Ind. Eng. Chem. Res. 1994,33, 125-136
125
Oxidation of Mecoprop in Water with Ozone and Ozone Combined with Hydrogen Peroxide F e r n a n d o J. Beltrln,' Manuel Gonzllez, Javier Rivas, and Manuel M a r i n Departamento de Ingenieria Quimica y Energetica, Universidad de Extremadura, 06071 Badajoz, Spain The aqueous ozonation of the herbicide mecoprop has been studied. A positive effect of pH and ozone partial pressure on the oxidation rate of mecoprop has been observed. Depending on the pH value, temperature has little or no influence on the oxidation rate probably due to the decrease of ozone solubility with the increase of this variable. The synergistic effect of ozone combined with hydrogen peroxide was also studied. Concentrations of hydrogen peroxide lower than 0.01 M do not affect the oxidation rate of mecoprop while higher concentrations inhibit the oxidation due to the consumption of both hydroxyl radicals and molecular ozone by hydrogen peroxide. From a mechanism of direct and radical reactions the rate constants of the reaction between ozone and mecoprop have been obtained. The mechanism has also been applied to the oxidation of mecoprop by ozone combined with hydrogen peroxide to determine the rate constant of the reaction between mecoprop and the hydroxyl radical. In all cases the importance of direct and radical pathways of mecoprop oxidation is discussed. Introduction Organic pesticides have been used extensively to control plagues and weeds and to increase the yield of crops in agricultural fields. However, pesticides have also contributed to the pollution of the environment due to their toxic character. Residual pesticide concentrationsin water constitute a menace for human health because of their facility to penetrate and attack the fat tissues of living beings. Among pesticides, chlorophenoxy herbicides are of great application to eliminate broad-leaved weeds in cereal crops and pastures (Worthing, 1987). These herbicides started to be used after World War I1 (Woods, 1974)having been found in many water environmentssince then. Nowadays there is a growing concern to remove these substances from the water. Due to their low concentration in aqueous media the best removal methods are adsorption and oxidation. Thus, some pesticides such as atrazine have been removed with activated carbon (Battaglia, 19791, but this procedure only transfers the contaminants from the water to the adsorbent, it does not destroy them. Therefore, chemical oxidation, especially carried out with noncontaminant oxidants such as ozone and/or hydrogen peroxide at low concentrations or with one of these two oxidants combined with UV radiation, is recommended (Mauk et al., 1976; Brunet et al., 1984; Reynolds et al., 1989). In this work the ozonation of mecoprop (242-methyl4-ch1orophenoxy)propionicacid) has been studied. This pesticide is one of the most important chlorophenoxy herbicides (Reynolds et al., 1989) together with MCPA ((4-chloro-2-methy1phenoxy)acetic acid), 2,4-D ((2,4dichlorophenoxy)aceticacid), 2,4,5-T ((2,4,5-trichlorophenoxy)acetic acid), MCPB (4-(2-methyl-4-chlorophenoxy)butyric acid). However, in spite of its great use in agriculture no works can be found about its treatment in literature. The literature studies on the ozonation of chlorophenoxy herbicides only focus on the four latter, MCPA, 2,4-D, 2,4,5-T,and MCPB (Reynolds et al., 1989). According to these works, this kind of herbicide can be degraded by ozone, though the rate of oxidation varies depending on the nature of the organic compound and the ozone dose applied (Dor6 et al., 1978; Erb et al., 1979;
* Author to whom correspondence should be addressed. 0SSS-5SS5/94/2633-0l25$04.50/0
