Environ. Scl. Technol. 1983, 17, 145-149
(9) Dressen, D. R.; Gladney, E. S.; Owens, J. W.; Perkins, B. L.; Wienke, C. L.; Wanger, L. E. Environ. Sci. Technol. 1977,11, 1017-1019. (10) Dudas, M. J. Environ. Sci. Technol. 1981, 15, 840-843. (11) Shannon, D. G.; Fine, L. 0. Environ. Sci. Technol. 1974, 8, 1026-1028. (12) Kanapilly, G. M. Health Phys. 1977, 32, 89-100. (13) Kanapilly, G. M.; Raabe, 0. G.; Goh, C. H. T.; Chimenti, R. A. Health Phys. 1973,24, 497-507. (14) Limbek, B. E.; Rowe, C. J.; Wilkinson, J.; Routh, M. W. Am. Lab. 1978,10,89-100. (15) Fisher, G. L.; Raabe, 0. G.; Prentice, B. A,; Silberman, D.,
Registry No. ",OH, 1336-21-6; HC1, 7647-01-0; EDTA, 60-00-4;TI%,77-86-1; Zn, 7440-66-6;Mn, 7439-96-5;Cr, 7440-47-3; Ni, 7440-02-0; Cu, 7440-50-8; V, 7440-62-2; Al, 7429-90-5; Fe, 7439-89-6; Ca, 7440-70-2; K, 7440-09-7; citric acid, 77-92-9; histidine, 71-00-1; glycine, 56-40-6.
Literature Cited Coles, D. G.; Ragaini, R. C.; Ondov, J. M.; Fisher, G. L.; Silberman,D.; Prentice, B. A. Environ. Sci. Technol. 1979, 13,455-459.
Davison, R. L.; Natusch, D. F. S.; Wallace, J. R.; Evans, C. A. Environ. Sci. Technol. 1974,8, 1107-1113. Campbell, J. A,; Laul, J. C.; Nielson, K. K.; Smith, R. D. Anal. Chem. 1978,50, 1032-1040. Hansen, L. D.; Fisher, G. L. Environ. Sci. Technol. 1980, 14,1111-1117. Linton, R. W.; Loh, A,; Natusch, D. F. S.; Evans, C. A.; Williams, P. Science (Washington,D.C.) 1976,191,852-854. Campbell, J. A.; Smith, R. D.; Davis, L. E. Appl. Spectrosc. 1978,32, 316-319. James, W. D.; Janghorbani, M.; Baxter, T. Anal. Chem. 1977,49, 1994-1997. Theis, T. L.; Wirth, J. L. Environ. Sci. Technol. 1977,11, 1096-1110.
Annual Report, Laboratory for Energy Related Health Research, UCD 472-125, 1978; pp 26-32. (16) Hansen, L. D.; Silberman, D.; Fisher, G. L. Environ. Sci.
Technol. 1981, 15, 1057-1062. (17) Eggett, J. M.; Thorpe, T. M. J . Environ. Sci. Health 1978, A13, 295-313.
Received for review June 23,1982. Accepted November 15,1982. This work was supported by the U.S. Department of Energy under Contract DE-AM03-76SF00472.
Solubility of Ozone in Aqueous Solutions of 0-0.6 M Ionic Strength at 5-30 OC Lynn F. Kosak-Channlng and Oeorge R. Helz"
Department of Chemistry, University of Maryland, College Park, Maryland 20742
rn The equilibrium constant of the reaction 03(1) + 03(g)
has been measured between 5 and 30 "C by analyzing equilibrated liquid and gas phases and has also been bracketed on both the high and low sides by interpreting the dynamics of transfer of ozone between the liquid and gas phases. Mildly acidic solutions were used in order to minimize ozone decay. The effect of ionic strength was determined by adding sodium sulfate to the solution. The following equation predicts all the measured values of Kh with a root mean square deviation in KhoM/Kh*d of 2.3% (2' = temperature in Kelvin; I.L = molar ionic strength): ln Kh = -2297T1 + 2 . 6 5 9 ~- 688.OpT' 12.19. Experience gained during this work with dynamic equilibrium methods of measuring Kh suggests that proof of equilibrium may be more difficult to attain than previously suggested. Particularly a t the upper end of the temperature range studied, the ionic strength effect on the value of Kh is large enough to be significant in natural masstransfer processes such as absorption of tropospheric ozone into seawater or into aqueous aerosols.
