Kinetics of Ozone Decomposition: A Dynamic Approach Mlrat D. Gurol"
Department of Civil Engineering, Drexel University, Phlladelphia, Pennsylvania 19104 Phlllp C. Slnger
Department of Environmental Sciences and Engineering, University of North Carolina, Chapel Hill, North Carolina 275 14
rn Decomposition kinetics of ozone in aqueous solution were studied at 20 "C and in the pH range 2-10, under both quiescent and dynamic conditions. The data obtained in a batch reactor under quiescent conditions were analyzed by integral, initial-rate, and differential methods. The results indicated that ozone decomposes by a second-order reaction with respect to ozone concentration. The second-order decomposition kinetics determined in the batch reactor was confirmed by using a mathematical model developed for ozone decompgsition in an ozone contact column and the experimental observations under such dynamic conditions. The rate of decomposition was found to be relatively insensitive to pH below 4 and relatively slow under such acidic conditions (kd= 0.27 L/(mol 8)). Above pH 4, the rate is pH dependent with kd = ko[OH-]o*55, where ko is specific to the chemical composition of the aqueous system.
Introduction Ozone is considered to be an attractive alternative to chlorine for water and wastewater disinfection and for the oxidation of various organic and inorganic contaminants. The rate of disinfection of microorganisms by ozone is proportional to the concentration of ozone in solution. Since ozone is unstable in aqueous solution, its effectiveness as a disinfectant depends upon the rate at which it decomposes. With respect to its behavior as an oxidant, ozone can react via a direct reaction pathway involving molecular ozone or by an indirect pathway involving various highly reactive radicals that arise from the decomposition. Therefore, to be able to design an efficient ozonation system, it is important to determine the kinetics of decomposition of ozone and the variables affecting the rate of decomposition. The decomposition of ozone in aqueous solution has been studied for several decades. In a review of the literature on the kinetics of ozone decomposition, the most common observation is the reported disagreement among different researchers as to both the order of the decomposition reaction and the magnitude of the reaction rate constant. The range of conditions investigated by various researchers and their conclusions concerning the reaction order relative to ozone were summarized by Peleg (1). This information is updated in Table I. Furthermore, although it is generally agreed that the reaction is catalyzed by the 0013-936X/82/09 16-0377$01.25/0
Table I. Summary of the Kinetics of Ozone Decomposition Decomposition in Water ref PH temp, "C reaction order
2 3 4
5 6 7 7 8 9 10 11 11 11 12 13
14 15 16 17
2 -4 5.3-8 acidic basic 1-2.8 7.6-10.4 0-6.8 8-10 5.4-8.5 10-13 9.6-1 1.9 6 8 2- 4 0.22-1.9 9 8.5-13.5 0.5-10.0 2.1-10.2 acidic basic
2 2
0 0
,12
1 1 1
0-27 1.2-19.8 25 25 5-25 25 25 10-50 10-20 30-60 5-40 20 18-27 3.5-60 25 25 25
)12
2 ,12
1 1 1 2 1 or 2 1 1 1
31z
1-2 1
Table 11. Rate Equations for Decomposition of Ozone rate eq, r = -d[O,]/dt
ref 3 4a 5 6 14b 15 16 17 a Formulated from column.
ko[OH-]0.36[03]2 k, [OH- I [O,] + k,[ OH]'.'[ 0, ]I.' k [ OH- ] [ 0, ] ko[OH-]0*'5[0 3 ko[OH-]'.'[ O,! k [OH[ 0, 3 ko[OH']X[0,]S'2, x is pH dependent h o [OH-]'.''[ O,] a mechanistic model. Packed
*
hydroxide ion, there are also substantial differences regarding the pH dependence of the reaction, as shown in Table 11. These major differences among the findings of the various researchers are believed to be due to (a) the use of different, sometimes questionable, analytical techniques to measure the concentration of dissolved ozone, (b) uncertainties in the data analysis and data interpretation,
0 1982 American Chemical Society
Environ. Sci. Technol., Vol. 16,
No. 7, 1982 377
-
,z
by pa ss valve
Reactor
exhaust
Rotameter
Needle va
,
,
Soap-film column
Ozone Generator
Rotameter
water in
-
Needle valve
Oxygen tank
-1
twater
trap
jacket by-pass valve
KI traps
Flgure 1. Experlmental apparatus.
