Ind. Eng. Chem. Res. 1987,26,39-43 Villadsen, J.; Michelsen, M. L. Solution of Differential Equation Models by Polynomial Approximation; Prentice Hall: Englewood Cliffs, NJ, 1978; p 304. Villadsen, J.; Stewart, W. E. Chem. Eng. Sci. 1967, 22, 1483. Vrinat, M. L. Ph.D. Dissertation, Claude Bernard University-Lyon I, Lyon, 1981. Vrinat, M. L. Appl. Catal. 1983, 6, 137.
39
Wakao, N.; Smith, J. M. Chem. Eng. Sci. 1962, 17, 825. Wang, J. B.; Varma, A. Chem. Eng. Sci. 1978, 33, 1549. Wang, J. B.; Varma, A. Chem. Eng. Sci. 1980, 35, 631. Yortsos, Y. C.; Tsotsis, T. T. Chem. Eng. Sci. 1982, 37, 237.
Received for review February 21, 1985 Revised manuscript received February 3, 1986
Ozone Decomposition in Water: Kinetic Study Jos6 L. Sotelo,* Fernando J. Beltrln, F. Javier Benitez, and Jesiis Beltrkn-Heredia Departamento de Quimica Tgcnica, Facultad de Ciencias, Universidad de Extremadura, 06071 Badajoz, Spain
Kinetic studies on ozone decomposition in water were performed over a range of temperatures from 10 to 40 "C and p H range from 2.5 t o 9. T h e ozone decomposition rate was determined in eq 24. This expression is supported by a reaction mechanism. The second term is negligible, a t p H values below 3, leading to a first-order kinetic expression. The direct ozonewater reaction and the hydroxide ion initiation step are the main causes of ozone decomposition. Experimental and calculated ozone concentrations agree within &lo% for 95% of the experiments. The use of ozone as an oxidizing agent in both drinking and wastewater treatment and in several processes for organic synthesis has been growing. In water treatment the rate of ozonation of dissolved pollutants is affected considerably by temperature and pH. These variables also influence ozone's self-decomposition,whose kinetics must be known in order to achieve a satisfactory design. This subject has been widely studied by many authors, as can be seen in Table I. Recently, in 1982, three new reports about this subject have been published (see Table I). Forni's work (1982) studies the ozone decomposition under conditions different from those in drinking water processes, using extremely high pH values above 10. Gurol and Singer (1982) indicated that ozone decomposes by a second-order reaction with respect to ozone concentration, but no mechanism was presented in order to support their kinetic equation. Finally, Staehelin and HoignVs work deals mainly with the hydroxide ion's initial action on ozone decomposition. For this purpose they inhibited the action of very active radical species, such as hydroxyl radicals, OH*, using radical scavengers such as carbonate or methylmercury ions. In these circumstances, there is competition between the reactions k = 3 X los mol/(L.s) OH' + O3 0 2 + HOz' (1) k = 4.2
OH' OH'
+ C032k = 1.5 + HCOB-
X lo* mol/(L.s)
X 10'
mol/(L.s)
OH-
+ C03'-
(2)
OH-
+ HC03'
(3)
So, the kinetic equation presented in this work, deduced only a t 20 "C,corresponds to the initiation reaction between ozone and the hydroxide ion. These authors also study the hydrogen peroxide action on ozone decomposition. This is particularly important when the water to be treated has organic micropollutants whose decomposition by ozone produces hydrogen peroxide via organic peroxides (Schultze and Schultze-Frohlinde, 1975). In general, from the works listed in Table I, it can be concluded that there is disagreement among the previous results. Most of these investigations are not supported by reaction mechanisms, and generally, the range of conditions investigated was very narrow; some used only one pH and one temperature. The reaction orders for the ozone and hydroxide ion varied very much, being the main point of disagreement (Table I). The only agreement reached by all the works is that the reaction is catalyzed by the 0888-5885/87/2626-0039$01.50/0
hydroxide ion. Hence, the aim of this work is to deduce a general kinetic expression for ozone decomposition in water involving the action of different species which greatly accelerate the process, such as OH' and OHz' radicals, and which are supported by a reaction mechanism involving these radicals. In addition, since temperature is generally the most important variable affecting reaction kinetics, it would also be interesting to study its influence over a wide range of values.