+
Introduction Knowledge of the solubility of ozone in water and salt solutions is needed to design efficient ozonation systems for water treatment (I) and to predict rates of ozone removal from the troposphere into the oceans (2) or into aqueous aerosols (3). Published measurements of the Henry's law constant for ozone in pure water disagree by as much as a factor of 2 (4-12).Many of the older values were obtained in neutral or alkaline solutions and may be questionable because ozone decomposes at appreciable rates under these conditions (9). Nevertheless, the nineteenth century measurements of Mailfert (lo),which are systematically lower than most subsequent ones, have found their way into standard reference books (11)and therefore continue to be widely used. A particular defi0013-936X/83/0917-0145$01.50/0
ciency of the data in the literature is their inability to predict the effect of dissolved salts on ozone solubility. This effect can be important in the case of ozone dissolution into the ocean or into sulfate-laden aerosols. In this paper, we report values of ozone's Henry's law constant, measured in a continuously flowing bubble column after the liquid and gas phases had come to equilibrium. Dissolved ozone was stabilized by conducting the measurements in mildly acidic solutions (pH 3.4 f 0.1). We also attempted to confirm these data by an independent, dynamic equilibrium method similar to that described by Mackay et al. (13). However, because mass transfer was too slow to maintain equilibrium when ozone was being added to or removed from the column, the dynamic equilibrium measurements serve only to bracket the true value of the Henry's law constant.
Experimental Section A diagram of the apparatus is shown in Figure 1. Ozone gas from a Welsbach ozone generator (Model T-408) in a stream of dried O2 (A) was bubbled through a fine, cylindrical frit at a typical flow rate of 60 mL/min into a thermostated column containing 500 mL of an aqueous solution. The solution circulated at 23 mL/min by a peristaltic pump (B) through a flow cell for measuring ozone's optical absorbance at 260 nm (S). The ozone gas was introduced until a steady-state O3 concentration in the liquid was observed on the recorder (C). The exit gas above the solution was partly exhausted (D) and partly recirculated at approximately 250 mL/min by a bellows pump (E). Recirculating the exit gas was intended to facilitate the system's attaining equilibrium by increasing the contact of the gas and liquid. When a steady-state ozone concentration was observed in the liquid, the exhaust gas was diverted for 2 min to a trap (F) containing 500 mL of 2% KI. The flow of O3to
0 1983 American Chemical Society
Environ. Sci. Technol., Vol.
17, No. 3, 1983
145
I
Steadv ._ _ _
Y
state'
-Decay
/
I
I
,
1
0
20
40
60
80
__
TIME
(MIN)
Flgure 2. Characteristicappearance of data for determination of K,; temperature of this trial, 25.0 OC.
stant flow rate. Again, absorbance measurements were recorded with time. A typical record of ozone in the liquid phase during the entire sequence of steps is given in Figure c
2.
The procedure was applied to distilled water, 0.05,0.1, and 0.2 M NazS04solutions each at 5, 1 0 , 1 5 , 2 0 , 2 5 ,and 30 f 0.1 O C . The pH of each solution was adjusted to 3.3-3.5 with H2S04to inhibit ozone decomposition.