Dor6 et al., 1980; Legube et al., 1980). These works, on the other hand, deal only with ozone doses, yields of herbicide destruction, and product identification. They report aromatic compounds as first intermediates and carbon dioxide,chloride ion, and saturated carboxylic acids as main final products of ozonation (Benoit-Guyod et al., 1986). Regarding the kinetics, in recent works, rate constants of the direct reaction between ozone and 2,4-D and 2,4,5-T at 20 "C and a pH value between 1.5 and 5 obtained from homogeneous ozonation have been reported by Yao and Haag (1991). The important feature about ozonation is the two ways that ozone can oxidize the organic matter in water. Ozone can react directly with compounds or can act via radicals formed from different routes (Staehelin and Hoign6,1985). This latter way is called an advanced oxidation process (Glaze et al., 1987). The oxidation of organics in water can be increased by the combination of ozone with hydrogen peroxide that leads to the formation of hydroxyl radicals, species of extremely high oxidation character. Therefore, in this paper the synergistic effect between ozone and hydrogen peroxide on the oxidation of mecoprop has also been studied. Identification of intermediates and their competition for ozone and hydroxyl radicals will be the subject of future work. Experimental Section Materials. Mecoprop and hydrogen peroxide were obtained from Aldrich and Merck, respectively. Ozone was produced from oxygen in a SLO Constrema ozonator. Experimental Setup. Experiments were carried out in a 4000 cm3standard glass agitated reactor operating in semibatch mode. A detailed description of the reactor can be found elsewhere (Beltrh etal., 1993). Ozone diluted in oxygen was fed into the reactor through a 6 cm diameter porous plate (mean porosity: 16 pm). Analytical Methods. The remaining concentration of mecoprop was determined by HPLC using a 15 X 0.4 cm i.d. Spherisorb C18 column with particles of 5 pm and a 441 Waters UV detector. A mixture of acetonitrile and water (4:6 by volume) was used as the mobile phase, and detection was made at 254 nm. Ozone was analyzed
0 1994 American Chemical Society
126 Ind. Eng. Chem. Res., Vol. 33, No. 1, 1994 Table 1. Ranges and Values of Variables CMO,M le gas flow rate, m3 B - ~ Pod, Pa 550-1970 agitation speed, rpm buffer system temperature, O C 2-20 PH 2-12 ionic strength, M CHzOzt, M 0.3
1.39 X 1W lo00 H3P01and NaOH 0.1
iodometrically in the gas phase, both at the reactor inlet and outlet, and colorimetrically in the water (Bader and HoignB, 1981). The concentration of hydrogen peroxide was determined colorimetrically using the method of Eisenberg (1943). Total organic carbon was determined using a 915B Beckman Carbon Analyzer. Procedure. The reactor was charged with 2650 cm3of an aqueous buffered solution of mecoprop (lo4 M). The buffer was prepared with phosphoric acid and sodium hydroxide (see Table 1 for conditions). In some experiments hydrogen peroxide, at a known concentration, was also fed to study its synergistic effect with ozone. Once the ozone concentration in the gas leaving the ozonator and the reactor temperature were stabilized the mixture of ozone-oxygen was directed to the reactor and samples were withdrawn at regular intervals for analysis. Determination of Stoichiometry. The stoichiometric ratio, z, of the direct reaction between ozone and mecoprop, calculated following the procedure described in a previous paper (Sotelo et al., 1990),was found to be 1mole of ozone consumed per mole of mecoprop consumed and independent of pH (between pH 2 and 12). Determination of Fluidodynamic and Physicochemical Data. The volumetric mass-transfer coefficient, k ~ awas , obtained following a procedure similar to that of Ridgway et al. (1989). It was found to be 2.02 X s-l. The individual mass-transfer coefficient,k ~was , calculated after dividing k ~ aby the specific interfacial area, a, determined from the absorption of oxygen in aqueous sodium sulfite solutions in the presence of different concentrations of cobalt(I1) sulfate (De Waal and Okeson, 1966). Values of kL and a were 2.32 X lo4 m s-1 and 87.02 m-l, respectively. A summary of results from both methods is presented in the Appendix. Diffusivities of mecoprop and ozone were calculated by using the Wilke-Chang correlation (Reid et al., 1977)and data of Matrozov et al. (1976), respectively. Ozone solubility was determined as a function of ozone partial pressure, pH, temperature, and ionic strength using the equation of Sotelo et a1 (1989).