(c) the effect of solution composition, e.g., the ionic medium, on decomposition, and (d) the possible presence of impurities in the reagents used. Additionally, most of the researchers studied the decomposition reaction in batch reactors and analyzed their data only by the integral method. Method of Approach In this investigation, the kinetics of ozone decomposition were studied under both static and dynamic conditions. As a first step, the reaction was carried out in a batch reactor under quiescent conditions. Measurement of the ozone concentration with time allowed the determination of the order (n) and the reaction rate constant (kd)in accordance with ros = -d[O,]/dt = kd[031n
(1)
The decomposition kinetics were also studied in an ozone contact column at constant temperature and under dynamic conditions whereby ozone was bubbled continuously through buffered solutions of known ionic composition to provide complete mixing. Under such conditions, the rate of change of ozone with time can be expressed as d[O,l/dt = k d o , l * - [OS])- Tos
(2)
where kLa is the mass transfer coefficient governing the absorption of ozone by the solution, [O,]* is the saturation concentration of ozone in accordance with Henry's Law and the applied partial pressure of ozone in the gas stream contacting the solution, and ro3is the rate of decomposition of ozone. Here, it was assumed that the mass transfer of ozone is controlled by the liquid film since ozone is slightly soluble in water. At steady state, the rate of ozone absorption is equal to the rate of ozone decomposition, or kLa([O,I* - [O,la,) = (r0Jss
(3)
where [O,], is the concentration of ozone at steady-state. The saturation concentration of ozone and the masstransfer coefficient of the system were measured independently in a solution of pH where the rate of ozone decomposition is negligible. Equation 2 for this system would be modified to d[o,l/dt = k~a([o,l*- [Os])
(4)
Concentration of ozone was followed with time until the equilibrium concentration of ozone was reached. This 378
Environ. Sci. Technol., Vol. 16, No. 7, 1982
concentration was taken as [OJ*. kLa was determined from the slope of the straight line obtained according to the integrated form of eq 4
where [O3lOis the concentration of dissolved ozone at t = 0. This mass transfer information from eq 5 and the decomposition kinetics as determined under quiescent conditions (eq 1)were used in eq 3 to predict the steady-state concentration of ozone at various pH values in the ozone contact column. [O,], values were also measured experimentally and compared to the predicted values to confirm the kinetic expression of the decomposition reaction as determined from the batch experiments. Experimental Setup The experimental setup used in this study is shown in Figure 1. Oxygen gas was supplied in standard cylindrical tanks and fed to an ozone generator. The percentage of ozone produced in the oxygen gas was controlled by changing the power input to the generator. The rotameters before and after the ozone generator were frequently calibrated by a soap-film technique during the experimental runs. The soap-film column was connected to the system downstream of the reactor and the KI traps. A bypass valve located after the second rotameter was used to divert the flow to either the reactor or the traps. The needle valves were used to balance the gas flow and maintain it constant into the traps and the reactor. The reactor consisted of a gas washing bottle fitted with a ground-glass joint. A coarse sintered-glass dispersion tube delivered ozone gas to the bottom of the reaction vessel. The height to diameter ratio of the vessel was about 4 1 to maintain good mixing of the solution by the bubbling gas. Since optimization of gas transfer was not the objective of this investigation, no attempt was made during the study to modify the reactor system to improve gas transfer. To keep the temperature of the contents of the reactor constant at 20 OC,water was continuously circulated by a water pump from a 20 "C water bath through a water jacket around the reactor. A liquid sampling port was located about 1 cm from the bottom of the bottle. Excess ozone gas passed out through the top of the reactor into a gas washing bottle containing 2% KI solution. Two gas absorption bottles in series were connected to the system in parallel to the reactor to determine the quantity
of ozone generated and applied to the column per unit time and the partial pressure of ozone in the feed stream.