Experimental Section Materials. Water was first deionized by ion exchange and then distilled. A mixture of KH,PO, and Na2HP04 was used to adjust both the ionic strength (0.15) and the pH of the solutions. The ionic strength was kept constant because this parameter affects ozone decomposition (Gurol and Singer, 1982). Potassium iodide, analytical grade, was used for the analytical measurements. Oxygen was taken directly from a commercial cylinder and dried with silica gel traps before entering the ozonator. Procedure. An ozone-oxygen mixture was produced in an ozonator (SLO Constrema) which is able to operate a t a maximum rate of 6 g of 03/h. The reactor was a 750-cm3 glass vessel with inlets for bubbling, stirring, sampling, venting, and temperature measurement and was submerged in a bath equipped with stirring, heating, and refrigeration systems, which allowed us to keep the temperature constant within rt0.5 "C. Stock buffered water (500 cm3)was added to the reactor and then ozonated with an 02-03 stream, containing about 4% (v/v) ozone, at a flow rate of 40 L/h for 30 min. About this time, saturation of water was reached, and the gas stream was stopped. Subsequently, the ozone decomposition chemical reaction was followed by determining the concentration of dissolved ozone by reaction with a buffered potassium iodide solution, measuring the triodide ions liberated spectrophotometrically at the wavelength of 352 nm (Shechter, 1973). Results and Discussion Effects of Agitation Speed, Temperature, and pH. Though the process studied is only a homogeneous chemical reaction, agitation was provided to keep the temperature of the water uniform during the experiment. Agitation speed was varied between 100 and 700 rpm. A t speeds below 300 rpm, the ozone conversion remained constant at a given reaction time. However, the decom0 1987 American Chemical Society
Ind. Eng. Chem. Res. Vol. 26, No. 1, 1987
40
Table I. Summary of the Kinetic Studies of Ozone Decomposition in Watern reaction order with respect to activation energy, temu, "C ozone hydroxide ion kJ/mol 2-4 0 2 0.37 2 0.37 5-8 0 1.5-1 0.5-1 2-8* 0 58.6 0.27 1 0.5 0.5-3 111.8 1.2-19.8 1 0.75 7.5-10.5 1.5 0-7 25 2 0.63 8-10 25 10-13' 25 1 1 0-2 5-40 1 or 2 59.0 2d 40-50 2 81.2 4d 30-60 2 96.7 6d 10-50 1.5-2 96.2 8d 10-20 1 70.5 12-13.5' 18-27 1 1 96.3 77.5 3-35 1.5-1 0.28-1 2-95 46.4 3.5-60 1 0.12 0.5-10 1 0.88 1-13.5' 25 11-13'" 20 1 1 2-10 20 2 0.55 20 1 1 8-10h
ref Rothmund et al., 1913 Sennewald, 1983 Weiss, 1935 Alder and Hill, 1950 Stumm, 1954 Kilpatrick et al., 1956 Kilpatrick et al., 1956 Czapski et al., 1968 Merkulova et al., 1971 Hewes and Davison, 1971 Hewes and Davison, 1971 Hewes and Davison, 1971 Hewes and Davison, 1971 Rizzuti et al., 1976 Morooka et al., 1978 Sullivan and Roth, 1980 Teramoto et al., 1981 Forni et al., 1982 Gurol and Singer, 1982 Staehelin and HoignB, 1982
5 , supported from a reaction mechanism. Reactions carried out in agitated tanks. *Kinetic expression: k,[OH-] [O,] + kz[OH-]0~5[03]1 Stopped-flow reactor. dTubular reactor. "Packed column. f Kinetic expression: kl[OH-]0~28[03]1~5 + k z[OH-][O,]. BKinetic expression: 2880[0H-][O3] in mol/(L.min). ,+Kineticexpression: 4200[OH-] [O,] in mol/(Lmin).
1.o
1.o
m -
0 0.8
2m 0.8
20
> m
,-.
.-.
9
-e.-
g,
-..