Theoretical Background Flgure 1. Apparatus for Henry's law constant determination: ozone generator, A; peristallc pump for liquid circulation, 6;strip chart recorder, C; exhaust, D; bellows pump for gas circulation, E potassium iodide trap, F; gas collection (water displacement)apparatus, W; oxygen supply, H; spectrophotometer, S; thermometer, T. Height of column, 62 cm; height of jacketed portion of column, 47 cm; diameter of Inner column, 4.3 cm.
the column was then shut off. The gas that exited the KI trap during the 2-min period was collected (W), and its volume was measured by water displacement. Subsequently, the quantity of ozone in the exhaust gas was determined by subjecting 25.0 mL of the trap solution to amperometric titration with phenylarsine oxide at pH 2. It was found that some IO; as well as I; was formed when ozone reacted with 2% KI solution, but adjustment of the trap solution to pH 2 allowed the sum of both products to be determined (14). From this titration, the concentration of ozone in the exhaust gas was calculated. The corresponding concentration of ozone in the liquid was evaluated from the optical absorbance measurement of the solution. A molar absorptivity of 2930 M-l cm-l (7) was used. This value is in agreement with the value measured by Hoigne and Bader (15) and was confirmed during our work by analysis of the mildly acidic test solutions. It is assumed to be independent of temperature, in the range 0-30 "C,on the basis of studies of gaseous ozone, which has a very similar molar absorptivity and absorption maximum (16 ) . After the flow of ozonated gas was stopped, a slow decrease in the ozone concentration in the liquid phase was observed while the gas and liquid continued to circulate within the system. No decrease in ozone absorbance in the liquid phase was observed in an approximately 10-min period, however, if the gas recycling pump (E) was stopped while the liquid continued to circulate. Therefore, the observed decrease was attributed to ozone loss in the gas phase. Catalytic decomposition of O3in glass frits has been reported (17-19) and is probably the source of the observed loss. To determine the rate of loss, we recorded absorbance measurements for at least 10 min after stopping the flow of 03-O2 gas. Subsequent to measuring the rate of loss, O2 gas (H in Figure 1)was bubbled into the column at a rate of 25 f 0.25 mL/min to strip ozone out of solution. A flow controller was used ahead of the column to maintain a con146
Environ. Sci. Technoi., Vol. 17, No. 3, 1983
The ratio of the measured concentrations of ozone in the gas phase ([03(g)])and theliquid phase ([03(1)])defines a dimensionless partition coefficient:
z = C03(g)l/[O3(l)l
(1)
When the system reaches a steady state, the observed partition coefficient will assume ita equilibriumvalue (Zq). The Henry's Law constant, Kh,can be computed from :2 , (2) Kh = Po,/Mo, = RTZ,, where Po, is the partial pressure of ozone in the gas phase, Mo, is the molar concentration of ozone in the liquid phase, T is the absolute temperature, and R is the gas constant (0.0821 L atm K-l mol-l). The value of Z at steady state must be an equilibrium partition coefficient because Figure 2 demonstrates that exchange of ozone between the phases is readily reversible on the time scale of the experiment. The only ambiguity in obtaining Kh by eq 1 and 2 would arise if chemical decay of ozone in one of the phases occurred too rapidly for mass transfer between the phases to keep up. Although slow decay was observed in the gas phase in these experiments, it will be shown that ozone was being replaced at such a rate relative to decay that [03(g)]could not have been depleted relative to the inlet gas by more than a few percent. The accuracy of Kh obtained by means of eq 1 and 2 depends on the accuracy of the molar absorptivity value used to compute [O,(l)] from the absorbance measurements and the accuracy of analyses of the gas. In principle, it is possible to obtain Khindependently of these quantities by studying the rate at which the liquid composition responds when exposed to a gas with which it is not in equilibrium. Mackay et al. (13)discussed this method for the case in which the inlet gas contained none of the solute of interest and in which no chemical decay occurred within the system. For a more general case, the total ozone in the system (EO,) and the rate at which it changes with time can be expressed by (3) c 0 3 = vg[o3(g)l ivl[o3(l)l and
..
