower temperatures, we used a value for ASmlof -23 f 3 cal K-' mol-'. This results in a change in the value of AGml from that at 298 K of -0.5 and -1.1 kcal mol-' at 275 and 249 K,respectively, assuming AHd does not change with temperature. Thus the values for HOH are 100 and 500 M atm-' and for HHol are 2 X 104 and 2 X lo5M atm-' at 275 and 249 K,respectively. The uncertainty in these estimates for the Henry's law coefficient is substantial: a f0.5 kcal mol-' uncertainty in AGml results in a factor of 2.5 uncertainty in H. Registry No. Hydroxyl radical, 3352-57-6; hydroperoxy radical, 3 170-83-0; sulfuric acid, 7664-93-9.
References and Notes (1) Utter, R. G.; Burkholder, J. B.; Howard, C. J.; Ravishankara, A. R. J. Phys. Chem., preceding paper in this issue. (2) Danckwerts, P. V. Gas-Liquid Reactions; McGraw-Hill: New York, 1970. (3) Schwartz, S.E. J. Geophys. Res. 1984,84, 11589. (4).Baldwin, A. C.; Golden, D. M. J . Geophys. Res. 1980, 85, 2888. Baldwin. A. C.: Golden. D. M. Science. 1979. 206. 562. (5) Gershenbn, Y. M.; Ivanov, A. V i Kucheryavyi, S.I.; Rozenshtein, V. B. Kinet. Karal. 1987, 27, 923. (6) Mozurkewich, M.; McMurry, P. H.; Gupta, A.; Calvert, J. G. J . Geophys. Res. 1987, 92, 4163. (7) Monchick, L.; Mason, E. A. J. Chem. Phys. 1961,35, 1676. (8) Mason, E. A.; Monchick, L. J. Chem. Phys. 1961, 36, 2746. (9) Emmert, R. E.; Pigford, R. L. Chem. Eng. Prog. 1954, 50, 87. (10) Schlichting, H. Boundary Luyer Theory; McGraw-Hill: New York, 1955. (11) DeMore, W. B.; Sanders, S. P.; Molina, M. J.; Golden, D. M.; Hampson, R.F.; Kurylo, M. J.; Howard, C. J.; Ravishankara, A. R. Chemical Kinetics and Photochemical Data for Use in Stratospheric Modeling Evaluation No. 8. NASA JPL Publication 90-1; Jet Propulsion Laboratory: Pasadena, CA, 1990. (12) Vaghjiani, G.L.; Ravishankara, A. R. J. Phys. Chem. 1990,93,7833. ~
4985
(13) Keyser, L. F. J . Phys. Chem. 1984,88, 4750. (14) Handbook of Chemistry and Physics, 60th ed.;Weast, R. C.,Ed.; CRC Press: Boca Raton, FL, 1979. (15) Gable, C. M.; Betz, H. F.; Maron, S.H. J . Am. Chem. SOC.1950, 72, 1445. (16) Perry, J. H. Chemical Engineer's Handbook; McGraw-Hill: New York, 1950. (17) Stirba, C.; Hurt, D. M. AIChE. J. 1955, 1, 178. (18) Lynn, S.;Straatemeier, J. R.; Kramers, H. Chem. Eng. Sci. 1955, 4, 49. (19) Brown, R. L. J . Res. Narl. Bur. Stand. (US.)1978, 83, 1. (20) Motz, H.; Wise, H. J . Chem. Phys. 1960, 32, 1893. (21) Murphy, D. M.; Fahey, D. W. Anal. Chem. 1987.59, 2753. (22) Schell, M.;Kapal, R. J . Chem. Phys. 1981, 75, 915. (23) Howard, C. J. J. Phys. Chem. 1979, 83, 3. (24) Schwartz, S . E. In Chemistry of Multiphase Atmospheric Systems; Jaeschke, W., Ed.; NATO AS1 Series; Springer-Verlag: Berlin, 1986; Vol. G6, p 415. (25) Danckwerts, P. V. Trans. Faraday SOC.1951, 47, 1014. (26) Worsnop, D. R.; Zahniser, M. S.;Kolb, C. E.; Gardner, J. A.; Watson, L. R.; Van Doren, J. M.; Jayne, J. T.; Davidovits, P. J. Phys. Chem. 1989, 93, 1159. (27) Golden, D. M.; Bierbaum, V. M.; Howard, C. J. J. Phys. Chem. 1990, 94, 5413. (28) Buxton, G.V.; Greenstock, C. V.; Helman, W. P.; Ross,A. B. J. Phys. Chem. Ref. Data 1988,17, 513. (29) Farhatziz; Ross, A. B. Selected specific rate of reactions of transients from water in aqueous solutions. III. 