P,
0.6
0.6
I
0.4
0.4
0.2
0.2
0
5
10
15
20 t,min
Figure 1. Influence of agitation speed on ozone decomposition. Experimental conditions: temperature, 30 "C; pH 7; agitation speed, (0) 100, (A) 300, (0)500, (v)700 rpm.
position rate was found to be higher when agitation speed was above 300 rpm, probably due to ozone desorption (Figure 1). Hence, the rest of the experiments were performed a t 100 rpm. Experiments were carried out while varying the temperature from 10 to 40 "C and pH from 2.5 to 9. When these variables were increased separately, a higher ozone conversion a t a given reaction time was observed (see Figures 2 and 3). Decomposition was extremely fast in experiments performed a t temperatures and pHs higher than 30 "C and 8, respectively. Interpretation of Kinetic Data. It was previously assumed (Table I) that ozone decomposition follows simple first-order kinetics with respect to ozone concentration; if so, its decomposition rate can be expressed as (4)
where
0 25 50 75 100 t, min Figure 2. Influence of temperature on ozone decomposition. Experimental conditions: pH 7; agitation speed, 100 rpm; temperature, (0) 10, ( A ) 20, (0)30, (V)40 "c. 1.0
1
h
-20.8 0
Y
Lm
B -70.6 ....
0.4
0.2
0
10
20
30
40
50
t,min
Figure 3. Influence of pH on ozone decomposition. Experimental conditions: temperature, 30 "C; agitation speed, 100 rpm; pH, (0) 2.5, (A)5 , (0) 7, (V)8, ( 0 )9.
Ind. Eng. Chem. Res. Vol. 26, No. 1, 1987 41 Table 11. Values of the Two Reaction Rate Constants, kT and k A kT, L1/2/(mol'/z.min),at pH temp, values 1o2kA,min-', "C 5 7 8 9 at pH 2.5 1.498 7.23 10 0.071 0.422 0.82 40.2 1.41 2.405 7.39 20 0.638 25.62 70.5 2.32 30 0.881 4.66 40 2.943 11.73 84.92 200.3 4.50
which represents the initial hydroxide ion attack. From the sequence of reactions 7-11, we obtain 4031
-03
= -- dt
- k,[O,l[OH'l
+ kz[O,l[HO,'l (12)
By assuming steady states for the concentrations of HO' and HOz' radicals, eq 12 leads to
t, min
Figure 4. Ozone decomposition in water for first-order kinetics. Experimental conditions: pH 7; agitation speed, 100 rpm; temperature, (0) 10, (A)20, (0)30, (V)40 O C .
(
-ro3 = 2 k i [ 0 3 ]+ 2kz
)
2kj + kjJOH-] '/' k,
(13)
Since ozone decomposition rate is very fast at a high pH, it can be assumed that 2ki is negligible when compared to ki' (reactions 7 and €9, so that eq 13 becomes -ro3 = 2ki[0,]
+ 2k2(
E)
'/2
[OH-]'/2[03]3/2 =
k ~ [ 0 3 ] kB[OH-]1'z[0,]3'2 (14) where and
u 3
2
4
5
6
7
8
By use of
9,,~
Figure 5. kd-pH dependency. Experimental conditions: pH range, 10, (A)20, (0) 30, ( 8 )40 O C . 2.5-9; temperature, (0)
because hydroxide ion acts as a catalyst. At constant pH, eq 1 can be integrated using the initial concentration of ozone, [O3lO,to yield [O,] = [0310 exP(-kdt) (6) Experimental data corroborated eq 6. (See Figure 4 as an example.) However, when In kd was plotted against pH, experimental points were not on a straight line, as shown in Figure 5. This plot does not agree with eq 5; moreover, it can be deduced that kd remains nearly constant when the pH is lower than 3. Taking into account these results and the reaction mechanism proposed by Weiss (Weiss, 1935), the following reaction sequence can be written: k
O3 + H 2 0 A 2HO'
+ O2
I I
+ OH' 0 2 - + HO2' k o3+ HO- -t, o2+ H O ~ * kn O3+ H02' 2 0 2 $ HO' k,'
0 3
-
2H02' -?+
k 0 2
+ H202
(7) initiation
(8) propagation
(9) (10)
termination
(11)
In this mechanism, reaction 8 is one of the two possible initiation steps proposed by Staehelin and Hoign6 (19821,
integration of eq 14 leads to
(
[:fl/'
) (
1 + -[03]1/2 kA
=
[:tO1/z
)
- ItA -t2
1 + -[03]01/2 kA
and by rearrangement
When the pH is lower than 3, [OH-] has hardly any influence on the ozone decomposition rate and the second term in eq 14 is negligible compared to the first. Whence, at these pH values the process follows first-order kinetics and eq 14 becomes eq 4 and kA = kd. Consequently, the values of kA for different temperatures can be obtained from experiments carried out at pH 2.5. Once k A values are known, eq 19 can be checked. Experimental data fit well a t 10 "C, as shown in Figure 6. Similar results are obtained for other temperatures. A least-squares analysis provided the pseudokinetic constants, kT (see Table 11). These values were plotted against hydroxide ion concentrations in logarithmic coordinates for all temperatures (see Figure 7). The slopes of the four lines, determined by least-squares analysis,
42
Ind. Eng. Chem. Res. Vol. 26, No. 1, 1987 Table IV. Ozone Decomposition Rate Contributions" hvdroxide ion action pH direct decb initialc overaF 2.5 1.43 x io+ 1.12 x 10-12 9.77 x 10-9 5 1.43 X 10" 3.54 X lo-'' 1.74 X 10.' 7 1.43 X lo4 3.54 X lo-' 1.74 X lo4 9' 1.43 x 10-7 3.54 x 10-7 5.50 x 10-7 "Temperature = 20 "C. [Os] = mol/L. (Mean value obtained experimentally at pHs between 2.5 and 7.) bFrom eq 20 and 22. Mean values from Forni et al and Staehelin and Hoign6 kinetic expressions given in Table I. dFrom eq 21 and 23. e [O,] = mol/L. (Mean value obtained experimentally at pH 9.)
