Advanced Oxidation of Atrazine in Water—I. Ozonation

Pergamon 0043-1354(94)E0019-3

Wat. Res. Vol. 28, No. 10, pp. 2153-2164, 1994 Copyright 1994 Elsevier Science Ltd

Printed in Great Britain. All rights reserved 0043-1354/94 $7.00 + 0.00

ADVANCED OXIDATION OF ATRAZINE IN WATER--I. OZONATION

FERNANDO J. BELTR,~,N* ~), JUAN F. GARCiA-ARAYA and BENITO ACEDO

Departamento de Ingenieria Quimica y Energ6tica, Universidad de Extremadura, 06071 Badajoz, Spain

(First received February 1993; accepted in revised form January 1994)

Abstract--The ozonation of atrazine in water has been studied under different conditions of ozone partial pressure, pH, temperature and presence of hydroxyl radical scavengers. The process mainly develops through radical reactions, even at pH 2, probably due to the presence of traces of impurities in the starting material. The rate constant of the direct reaction between ozone and atrazine at pH 2 has been determined and expressed as a function of temperature. Thus, at 20C the rate constant was found to be 4.5 M- t s-t the energy of activation being 18.8 kJ mol-t . A model formed by mol balance equations of atrazine and ozone, the latter in the liquid and gas phases, with the ozone decomposition rate based on the Staehelin and Hoign6 mechanism allowed the determination of the concentrations of these species at different experimental conditions. The aqueous solution of atrazine is characterized by a hydroxyl radical initiating rate factor which accounts for other radical reactions not included in the basic ozone decomposition mechanism. With this factor deviations between experimental and calculated concentrations of atrazine are less than + 15% in most of the cases.

Key words--atrazine, advanced oxidation, ozone, ozonation, kinetics, water treatment

NOMENCLATURE

CA = concentration of atrazine (tool 1 -~ ) CAo = initial concentration of atrazine (mol 1- ~ ) (;'As = concentration of atrazine defined in equation (15)

(moll a) CB = concentration of species B (mol 1 - z ) C~ = concentration of species I that initiates the for-

mation of hydroxyl radicals (mol 1- J ) Con = concentration of hydroxyl radicals (mol 1-1)

Con~p = experimental concentration of hydroxyl radicals calculated by equation (41) (tool 1-1)

Co~-= concentration of superoxide ion radical (tool 1-1 ) Cos = concentration of dissolved ozone (tool 1- ~ ) C~ = concentration of ozone at the gas-water interface

or solubility (tool I 1) Co,g = concentration of ozone in the gas at the reactor

outlet (tool 1 ~) Costa = average value of dissolved ozone concentration

(mol 1- J ) Cojs = stationary concentration of dissolved ozone

(mol I l) Dos = diffusivity of dissolved ozone (m 2 s- i ) H = Henry constant (Pa I mol -t )

Ha = Hatta number defined in equation (28) (dimensionless)

Koa =overall mass transfer coefficient [mol (lgas) -~ Pa 1 s- l ]

Koa' = overall mass transfer coefficient (tool 1-i Pa-1 s- l ) k = rate constant of the irreversible reaction between

ozone and species B (1 tool- 1 s- 1 ) kA = actual rate constant of the direct reaction between

unprotonated form of atrazine and ozone (1 tool -I s - l )

k d = rate constant of the direct reaction between atrazine and ozone at pH 2 (1 tool -j s -~)

*Author to whom correspondence should be addressed.

k~ = pseudo first order rate constant defined by equation (14) (s - i )

k~ = rate constant of the hydroxyl ion catalysed decomposition reaction of ozone (1 tool- i s- ~ )

k t = rate constant of reaction (38) (1 tool -I s - I) ke = individual liquid phase mass transfer coefficient

(ms- ' ) ko~t, A = rate constant of the reaction between atrazine and

the hydroxyl radical (1 tool -i s- ~ ) k~ = rate constant of the reaction between the intermedi-

ate i and the hydroxyl radical (1 tool- ~ s- ~ ) k~ = rate constant of the reaction between the scavenger

i and the hydroxyl radical (1 tool- 1 s - l ) k z = apparent pseudo first order rate constant of the

ozonation of atrazine defined in equation (40) (s-i ) kt = rate constant of second order termination for

hydroxyl radicals (1 tool- 1 s - l ) Posi = ozone partial pressure at the reactor inlet (Pa)

m = reaction order of ozone in its irreversible reaction with species B (dimensionless)

n = reaction order of species B in its irreversible reaction with ozone (dimensionless)