Results and Discussion Table 1shows the ranges and values of variables applied in this work. Ozonation of Mecoprop. Since ozonation of organics in water usually involves a complex reaction system (the ozone absorbed is consumed by the initial organic and intermediates) it was first decided to check the importance of secondary reactions. Thus, at different reaction periods the amounts of ozone absorbed, Dab,reacted with mecoprop, D,, and accumulated in water, D,,, were measured from the difference between the amounts of ozone measured at the reactor inlet and outlet, from the decrease of mecoprop concentration, and from ozone concentrationtime data, respectively. Some of the results are presented in Table 2 as percentages with respect to the total ozone fed, Df, for the same reaction period considered. From Table 2 it can be observed that the percentage of ozone absorbed varies between 60 and 95 % of the ozone fed to the reactor. It is also deduced that at pH I 7 contribution of direct reaction between ozone and mecoprop represents less than 20% of the ozone fed and increases up to 39%
Table 2. Percentages of Ozone Absorbed, Consumed by Mecoprop, and Accumulated in Water during the Ozonation of Mecoprop. 20 206 20c 2od 20e
201 10 2 20 10 2
20 10 2 20 20
7 7 7 7 7 7 7 7 4 4 4 2 2 2 10 12
75.8 74.4 67.8 84.9 73.8 80.1 66.8 87.1 71.3 69.7 64.5 83.1 75.7 70.4 74.2 93.7
19.8 18.5 13.2 19.4 17.7 16.0 20.8 19.2 20.1 17.8 11.5 16.9 14.6 9.2 23.9 38.0
19.8 14.2 7.4 23.3 33.1 45.0 30.2 47.2 27.5 33.1 48.2 37.3 41.5 55.5 8.1 0.0
Data correspond to the first 2 min of ozonation and 550 Pa of ozone partial pressure unless indicated. t = 4 min. t = 10 min. PO,^ = 1028 Pa. e Po4 = 1520 Pa. f P%i = 1966 Pa.
*
at pH 12 for the first 2 min of reaction (in all cases the percentage is decreasing with the increasing reaction period). On the other hand, the percentage of accumulated ozone decreases with the increase of temperature and pH, becoming zero at pH 12. For the reaction periods considered in Table 2, it is therefore deduced that percentages of secondary reactions (Dab - (D, + D,,))/Df can vary from 6% (pH 2 and 2 "C) to 56% (pH 12 and 20 "C) of the ozone fed. These figures clearly indicate the highly reactive intermediate compounds which are formed in this process whose identification and kinetics will be treated in future works. Influence of Operating Variables. The influence of ozone partial pressure, pH, and temperature was then investigated. Influence of Ozone Partial Pressure. The ozone partial pressure has a positive effect on the oxidation rate of mecoprop which increases when this variable is increased. In most of the experiments dissolved ozone was detected and its concentration increased with the increasing ozone partial pressure. This suggests that the kinetic regime of ozone absorption was moderate or slow (Charpentier, 1981). Concentrations of mecoprop were below the detection limit of the HPLC equipment (10-8 M) in 15 and 5 min when the ozone partial pressure applied was 550 and 1966 Pa, respectively (at 20 "C and pH 7). Influence of Temperature. A positive effect of temperature on the rate of oxidation of mecoprop between 2 and 10 "C at pH 2 and 4 was observed. At temperatures between 2 and 20 "C and pH 7, however, no change was observed on the oxidation rate of mecoprop (see Figures 1and 2a). These results can be attributed to factors such as changes in ozone driving force and rate constants and the importance of secondary reactions. Thus, the increase of temperature leas, on one hand, to an increase of the rate constant of mecoprop-ozone reaction and, on the other hand, to a decrease of ozone solubility. The first effect would increase the oxidation rate of mecoprop while the second one would have an opposing result (as shown in Figure 3b, concentration of dissolved ozone decreases with the increasing temperature). In addition, as observed from Table 2 secondary reactions are much less important for the consumption of ozone at 2 "C and pH values of 2 and 4 (about 5% of the ozone fed). Influence of pH. The effect of the third important basic operating variable, pH, is shown in Figure 3. As can be seen from Figure 3 the increase of pH yields an increase
Ind. Eng. Chem. Res., Vol. 33, No. 1, 1994 127
Time, rnin 8)
10
I
I
t
I
I
5
I
I
I
10
I
Time, min b)
1
1
2
1
1
4
1
1
6
Time, rnin
1
1
a
1
1
,
10
b)
Figure 1. Ozonation of mecoprop. Influence of temperature on the variation of the normalized remaining concentration of mecoprop with time at (a) pH 2 and (b)pH 4. Conditione: Po4 = 550 Pa, CMO 1W M. Temperature, OC: 0 , 2 ; A, 10; 0,20.