Here, V, and V, are the volumes of the gas and liquid phases, G is the volumetric flow rate of the inlet gas (measured at temperature To),T i s the temperature of the system, [03(1)]is the ozone concentration of the inlet gas (which may be zero), and kD is the chemical decay constant for ozone in the gas phase. If the gas flow rate is calibrated at the temperature of the system, then T/To drops out of eq 4. Equations 1, 3, and 4 may be solved to eliminate the variables EO3and [03(g)],yielding d[0,(1) 1 -
0
60 0
$1
--
dt
(5) When G = 0, kD may be evaluated in terms of 2 from measured values of d[O,(l)]/dt by using this equation. At steady state, when d[03(l)]/dt = 0: (6) [03(g)l/[03(I)1 = 1/(1 + VgkDTo/GT) This equation indicates the extent to which chemical decay causes [03(g)]to differ from [03(1)]at steady state. For the present experiment, [03(g)]/[03(I)]was always between 0.94 and 1.00. Integration of eq 5 yields
Here [O3(l)lO is the initial ozone concentration in the liquid phase, [O3(l)Imis the final, steady-state concentration, and [O,(l)lt is the concentration at some point in time after the liquid was first exposed to a gas phase with which it was not in equilibrium. The partition coefficient, 2, can be evaluated from eq 7, which applies equally well whether O3 is being added to or removed from the liquid phase. The various ozone concentrations in eq 7 may be replaced with the corresponding optical absorbances, so the accuracy of Z, calculated in this manner, is not dependent on the accuracy of the molar absorptivity value for ozone. Derivation of eq 7 assumes that the gas and liquid phases in a completely mixed reactor are effectively homogeneous; otherwise, the definition of 2 (eq 1) is ambiguous. Also 2 is assumed to have a constant but not necessarily an equilibrium value. Mackay et al. (13) suggest that it is possible to determine if the value of 2, obtained from an equation similar to 7, is an equilibrium value by varying the height of liquid in the bubble column. This changes the gas-liquid interfacial area because the deeper the liquid, the greater the number of bubbles at any instant. If the value of 2 does not change with interfacial area, then it is argued that the gas leaving the liquid column must have reached equilibrium with the liquid. In the early phases of this work some experiments were done with a purge-without-recycleapparatus very similar to that of Mackay et al. (13). Variation of the liquid volume produced constant 2 values when the liquid volume exceeded 450 mL (Figure 3). However subsequent work showed that the Kh values computed from 2 at V, > 450 mL could not be confirmed from the reverse direction, by measuring O3uptake rates. We have concluded that variation of liquid volume provides unreliable evidence that an observed value of 2 is an equilibrium value. The gas-liquid contact area does not vary in a simple way with V,, because bubbles of gas fuse as they rise. If mass transfer is not fast enough to maintain equilibrium during a dynamic experiment, then the liquid con-
x
x
20
0
x
X 0
0
0
200 400 Volume ( m i )
600
Flgure 3. Test for equilibrium: system at 30 O C , 0 ; 15 O C , X; 1 O C , 0.The expected values of Kh (from eq 8) at 30 and 1 O C were 100 and 45 atm M-', respectively.
Table I. Unitless Partition Coefficients, 2,from Uptake, Steady-State, and Purge Methods" at Experimental Temperatures and Ionic Strengthsb
z
T,O C
p, M uptake steady state purge 5.0 0.00 2.33 2.20 1.85 10.0 2.74 2.71 2.07 15.0 3.40 2.86 2.51 20.0 6.48 3.30 2.72 25.0 7.63 3.68 2.86 30.0 14.6 4.15 3.30 5.0 0.15 2.83 2.33 1.95 10.0 3.90 2.70 2.55 15.0 3.97 3.02 2,79 20.0 4.68 3.28 3.11 25.0 3.40 3.66 2.93 30.0 47.2 4.35 3.53 5.0 0.30 2.40 2.12 1.32 10.0 3.40 2.66 2.39 15.0 2.77 2.84 2.66 20.0 6.82 3.50 3.18 25.0 5.05 4.01 3.37 30.0 4.52 4.39 4.70 5.0 0.60 3.63 2.51 2.43 10.0 5.54 2.83 2.78 15.0 5.20 3.50 3.24 20.0 6.57 3.90 3.23 25.0 5.45 4.54 3.54 30.0 13.0 5.15 3.89 a Z = (molar concentration of 0, in the gas hase)/(moExperilar concentration of 0, in the liquid phase). ments were performed with distilled water, 0.050 M, 0.099 M, and 0.20 M Na,SO, solutions adjusted to pH 3.3-3.5 with H,SO,.