Experimental Procedure Ozone-demand-free water was prepared by ozonating distilled and deionized water for about 15 min. Residual ozone was then stripped out of solution by nitrogen gas; the resulting water was used to prepare the test solutions. At pH values less than 4, pH was controlled with dilute sulfuric acid. Sodium sulfate, which has been observed to have no effect on ozone decomposition in water (HoignB and Bader (18)),was used to adjust the ionic strength, p , of these solutions of low pH. A KH2P04/K2HP04buffer system was used over the pH range 6-9 and a H3B03/ NaOH buffer system was used over the range 7-10 for pH control. Some overlap of the two buffer systems was examined to test the reproducibility of the kinetic results with respect to the two buffer systems. The ionic strength of the borate solution was adjusted with Na2S04. The KH2P04and KzHP04salts were mixed in precalculated amounts to adjust both the ionic strength as well as the pH of the buffer solutions. In order to observe any possible ionic-strength effects, the decomposition reaction was studied in solutions of ionic strength 0.1 and 1.0. For the first phase of the investigation in the batch reactor, the buffered solutions were brought to 20 OC and then ozonated with an oxygen stream containing about 5% (by weight) of ozone at a flow rate of approximately 0.4 L/min for approximately 15 min. Depending on the pH of the test solution, dissolved ozone concentrations of 5-40 mg/L were reached. The gas flow was then turned off, and the residual dissolved ozone concentration was followed iodimetrically in accordance with "Standard Methods" (19) (procedure 143 B). Potassium iodide solution was buffered at pH 7 by phosphate buffer. The concentration of residual dissolved ozone was also followed by the indigo procedure (HoignB and Bader (20)). This procedure is more sensitive to low ozone concentrations and is believed to be subject to less interference from other oxidizing species that might arise as products of ozone decomposition. The water-soluble trisulfonated potassium salt of indigo (Cl6H7K3NzOl1S3, MW = 616.7) was prepared by treating the water-insoluble indigo (Cl6HI0N202)with concentrated sulfuric acid according to the procedure of Dorta-Schaeppi and Treadwell (21). The procedure was calibrated by using ozone solutions standardized by the KI procedure. So that any interference in the KI procedure arising from decomposition products of ozone would be minimized, the calibrations were made at pH 3. The change in absorbance of the indigo solution was also calibrated against UV absorption measurements of ozone at 260 nm to assure no interference of decomposition products. In the ozone contact column, kLa and [Os]* were measured in a solution with pH 3, a constant ionic strength, and a temperature of 20 "C. The solution was ozonated a t a constant gas flow rate, and the ozone concentration was followed by taking frequent samples of solution at small time intervals and analyzing them iodimetrically. For measurement of the concentration of ozone in the gas stream, the ozonmxygen mixture was diverted around the reactor into the traps containing neutral buffered 2% KI solution for a constant period of time (30-75 s) at the termination of the experimental run. The ozone generated per unit time was assumed to be constant throughout the experimental run. The contents of the traps were titrated with standardized Na2S203,and the ozone concentration in the gas stream was calculated by knowing the flow rate through the traps.
J
P3 mqA
pH = 8.0
8.0
T = 2OoC Borate buffer
7.0
6.0
5.0
4.0
3.0
2.0
0
0.5
1.5 2.0 Time (min.)
1.0
2.5
Figure 2. Decomposition of ozone in water at pH 8.0, at different initial ozone concentrations.
The experimental procedure was repeated for pH values in the range 2-10 to observe the effect of pH on the steady-state concentration of ozone in water at 20 "C. For each experimental run, the gas flow rate, partial pressure of ozone in the gas mixture, and ionic strength of the solution were kept constant. The pH was buffered as described above.
Results and Discussion Kinetics of Ozone Decomposition in Batch Systems. The data collected from the batch experiments over the pH range 2-10 were analyzed by the integral, initialrate, and differential methods. When the integral method was applied, the correlation coefficients for least-square fit to second-order kinetics were always higher than those calculated for first- and 3/2-orderkinetics. For pH values below pH 7, the second-order fit was significantly better than a first- or 3/2-orderfit in a 95% confidence interval. However, statistical analyses of the experimental results at pH values above 7 were unable to distinguish between first-, 3/2-, and second-order kinetics, even in the 90% confidence interval. Consequently, to test the decomposition kinetics further in the basic pH range, the initial-rate method was applied to the experimental data. It was assumed that the initial rate of decomposition (roJ0 is (r03)0= kd[0310n
(6)
where [O3I0is the initial ozone concentration. Then
1% (rO3)O = 1%
k d -k
n 1%
10310
(7)
The initial rates were determined by fitting a polynomial to the experimental data and differentiating at t = 0. The resulting initial rates are shown in Figure 2. In Figure 3, log (roJ0is plotted against log [O3lOfor the experiments at pH 8,9, and 9.5. The slopes of the three lines, determined by least-squares analysis, indicate reaction orders of 1.8, 1.8, 2.0, respectively. Thus, for practical purposes, second-order kinetics represents the data much better than does first-order or 3/2-order kinetics. For a check on whether the application of eq 7 to the rest of the data for t > 0 will also fit second-order kinetics, Environ. Sci. Technoi., Voi. 16,
No. 7, 1982 379
/ I
~
E"
1 -
10-
CO 3 = 4.4 - 8.6 mg/l
30 K I Procedure" *, - - - b y Indigo Procedure
-by
pH 9.5
0 8 0
-m
0
06
-
04
-
v L
cn
0 -
I
/
a3 0.2
-
0 -
0.2 -0.2 I 04
J 0.5
0.6
0.7
log
0.8
IO
0.9
[O3lO
Figure 3. Application of initial rate method ( p = 0.1, T = 20
"C).