-
Figure 2. Ozonation of mecoprop. Influence of temperature on (a) the variation of the normalized remaining concentration of mecoprop with time and (b)the variation of the dissolved ozone concentration with time. Conditions: Po4 = 550 Pa, pH = 7, CMO= lo-' M. Temperature, O C : 0 , 2; A, 10;0,20.
of the mecoprop oxidation rate although the differences are negligible at pH 7 and 10. This effect of pH is expected to occur in the ozonation of dissociating compounds like mecoprop. Thus, previous works (HoignBand Bader, 1983; Beltrh et al., 1993) report that dissociating compounds show a different reactivity toward ozone according to their pK value and the pH of water. In fact, direct ozonation of dissociating compounds involves reactions of ozone with both their neutral and ionized species, the latter usually being much more reactive toward ozone than the former. At a given pH, both species are present in water their percentages being pH dependent. In the case of compounds with acid groups, like mecoprop, the percentage of the ionized species increases with the increase of pH and, hence, the ozonation rate increases. This effect disappears when the compound is completely dissociated, that is, when pH >> pK (Hoign6 and Bader, 1983). In summary, in this type of reaction the rate constant does not vary at very low pH, then increaseswhen pH increases, and finally, when pH is much higher than the pK, of the acid, remains constant. If it is assumed that the pKa of mecoprop is similar to those of 2,4-D (PKa = 2.6) and 2,4,5-T (PKa = 2.7) (values taken from Yao and Haag, 1991),which is a reasonable hypothesisgiven the similarity of their molecular structures, the reactivity of mecoprop with ozone should increase when the pH goes from 2 to 7 as observed experimentally (see Figure 3a). A t pH 7
mecoprop is totally dissociated and a further increase of pH, up to 10, does not lead to a higher oxidation rate of mecoprop. At these conditions the pH is already much higher than the PKa and the direct rate constant does not vary. At pH 10,however,decompositionof ozone catalyzed by the hydroxide ion to generate radicals cannot be neglected (Sotelo et al., 1987) and this reaction will likely compete with the direct ozone-mecoprop reaction. Therefore, at pH > 7 mecoprop can be oxidized both directly by ozone and via hydroxyl free radicals and an increase of the oxidation rate of mecoprop at pH 10 with respect to that at pH 7 would be expected. On the other hand, at pH 10 the ozone available to react directly with mecoprop decreases because of the competitive action of the hydroxide ion and hydroxyl radicals (Staehelin and HoignB, 1985). As a result, the increase of the oxidation rate due to the action of radicals is counterbalanced by the decrease of the direct oxidation rate. The net result, as shown in Figure 3a, is that the oxidation rate of mecoprop at pH 7 and 10 is similar. When pH is further increased up to 12 the oxidation rate increases,the radical reactions (hydroxyl radicals) being the main responsible species of oxidation (see Kinetics section). With respect to the dissolved ozone, as shown in Figure 3b, at a given time its concentration diminishes when pH increases. At pH 12 there was not dissolved ozone at all
128 Ind. Eng. Chem. Res., Vol. 33, No. 1, 1994
I and I'
I and I'
+ 0,
+ OH'
kdi
P
koni
(3)
P'
(4)
The hydroxyl free radicals, 'OH, the species responsible for the radical oxidation (Buxton et al, 1988), are formed through the well-knownmechanism (Staehelinand HoignB, 1985) of the decomposition of ozone. The main steps of this mechanism are shown below. initiation step: 0, + OH01
I
I
I
I
I
I
0
I
I
I
Time, min
i 10
I
I
I
I
I
1
15
ki = 70 M-'&
02*+ HO,'
with PK = 4.8
a)
+ O,'-+
HO,'
H+
propagation steps:
-
ka = 1.4 X l@ E - ~
HO,' OH' 0
4
,
6 Time, rnin b)
a
10
Figure 3. Ozonation of mecoprop. Influence of pH on (a) the variation of the normalized remaining concentration of mecoprop with time and (b) the variation of the dissolved ozone concentration with time. Conditions: Pod = 550 Pa, temperature = 20 OC; CMO= l eM. pH: 0, 2; A, 4; 0,7; V, 10; 0 , 12.