centration of a solute will lag behind the equilibrium value during uptake, whereas the exit gas concentration will lag behind the liquid during purging. This will cause the observed values of 2 to be larger than the equilibrium value during uptake and smaller than the equilibrium value during purging. The observed values obtained in each direction will bracket the equilibrium value. If the values obtained in each direction agree, then equilibrium is proven. Attainment of equilibrium in dynamic experiments is much more difficult for highly volatile solutes such as ozone than for relatively low volatility solutes such as those of interest to Mackay et al. (13). Results and Discussion The unitless partition coefficients, 2, calculated from Environ. Sci. Technot., Vol. 17, No. 3, 1983
147
Table 11. Decay Constants, k D , at Experimental Temperatures and Ionic Strengths 104kD, s-l T,O C 0.00 M 0.15M 0.30 M 0.60M 5.0 1.98 1.09 1.82 1.30 10.0 1.11 1.38 1.54 1.43 15.0 1.54 1.44 1.75 1.40 20.0 2.03 1.64 1.57 1.82 25.0 2.25 1.66 1.88 1.82 30.0 2.50 2.42 2.30 2.20 4.9-
Table 111. Smoothed Values of Henry's Law Constants, Kh,at Experimental Temperatures and Ionic Strengths4 Kh, atm M-I T,O C 0.00 M 0.15 M 0.30 M 0.60 M 56.3 51.8 53.2 5.0 50.5 59.3 61.3 63.4 67.9 10.0 73.8 80.2 68.0 70.8 15.0 93.9 81.8 85.8 20.0 78.0 97.0 108 25.0 87.5 92.2 113 127 106 30.0 100 These Henry's law constants were computed from regression eq 8 for the experimental temperatures and ionic strengths. Table IV. Bunsen Absorption Coefficients, a, at Experimental Temperatures4 T,OC a ,atm-' T,OC a ,atm-' 20.0 0.287 0.444 5.0 25.0 0.256 10.0 0.378 30.0 0.224 0.330 15.0 a = 22.414/Kh, where Kh is the Henry's law constant in atm M-I.
4.7-
4.5-
r
y.
c 4.3_I
4.1
pirical equation below by multivariable linear regression analysis:
-
In Kh = Up'
3 91
8
00
0 2 04 Ionic Strength ( M I
06
Flgure 4. Comparison of measured and predicted values of K,. Symbols represent raw experimental values. Solid lines represent smoothed values (Table III), computed from eq 8.
the uptake, steady-state, and purge experiments are presented in Table I. The disagreement between uptake and purge results indicate that the gas was not in equilibrium with the liquid phase while ozone was being added to, or removed from, the system. In principle, equilibrium would be more closely approached during the dynamic experiments if the recycle rate were increased or if the inlet flow rate (G) were decreased. However the recycle rate used here (250 mL/min) was the fastest practical without creating frothing in the column and loss of liquid. Also if the inlet rate were lowered much below the values we used, d[03(1)]/dt would become undesirably sensitive to chemical decay. The steady-state ratios in Table I are deemed to be the best estimates of the equilibrium values and the most precise values as judged by error propagation. For example, at 25 "C and p = 0, Z = 7.63 f 3.69 by the uptake method whereas Z = 3.68 f 0.22 by the steady-state method. Similarly, at 10 "C and p = 0.60, 2 = 2.78 f 0.41 by the purge method and 2.83 f 0.15 by the steady-state method. Although the uptake and purge data do not represent true equilibrium constants, they are useful in bracketing the true values, particularly at the lower temperatures. Values for the decay constant, kD,are listed in Table 11. The values of Z used in the computation of the decay constants were those obtained from the steady-state measurements. The data indicate that ozone loss in the gas phase followed first-order kinetics. The Henry's law constants (Kh), experimentally determined by the steady-state method, were fitted to the em148
Environ. Sci. Technol., Vol. 17,No. 3, 1983
bp
+ Cpp' + d
(8)
where T is the absolute temperature and p is the ionic strength. The regression coefficients a-d with their respective standard errors for 21 cases are -2297 f 88,2.659 f 0.852, -688.0 f 247.7, and 12.19 f 0.30, respectively. As graphically illustrated in Figure 4, a good fit of the measured Henry's law constants to the regression equation was obtained. The values of In Kh lie within f2.3% of the predicted values. Allowing for curvature in the In Kh - p relationship by inclusion of a p z T 1term in the regression equation provided negligible statistical improvementin the fit of the data over the ionic strength range investigated. Smoothed values of the Henry's law constants (Kh) at the experimental temperatures are listed in Table 111. Bunsen absorption coefficients (a),which are directly proportional to the concentration of O3 in the liquid, are provided in Table IV. The standard enthalpy and entropy changes for the reaction p = 0.00 03(1) 03(g) were evaluated from 8 according to d(ln Kh)/d(T') = a + C p = -AHo/R (9)