1.0
1.1
I
,
1.2
1.3
I
1.4
I
1.5 1.6 ew(Y)
Figure 6. Check of eq 19. Experimental conditions: temperature, 10 "C; PH, (0) 5, (A)7, (0) 8, (VI 9. k,
1
I
10
1
lo-'
-
-
1o - ~
10-9 2
5
168
10.7
6
7
5
[OH-],mol 8
9
PH
L
Figure 7. k ~ p H dependency. Temperature, (0) 10, (A) 20, ( 0 )30, (V) 40 "C. Table 111. Hydroxide Ion Reaction Order and Rate Constant, kg,by Least-Squares Analysis of Lines Plotted in Figure 7 temp, hydroxide ion rate const, kB, "C reaction order L/(moLmin) 10 0.493 1603 20 0.434 3 929 30 0.485 17 545 40 0.475 44 205
indicated that the data consistently agreed with eq 17. From their corresponding intercepts, the true kinetic constant values, kg, were obtained (see Table 111). The influence of temperature on the rate constants, k , and k g , was derived from Arrhenius plots to yield
k, = 3.26 and kB = 5.69
X
X
lo5 exp(-4964/n
10ls exp(-10130/T)
(min-')
(L/(mobmin))
(20) (21)
The activation energy for k,, 41.23 kJ/mol, is very close to values obtained in other work in which a first order with respect to ozone was assumed (Table I). So, the ozone decomposition in water involves two major contributions: direct ozone decomposition, reaction 7, and the overall hydroxide ion attack, initiated by reaction 8,
0
Gurol b S i q e r
A
HEW- %. Davison
0
M o r o o k a et
V
Rizruti et a1
0
St""
CI
I
Sullivcn b Roth
A
This w o r k
V
Forni et CI
a
S t c e h e h n g H o i g nI i
,
i
:-":
Corresponding t o rocction , 181
Ind. Eng. Chem. Res. 1987,26, 43-50 the process. At lower pH, the major contribution is due to direct decomposition (eq 7). Finally, Figure 8 shows the comparison among the ozone half-lives obtained from this work and previous reports updated in Table I. It can be seen that, a t a pH ranging from 7 to 9, the results from this work agree with the others. However, a t pHs below 7, there are discrepancies among the data from different sources, probably due to the different ionic species present in the water in each case. Thus, the works of Hewes and Davison, Morooka et al., and Gurol and Singer, carried out in the presence of dilute sulfuric acid, have ozone half-lives greater than those from Sullivan and Roth's work and our, which were developed with phosphate ions. Similar results are observed a t other temperatures. It can also be observed that ozone half-lives in the presence of ion scavengers (carbonates) are obviously higher.
Conclusions Ozone decomposition at pHs and temperatures ranging between 2.5 and 9 and 10 and 40 "C, respectively, follows the two-term rate eq 24 supported by a reaction mechanism. It was shown that at pHs lower than 7, whatever the temperature was, direct ozone decomposition and the initiation step involving the hydroxyl radicals are the main cause of ozone decomposition. At higher pHs, the importance of the peroxy radicals and of the hydroxide ion initiation step increases. Thus, at pHs around 9, the ozone decomposition rate depends on two major contributions: direct ozone decomposition, which leads to the formation of hydroxyl radicals, OH', and hydroxide ion action, which produces not only peroxy radicals, HOz', but also hydroxyl radicals. Finally, the nature of the ionic species (carbonate, sulfate, phosphate, etc.) present in the water greatly influ-
43
ences ozone decomposition, inhibiting some of the reactions in the mechanism proposed.