which confirms its decomposition via radicals and the absence of the direct reaction with mecoprop. The presence of dissolved ozone at pH 2-10 indicates that the regime of ozone absorption can be slow or moderate, and its absence at pH 12 suggests that the ozone absorption corresponds to a fast or moderate reaction (Charpentier, 1981). Mechanism and Kinetics of the Ozonation. As deduced from data shown in Table 2 and previous works (Staehelin and HoignB, 1985) ozone is consumed in multiple reactions constituted by direct and radical steps. Regarding the disappearance rate of mecoprop two contributions have to be considered:
direct reaction: radical reaction:
M + 0,
kd
M + 'OH
I
(1)
kOHM
I'
(2)
Intermediate compounds, I and 1', undergo similar reactions:
+ 0,
OH'
k8
p
2.8 X l@M-'
--A
+ 0,
E-'
HO,'
+0 2
Termination steps are the reaction between hydroxyl radicals and the organics, mecoprop, and intermediates I and/or 1'. A. Kinetics at pH I7. The importance of both contributions (direct and radical, reactions 1 and 2) is different according to the pH of the water. Thus, following previous results (Sotelo et al., 1987) it can be accepted that the decomposition of ozone starts to be important at pH values higher than 7 though depending on the value of the direct rate constants of the reactions involved, the radical pathway could be competitive even at acidic pH (HoignB and Bader, 1983). For the case of mecoprop ozonation, it was assumed that the radical contribution is negligible at pH 7 or lower. Therefore, for these pH values the oxidation rate of mecoprop is exclusively due to the reaction (1):
According to eq 10 a plot of -FM vs CO$M should yield a straight line of slope equal to k d . This plot, prepared from experimental data shown in Figures 1-3, is presented in Figure 4. Values of -FM were calculated from polynomial regression analysis of the curves C w t . As can be observed from Figure 4 points follow straight lines regardless of temperature and pH except in some cases corresponding to high reaction times for which the dissolved ozone concentration is close to its stationary value. Values of k d were determined by least squares fittings and are presented in Table 3. With these data and the experimental conditions used the Hatta number, Ha,corresponding to the irreversible gas-liquid direct reaction between ozone and mecoprop was obtained. This dimensionless number that relates the chemical and physical rates of ozone is
Ind. Eng. Chem. Res., Vol. 33,No. 1,1994 129
/ I
I
2
DH 7
I
I
4
I
6
/ f
Figure 4. Variation of mecoprop ozonation rate with CO$M at different temperatures and pH 2 (top left), 4 (top right), and 7 (bottom). Conditions: Pod = 550 Pa, CMO= lo-' M. Temperature, O C : 0,2;A, 10;0,20. Table 3. Calculated Rate Constants of the Direct Reaction between Ozone and Mecoprop and Initial Rates of Ozone Absorption pH T,OC kd,' M-'8-l NoJt-0 X lo', M s-' kd,b M-'6-l 1.07 15.3 2 2 13.7 33.0 29.2 1.62 2 10 37.9 40.2 1.64 2 20 1.49 22.1 21.7 4 2 54.8 50.4 2.28 4 10 77.0 78.3 2.41 4 20 38.4 41.5 2.25 7 2 3.05 86.9 7 10 92.8 2.70 101.2 111.2 7 20 a
Calculated from eq 10. b Calculated from eq 14.