0.1
/
0.5 pH = 8.0
0.4
[0,1,=38-8 2 m g / i
CI
0
T = 20°C
0.3
20
1.6
slope
0
L
80
100
Time (sec.)
.
1.9 t 0 . 2
0 oTime = 0 . 5 min
-0.1
ATime = 1.0 min
-0.2
D
d
60
o Borate, p'0.I Borate, p = 1.0
m 0 .I
-0.3
I 40
4/
0 -
pH 7.8
Figure 6. Decomposition of ozone in water at pH 28,demonstrating second-order kinetics (T = 20 OC,p = 0.1, 1.0).
p'ol
0.2
0
II
_c
--
08
-
0.4
-
Time = 1.5 min
/
0.4
0.5
0.6
0.7
0.9
0.8
log LO3] Figure 4. Log-log plot of instantaneous rate of decomposition vs. ozone concentration. -I .6
5
6
7
8
9 PH
Figure 7. pH dependency of secondorder decomposition rate constant over the pH range 6.0-9.5 ( T = 20 "C). pH58 pH40
o
pH22 -123 v
2
0
Figure 5. Decomposition of ozone in water at pH 5 7 , demonstrating second-order kinetics ( T = 20 OC,p = 1.0).
log rOsvalues were plotted against log lo,] at t = 0.5,1.0, and 1.5 min for five different experimental runs at pH 8 (see Figure 4). The data consistently indicated secondorder kinetics for the decomposition of ozone, cofirming the results of the initial-rate analyses. This observation suggests further that the mechanism of the decomposition 380
Environ. Sci. Technol., Vol. 16, No. 7, 1982
reaction does not change during the course of the reaction. So that the second-order reaction rate constants could be determined more precisely, the data were plotted according to the integral method as shown in Figures 5 and 6. Figure 5 indicates that the rate of decomposition of ozone is essentially independent of pH between pH 2 and 4 ( k d = 0.24 and 0.27 L/(mol s) at pH 2 and 4, respectively). Above pH 4, the rate increases markedly with increasing pH. The second-order reaction rate constants are plotted in Figure 7 as a function of pH for the phosphate and borate buffers and for ionic strengths of 0.1 and 1.0. The kinetics of ozone decomposition can be expressed by the relationship ro3 = ko[03]2[OH-]0~65 (8) ko is specific to the chemical composition of the aqueous solution. The order with respect to hydroxide ion is not
,
;
!
E
pH = 2.5 - 3.0
2.0 -
1.5
-
1.0
-
0.5 Gas Flow Rata = 0.63I/mln
0
20
40
60
0
00
Figure 8. Effect of ionic strength on the mass-transfer coefficient of the system.