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. (30) Fairbanks, D. F.; Wilke, C. R. Ind. Eng. Chem. 1950,42,471. (31) Marrero, T. R.; Mason, E. A. J . Phys. Chem. Ref. Data 1972, 1, 3. (32) The dipole moments for H202 and H 2 0 from: Townes, C. H.; Schawlow, A. L. Microwave Spectroscopy; Dover: New York, 1975; p 639. H 0 2 from: Saito, S.;Matsumura, C. J . Mol. Spectrosc. 1980,80, 34. OH from: Peterson, K. I.; Fraser, G.T.; Klemperer, W. Can. J . Phys. 1984, 62, 1502. (33) Danon, F.; Amdur, I. J . Chem. Phys. 1%9,50,4718. (34) Weissman, S. Proceedings of rhe Fourth Symposium on Thermophysical Properties; Moszynski, J. R., Ed.; ASME: New York, 1969; p 360.
Silicaiite Characterization. 1. Structure, Adsorptive Capacity, and I R Spectroscopy of the Framework and Hydroxyl Modes A. Zecchina,* S. Bordiga, G. Spoto, L. Marchese, Dipartimento di Chimica Inorganica, Chimica Fisica e Chimica dei Materiali, Universitii degli Studi di Torino, via P. Giuria 7, IO125 Turin, Italy
G. Petrini, G. Leofanti, and M. Padovan ENICHEM ANIC Centro Ricerche di Bollate, via S. Pietro 50. 20021 Bollate, Italy (Received: August 26, 1991; In Final Form: January 22, 1992)
T h e physical properties (microcrystal morphology, crystallinity, internal perfection, adsorptive capacities toward CO and N2, and IR manifestations of the skeletal modes and of the hydroxyl groups) of a Na- and AI-free silicalite (S) prepared following a specifically designed method are investigated. Comparison is made with the silicalite containing Na and Al impurities prepared following a more conventional (classical) path (SNa). T h e experimental techniques are X-ray diffraction (XRD), high-resolution electron microscopy (HRTEM), infrared spectroscopy in reflectance mode (FTIR), and volumetric isotherms in a BET apparatus. The most relevant results are as follows: (i) S microcrystals have very regular octagonal prismatic habit, (ii) nanodefects and microcavita are more aboundant in S than in SNa, (iii) hydroxyl groups are present at the internal defects sites, and (iv) N2 and CO volumetric isotherms at 77 K differ from S and SNa.
htroduction Silicalite, a zeolite having a MFI structure,' is often considered a purely siliceous zeolite without Al heteroatoms in the framework. However, this picture is not completely true: in fact, to our best knowledge, silicalite synthesized according to the U.S.Patent of Grose and Flanigen2contains, as impurities, variable concentrations of A1 and Na. With the exception of refs 2 and 3, all the researchers reporting about silicalite do not give full details about t h e AI and Na impurity concentrations, probably because they are considered as not influencing the properties of the zeolites.
To our experience, A1 and Na impurities cannot be completely removed by treatments of the solid at the end of the synthesis, either before or after the calcination. We think that the presence of these impurities can affect in different manners the physical properties of the samplen4 We have therefore synthesized a high-purity sample, by changing the synthesis procedure (vide infra), and the physical properties of this high-purity silicalitewere then compared with those of the analogous SNa samples used as standard reference. Another important reason for studying the Na-free silicalite (S) is that only the full understanding of its
0022-3654/92/2096-4985$03.00/00 1992 A m e r i c a n C h e m i c a l Society