R = gas constant (J K - ~ tool- ~ ) T = temperature (K) t = time (s) t~ = time to reach the stationary concentration of ozone

(s) V o = gas volume accumulated in the reactor (1) V c = liquid volume in reactor (1) v 0 = gas volumetric flow rate (1 s- t ) z = stoichiometric ratio of the direct reaction between

atrazine and ozone (dimensionless) Greek letters

= degree of protonation of atrazine (dimensionless) fl = hydroxyl radical initiating rate factor (tool 1 - ~ s- ~ ),

defined by equation (42) )'o3 = percentage of the elimination of atrazine due to its

direct reaction with ozone, defined in equation (39) (dimensionless)

2153

2154 FERNANDO J. BELTR./~N et al.

INTRODUCTION

Atrazine is one of the most used herbicides worldwide and constitutes a priority pollutant of the water (Battaglia, 1989). Since atrazine has been found in many water environments at very low concentrations [although higher than its maximum contaminant level (Croll, 1989)], oxidation is one of the best methods to degrade it. In a previous paper (Beltrfin et al., 1993) it was shown that the combination of ultraviolet radiation and hydrogen peroxide lead to high degradation rates of atrazine. Accordingly, advanced oxidation processes can be a convenient way to oxidize atrazine.

This work, divided in two parts, deals with the oxidation of atrazine with ozone and u.v. radiation (03 and u.v./O3), processes that like the combination of u.v. radiation and hydrogen peroxide (Beltr~n et al., 1993), generate hydroxyl radicals. In the first part of this work, ozonation of atrazine is studied, the major aims being the determination of the rate constant of the direct reaction ozone-atrazine, the comparison of the importance of direct and radical ways of ozonation and the study of the oxidation kinetics. In the second part, the u.v./O3 oxidation results are presented (see subsequent paper, Beltr~n et al., 1994).

The ozonation of atrazine in water has already been treated by different authors to identify the main by-products formed and establish molecular mechan- isms. Chronologically, Erb et al. (1979) were the first to study this process to report some data about the degradation rate of the herbicide. Later, Legube et al. (1987) evaluated the ozone demand of atrazine solutions at slightly acid pH, but the main objective of their work was to identify secondary products. In another work Adams et al. (1990) also studied the ozonation of atrazine and derivative compounds, like deethylatrazine, deisopropylatrazine and 2-chloro- 4,6-diamino-s-triazine to quote a few. Like Legube et al., they proposed a molecular mechanism indicat- ing the importance of contact time, pH and ozone concentration on the rate of reactions. More recently, Adams and Randtke (1992) reported additional data about the ozonation by-products, highlighting the differences in the nature of these intermediates when carbonate ions are present.

As far as kinetics is concerned, Adams et al. (1990), Yao and Haag (1991) and Xiong and Graham (1992) have reported on values of the rate constant of the direct reaction between ozone and atrazine at specific conditions of temperature (20C) and pH. Table 1 lists their results. The lack of information about the contribution of radical reactions and discrepancies in the rate constant data made necessary further studies, which were one of the objectives of this work.

EXPERIMENTAL SECTION AND ANALYSIS

The experimental set-up basically consisted of an ozonator (SLO Constrema) and a 850 cm 3 glass reactor described

Table I. Rate constants of the direct reaction between ozone and atrazine reported in literature

k d Scavenger pH (M-Is ~ ) and concentration Reference 2 a 2.25 t-Butanol, 0.01 M Xiong and Graham (1992) 7.5" 12.24 Carbonate, 10-4M Xiong and Graham (1992) 4.1 b 6 t-Butanol, 0.01M Yao and Haag (1991) 5 b 73 c -- Adams et al. (1990) 7 b 146 -- Adams et al. (1990) 9 b 617 -- Adams et al. (1990) 7 ~ 12Y -- Adams and Randtke (1992) aBatch competitive ozonation with MCPA as reference compound.

bBatch ozonation. CReported as total rate constant, dContinuous ozonation.

in detail previously (Beltr~n et al., 1993). Ozone produced in the ozonator from oxygen was fed into the reactor, filled with 800 cm 3 aqueous buffered solution of atrazine, through a porous plate (of 16 #m mean porosity) situated at the bottom. The experimental conditions are given in Table 2.

Atrazine was analysed by liquid chromatography as described before (Beltr~in et al., 1993). Ozone both in the liquid and in the gas phase, at the reactor inlet and outlet, was measured colorimetrically (Bader and Hoignr, 1981) and iodometrically, respectively. In some experiments the ozonation was carried out in the presence of known amounts of t-butanol (pH2) or bicarbonate/carbonate ions (pH 7 and 12) that acted as scavengers of hydroxyl radicals.