given by eq 11. At pH S 7, the Hatta values calculated by
eq 11 were always found to be lower than 0.02 which corresponds to a slow reaction (Charpentier, 1981). In order to confirm the kinetic data obtained a second method was applied based on the film theory concept for a slow-moderate irreversible gas-liquid reaction (Van Krevelen and Hoftijzer, 1948). The ozone absorption rate is given by the following equation:
= kLuC03
cos
tanh(Hu)
Co,* cosh(Ha)
]
(12)
where No, represents the ozone absorption rate and Co3* is the ozone concentration at the gas-water interface. Since reactions were carried out in a standard well-agitated tank where perfect mixing was assumed to hold, Cos*represents the concentration of ozone in water in equilibrium with the gas leaving the reactor, that is:
Po,,
co,* = He PO,, and He being the ozone partial pressure at the reactor outlet and the Henry constant for the ozone-water system, respectively. Note, however, that kinetic eq 12 can only be applied to the start of ozonation ( t = 0) because of the competitive effect of secondary reactions in the consumption of ozone (see Table 2). At this moment, the concentration of dissolved ozone is negligible (Cos N 0) and eq 12,once eq 13 has been accounted for, simplifies to the following:
Nogltl0, the initial ozone absorption rate, was obtained assuming that ozone is only consumed by mecoprop and hence:
130 Ind. Eng. Chem. Res., Vol. 33, No. 1, 1994
where z is the stoichiometric ratio (1 mole of ozone consumed per mole of mecoprop consumed). Values of NO&O are given Table 3. On the other hand, the ozone partial pressure at the reactor outlet, Poso,was obtained from the gas law and the difference between the ozone molar flow rate at the reactor inlet and the amount of ozone consumed by mecoprop according to the stoichiometry. By a trial and error procedure values of k d were determined from eq 14. As shown in Table 3 there is a good concordance between the rate constants obtained from both methods. From Table 3 it is observed that k d increases with pH, which is in agreement with the dissociating character of mecoprop (HoignB and Bader, 1983) for the range of pH values considered. B. Kinetics at pH 1 7. Under weakly or strongly basic conditions, the ozone decomposition reaction cannot be neglected and the oxidation of mecoprop (and intermediates) is due to both contributions, direct and radical reactions, or simply to the latter as shown later for pH 12. Thus, the oxidation rate of mecoprop due to reactions 1 and 2 is now
dCM
-
-rM = - dt - kdCMC03 + kOHMCMCOH*
(16)
where COHOrepresents the concentration of hydroxyl radicals. The concentration of hydroxyl radicals can be expressed as a function of the concentration of molecular compounds (mecoprop, ozone, and intermediates) after applying the stationary state hypothesis in the mechanism formed by reactions 1-10. Thus, eq 16 becomes -rM =
dC M --= dt
kiCO,COHkdCOsCM
+ kOHMCM kOHMcM
d:r
- c03cM
[
kd +
2kikOHMCOH-]
x 1 O’O, M2
Figure 5. Variation of mecoprop ozonation rate with C ~ C M at pH 10. Conditione: Pod = 550 Pa, CMO= 1W M. Temperature = 20 OC.
Equation 18 was not applied to data corresponding to pH 12, since in this case the concentration of ozone was under the detection limit of analysis (X10-e M). This suggests that radical reactions are more important due to the increase of pH. As a consequence a different kinetic regime of ozone absorption develops. In fact at pH 12 the ozone decomposition reaction becomes a moderate reaction (3kiCo~-= 2.1 s-l) and its Hatta number is betweeen 0.3 and 3 (Beltrh et al., 1992). Under these circumstances, it was assumed that the oxidation of mecoprop was exclusively due to radical reactions (kdco,