Acknowledgment This research was supported by the Comisidn Asesora de Investigacidn Cientifica y T6cnica (Grant 1109/81). Registry NO.03,10028-15-6; HzO, 7732-18-5; OH-, 14280-30-9.
Literature Cited Alder, M. G.; Hill, G. R. J . Am. Chem. SOC.1950, 72, 1884. Czapski, G.; Samuni, A,; Yelin R. Isr. J. Chem. 1968, 6, 969. Forni, L.; Bahnemann, D.; Hart, E. J. J . Phys. Chem. 1982,86,255. Gurol, M. D.; Singer, Ph. C. Environ. Sci. Technol. 1982, 16, 377. Hewes, C. G.; Davison, R. R. AIChE J . 1971, 17, 141. Kilpatrick, M. L.; Herrick, C. C.; Kilpatrick, M. J . Am. Chem. SOC. 1956, 78, 1784. Merkulova, V. P.; Lovchikov, V. S.; Ivanovskii, M. D. Izv. Vyssh. Uchebn. Zaued., Khim. Khim. Tekhnol. 1971, 14,818. Morooka, S.; Ikemizu, K.; Kato, Y . Kagaku Kogaku Ronbunshu 1978, 4, 377. Rizzuti, L.; Augugliaro, V.; Marrucci, G . Chem. Eng. Sci. 1976, 31, 871. Rothmund, V.;Burgstaller, A. Monatsh. Chem. 1913, 34, 665. Schultze, H.; Schultze-Frohlinde, D. J. Chem. SOC.,Faraday Trans. 1 1975, 71, 1099. Sennewald, K. 2.Phys. Chem. 1933, A164, 305. Shechter, H. Water Res. 1973, 7, 729. Staehelin, J.; Hoign6, J. Environ. Sci. Technol. 1982, 16, 676. Stumm, W. Helv. Chim. Acta 1954, 37, 773. Sullivan, D. E.; Roth, J. A. AIChE Symp. Ser. 1980, 76, 142. Teramoto, M.; Imamura, S.; Yatagai, N.; Nishikawa, Y.; Teranishi, H. J . Chem. Eng. Jpn. 1981,14, 383. Weiss, J. Trans. Faraday SOC.1935, 31, 668.
Received for review February 22, 1985 Revised manuscript received February 3, 1986 Accepted March 20, 1986
Measurement of Competitive Adsorption Isotherms by Frontal Chromatography Jana M. Jacobson, John H. Frenz, and Csaba Horv&th* Department of Chemical Engineering, Yule University, New Haven, Connecticut 06520
Two methods are described for rapid measurement of competitive adsorption isotherms by frontal high-performance liquid chromatography (HPLC). T h e first method involves the measurement of the composition of the column effluent with a dual chromatographic system using an in-line analytical HPLC and determination of the isotherms by mass balance. This method is generally applicable t o any type of adsorption isotherm. The second method, based on the theory of multicomponent chromatography, involves regression of the experimentally determined velocities of composition changes in the column t o the parameters of a Langmuir model by a reversed h-transformation. The results indicate that the adsorption from aqueous solution of the model compounds studied conformed t o Langmuir behavior on the nonpolar sorbent studied. Separation by chromatography and many other purification processes is governed by equilibrium adsorption of a compound from a solution. The adsorption equilibrium is most commonly represented by the adsorption isotherm, which relates the solute concentration in the liquid phase to that a t the surface of the solid phase over the concentration range of interest. When the solution contains several sorbable components a t sufficiently high concentration, they mutually influence one another's adsorption, and the isotherms obtained under such conditions are
termed "competitive". In contradistinction, when only one sorbate species is present or the solute concentrations are so low that no mutual effects are manifest, noncompetitive isotherms of the individual species sufficiently characterize the adsorption. Adsorption isotherms have long been a focus of investigation in physical chemistry as a means to study solute-surface interactions and in chemical engineering for process design. They have lately gained attention as well in high-performance liquid chromatography (HPLC) where
08SS-58S5/81/2626-0043$01.50/0 0 1987 American Chemical Society