-
and bp
+ d = ASo/R
(10)
where R is the gas constant. The values of AHo and ASo are 19.1 f0.7 kJ mol-' and 0.101 f 0.002 kJ mol-' K-l, respectively. At 25 "C and p = 0.00 the standard freeenergy change (AGO) for the reaction is -11.1 f 0.3 kJ mol-'. For comparison, the temperature dependences of In Kh found in several other studies are compared to the present study in Figure 5. Brackets established by the uptake and purge experiments for the value of the Henry's law constant at 5 "C are shown in Figure 5. At higher temperatures, the brackets are less useful because their range is so large; however, at low temperature, some of the literature values clearly fall outside the brackets. The variation of Kh with temperature found by several previous workers
5201
Table V. Saltingout Coefficients, a , at Experimental Temperatures T, O C u, M-' T, O C u, M-' 20.0 0.311 0.184 5.0 25.0 0.350 0.228 10.0 30.0 0.388 15.0 0.270
5 00
4 801
+
4.401
A
I
. A
x c
Jz
t
4.20
*
+
d(ln Kh)/dp = b
O
A
Qc
4.001
3.80
3,601 340
3 20
+ c T 1 = u (constant )'2
(11)
0.0
L1
I
I
1
1
I
3 30
3 40
3 50
3 60
3 70
i o 3x
I/T
Comparison of In K h vs. T-' from previous studies to the present study: +, Rawson (6); (acid solution)and 0,Mailfert (70); A,Kawamura ( 4 ) ;0, Kilpatrick et al. ( 7 ) ;0, Matrazov et al. ( 8 ) ;X, Briner and Perrottet (5);0 , Rothmund from Rawson ( 6 ) ;V (p = O.l), Gurol and Singer (12);V (calculated at pH 3 4 ,Roth and Sullivan (9); A, Stumm from Roth and Sullivan ( 9 ) ; 0 , present study. Figure 5.
(e.g., Rawson (6) and Mailfert (10)) implies a AH value that is about twice the value found here and that is anomalously large compared to values for same other triatomic gases. For example, AHfo(g)- AHfo(l)for HzS, COz, CIOz and NzO are 19, 20.3, 27, and 26 kJ mol-l, respectively (11). Our value is 19.1 kJ mol-l, whereas that computed from Rawson's data is 37 kJ mol-l. The AH value obtained from Mailfert's data can be made to agree with ours if only his five lowest temperature points are used; these would be the points least affected by decomposition. Previous workers who tried to measure Kh in pure water, where ozone's chemical decay is much faster than at low pH, have had difficulties because of gas-phase inhomogeneity. For example, Matrazov et al. (8) measured O3in the inlet and outlet gases and assumed arbitrarily that the effective gas-phase O3concentration in the column was halfway in between the inlet and outlet gas concentrations. However, the values obtained by Matrazov et al. (8) at higher temperatures in acid solution, where gas-phase inhomogeneity is minimal, are consistent with the trend established at lower temperatures in the present study. The results of Stumm (9)are in good agreement with those of the present study although they produce a slightly lower AH. Excluding the results of Rawson (6) and Gurol and Singer (12)in 0.1 M ionic strength solution and Mailfert (10) in distilled water, the values of Kh and AH determined by several workers are consistent within a relatively narrow range. The dynamic measurements of Roth and Sullivan (9) in acid solution are in excellent agreement with ours. However in neutral and mildly alkaline solutions they found systematically higher Kh values. This pH dependence in their Kh values probably derives from incomplete correction in their computations for the rapid decomposition that occurs in alkaline solutions. The true Henry's law constant of a weakly solvated, nonhydrolyzing molecule such as ozone should be independent of pH at constant temperature, provided the pH adjustments do not produce changes in ionic strength. The dependence of the Henry's law constant on ionic strength, represented by the salting-out coefficient u, was assessed from eq 8. Values of the salting-out coefficients
at the experimental temperatures, listed in Table V, are similar to values for other substances of similar molecular weight (21). The magnitude of the ionic strength effect is large enough to have a significant influence on mass transfer rate of ozone in natural processes. For example, at 25 "C and an ionic strength of 0.7, which is representative of seawater, the rate of absorption of ozone into the ocean from the troposphere would be nearly 30% slower than in fresh water of the same temperature (see discussion in ref 2). Acknowledgments The paper benefitted considerably from the comments of reviewers Paul Chrostowski and Bruce Thomson. Registry No. Os,10028-15-6.