significantly different (within the 95% confidence interval) for the phosphate and borate buffers. Figure 7 shows that phosphate has a significant retardation effect on the rate of ozone decomposition. Phosphate acta as an hydroxyl radical (OH-) scavenger (22);as a result, the decomposition rate in solutions of higher phosphate concentration ( p = 1.0) is much slower than the rate in the more dilute ( p = 0.1) phosphate solutions. Although the decomposition rate appears to be faster in borate solutions of lower ionic strength, the difference is not significant within the 95% confidence interval. Expressions 9-11 were obtained for the second-order decomposition rate constant by applying the method of least squares to the data in Figure 7 for borate, phosphate with p = 0.1, and phosphate with p = 1.0, respectively, over the pH range 6.0-9.5, at 20 "C. k d = (2.2 f 0.2) x 105[OH-]0~60*0*05 (9) = (1.4 f 0.4) x 104[0H-]0.51'0.17
(10)
= (4.8 f 1.6) x 102[OH-]0~49'0~06 (11) Due to questions associated with the KI procedure (e.g., interferences due to decomposition products of ozone, stoichiometric ratio of iodine liberated/mol of ozone reacting with iodine), the indigo method was also used to measure the residual ozone concentration with time. The application of the integral method to data gathered at pH 7.8 and 8.6 using the indigo procedure is included in Figure 6. The second-order rate constants determined by using the indigo procedure fall within the 95% confidence interval of the observations made with potassium iodide. The conclusion that the decomposition of ozone over the pH range 2-9.5, under batch conditions, conforms to second-order kinetics adds to the confusion reported in the literature regarding the order of the ozone decomposition reaction (see Table I). In order to verify the results of the batch studies, the dynamic approach was adopted to provide an independent measure of the kinetics of the ozone decomposition reaction. Decomposition of Ozone in Dynamic Systems. Absorption and Solubility of Ozone. The mass-transfer coefficient and saturation concentration of ozone were determined for the ozone contact column in a pH 3 solution of Na2S04,where the rate of ozone decomposition is extremely slow. Figure 8 illustrates the application of eq 5 to determine the mass-transfer coefficient of the system. kd
40
60
80 100 (03)gos.mg/l
Figure 9. Solubility of ozone in water at 20
Time ( s e d
kd
20
OC.
Table 111. Comparison of Literature Values for Ozone Solubility with Solubility Coefficients Determined in this Study ref solubility coeff Mailfert: 1894& Luther: 1905& Kawamura,a 1932b Briner & Perottet: 1939& Seidell, 194OC Stumm (23) Nyberg, 1962d Kirk-Othmera ( 2 4 ) this work
0.41 (19 "C) 0.25 (20 "C) 0.31 (20 "C) 0.34 (19.8 "C) 0.35 (20 "C) 0.34 (20 "C) 0.23 (20 "C) 0.30 (20 "C) 0.35 (20 "C), p = 1.0 0.41 (20 "C), p = 0.1 Solubility coefficient calculated from absorption coefficient. Data given by Hoather (25). Data given by Data given by O'Donovan (27). Taylor (26).
It was observed that the ionic strength of the solution has a pronounced effect on the mass-transfer coefficient. This might be due to changes in the interfacial area, "a", caused by changes in the sizes of the gas bubbles, which is a function of the surface tension of the solution and is influenced by ionic strength. The saturation concentration of ozone, [O,]*,was measured for various ozone concentrations in the gas phase, [O,] The solubility coefficient, s (the ratio of [O,]* to [OJ,d, was also observed to be dependent on the ionic strength of the solution. Figure 9 shows that at 20 "Cand atmospheric pressure, s, in units of (mg of ozone/L of water)/(mg of ozone/L of gas), is 0.35 for p = 1.0 and 0.41 for p = 0.1 Na2S04. Table I11 compares the solubility coefficients determined in this study with the values reported in the literature. The results reported herein are essentially in agreement with the literature values, considering that most of the previous researchers have not reported the pH, chemical composition, and ionic strength of their solutions. In Figure 10, the observed steady-state ozone concentrations obtained in the dynamic systems are shown at various pH values when kLa = 60 h-l, [O,],, = 83 mg/L, T = 20 "C,and p = 0.1. The data are represented by the discrete points shown in the figure. At pH values less than 6, decomposition is relatively slow, and under the experimental conditions, [O,],,is equal to [O,]*. Above pH 6, the rate of decomposition becomes significant, and [O,], is determined by both the rate of ozone decomposition and the rate of ozone absorption. The rate of decomposition is not sufficiently fast, however, to make the decomposition Environ. Sci. Technol., Vol. 16, No. 7, 1982
381
Table IV. Calculated Decomposition Rate ConstantsQ Assuming First-, Three-halves-, and Second-Order Kinetics ( p = 0.1,T = 20 "C) n = 311 n= 2
34
30 T = 20'C
26
- 22 0 OBSERVATIONS
F
-Predicted [O3Issby
$ IS r-.
m
second-order klnetlcs
0
u
\
9.5
4
5
6
rz
rz
0.47
0.959
0.04
0.988
3.5
0.991
6.0
0.975
0.60
0.990
56.0
0.992
10.0
0.979
1.34
0.992
176.0
0.999
16.0
0.984
2.74
0.995
480.0
0.996
s)
(borate) a Constants were calculated from a least-squares analysis using the integrated kinetic expressions applied to the batch experimental data shown in Figures 5 and 6.
6 3
s)
(borate)
first -order klnetics
2
rz
s-
9.0
-"-"Predicted [03