Literature Cited (1) Shambaugh,R. L.; Melnyk, P. B. In "Proceedings-Forum on Ozone Disinfection"; Fochtman, E. G., Rice, R. G., Browning, M. E., Eds.; Ozone Press International: New York, 1976; pp 312-326. (2) Garland,J. A.; Elzerman, A. W.; Penkett, S. A. J . Geophys. Res. 1980,85, 7488. (3) Larson, T.; Horike, N. R.; Harrison, H. Atmos. Environ. 1978, 12, 1597. (4) Kawamura, F. J . Chem. SOC.Jpn. 1932,53, 783. (5) Briner, E.; Perrottet, E. Helv. Chim. Acta 1939, 22, 397. (6) Rawson, A. E. Water Water Eng. 1953,57, 102. (7) Kilpatrick, M. L.; Herrick, C. C.; Kilpatrick, M. J . Am. Chem. SOC.1956, 78, 1784. (8) Matrazov, V. I.; Kashtanov, S. A.; Stepanov, A. M.; Tregubov, B. A. J . Appl. Chem. USSR (Engl. Transl.) 1975, 48, 1902. (9) Roth, J. A.; Sullivan,D. E. Ind. Eng. Chem. Fundam. 1981, 20, 137. (10) Mailfert, M. C. R. Hebd. Seances Acad. Sci. 1894,119,951. (11) Perry, J. H., Ed. "Chemical Engineers' Handbook", 4th ed.; McGraw-Hill: New York, 1963; p 14. (12) Gurol, M. D.; Singer, P. C. Environ. Sci. Technol. 1982,16, 377. (13) Mackay, D.; Shiu, W. Y.; Sutherland, R. P. Environ. Sci. Technol. 1979, 13, 333. (14) Kosak-Channing, L. Ph.D. Thesis, University of Maryland, College Park, MD, 1981. (15) Hoigne, J.; Bader, H. Water Res. 1976, 10, 377. (16) Vigroux, E. Ann. Phys. 1953,8, 709. (17) Naimie, H. Ozone: Sci. Eng. 1981, 3, 139. (18) Flamm, D. L.; Anderson,S. A. Environ. Sci. Technol. 1975, 9, 660. (19) Perrich, J.; McCammon, J.; Cronholm,L.; Fleishman, M.; Pavoni, J.; Riesser, V. AICHE Symp. Ser. No. 166 1977, 73, 225. (20) Wagman, D. D.; Evans, W. H.; Parker, V. B.; Halow, I.; Bailey, S. M.; Schumm, R. H. National Bureau of Standards, Washington, D.C., Tech. Note 270-3, 1968. (21) Randall, M.; Railey, C. F. Chem. Rev. 1927, 4 , 273. Received for review February 10, 1982. Revised manuscript received August 16,1982. Accepted November 3,1982. This work was supported by a grant from the Maryland Power Plant Siting Program.
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