RESULTS AND DISCUSSION

First the influence of variables (see Table 2) on the degradation rate of atrazine was studied. Both the ozone partial pressure and pH exert a strong effect, increasing the atrazine conversion with the increase of each of these variables. Figure 1 shows, as an example, the effect of pH. These results are, on one hand, a logical consequence of the increase of the ozone driving force by increasing the ozone partial pressure and, on the other hand, of the higher decomposition of ozone into radicals when pH increases. As shown in Fig. l(b), the concentration of dissolved ozone increases with time, only for a few minutes, until reaching a stationary value. At any given time, this concentration decreases with the increasing pH probably due to the more important contributions of radical reactions at high pHs. In any case the most noticeable fact, the presence of dissolved ozone, suggests that the ozonation of atrazine, regardless of the mechanism (through direct or radical reactions), is a set of slow gas-liquid reactions. At this point, however, it has to be pointed out that at

Table 2. Experimental conditions applied in this work CAo(M) 4.2 10 5to 7.3x10 5 Po3i (Pa) 500 to 2000 T (C) 3, 10 and 20 pH 2, 7 and 12 C s as t-butanol (M) 0 to 0.05 C~ as total bicarbonate (M) 0 to 0075 Gas flow rate (l h -t) 50 Ionic strength (M) 0.01

Ozone oxidation of atrazine 2155

1.0

~ OJ o

I \a" !~ \ \ ' \ \

6 \ \ , , -t, \{ \ \ - I t a \ o x.

w

N

I l l i l 0 4 8

"~o ~

i i i i J i i i ~) i 12 16 20 24 3

time, mln (a)

0

10

-0

s, ~' o I

;

4 8 12 16 20 24 30

time, rain

(b)

o o

@ o o

e" 0 e" - - "~ - - - - -- - - - - " " ~

(o o

o

o o

0 4 8 1 6 20 24 30

time, mkl

(c)

Fig. 1. Influence of pH on the ozonation of atrazine. Variation of concentrations with time: (a) atrazine as CA/CAo; (b) dissolved ozone; (c) ozone in the gas. Conditions: C) pH 2; [] pH 7; ~ pH 12.

Po3~ = 1050 Pa, CA0 = 5 x l0 -5 M (average values).

pH 12 (see later) the very low concentration of dissolved ozone indicates a change in the kinetic regime of absorption that becomes moderate (Charpentier, 1981).

A similar pattern to that shown in Fig. l(b) was followed by the concentration of ozone in the exiting gas [see Fig. l(c)]; in this case, however, the stationary concentration of ozone is apparently independent of pH.

Figure 2 shows the effect of inhibitors of hydroxyl radicals on the degradation of atrazine. Two strong inhibitors were used, t-butanol and bicarbonate/ carbonate ions in addition to phosphates to buffer the aqueous solutions that are also weak scavengers (Staehelin and Hoign6, 1982). The effect, as observed from Fig. 2, was checked at three pH levels (2, 7 and 12). It is evident that atrazine is degraded through radical reactions even at pH 2, so that the ozonation of atrazine is mainly an advanced oxidation, although at acid pH the contribution of its direct reaction with ozone can be important as shown later.

Finally, the effect of temperature is shown in Fig. 3. Here, results corresponding to two series of ozonation reactions are plotted. In one series, the ozonations were carried out at the same ozone partial

pressure whilst in the other the effect of this variable on the ozone solubility was corrected and different ozone partial pressures were applied. Thus, when the ozone partial pressure was the same (about 1100 Pa) there were no appreciable differences between the values of the remaining concentration of atrazine with temperature. Since the degradation rate of atrazine is proportional to the rate constant, and hence to temperature, one should expect, a priori, a positive influence of this variable on the oxidation rate of atrazine. However, as ozonation is a gas liquid reaction, the oxidation rate also depends on the ozone solubility which decreases with the increasing temperature (Sotelo et al., 1989). Thus, both effects, increase of the rate constant and decrease of the ozone available for reaction, are opposed and the final consequence is that temperature does not exert any significance influence on the elimination of atrazine, at least under the conditions applied in this work (3-20C).

In the second series of ozonations, the ozone partial pressure applied was corrected so that the ozone solubility, and hence the ozone available for reaction, was the same at different temperatures. These experiments were also carried out in the

2156 FERNANDO J. BELTR~N et al.

1.0

0.5

O

0

1.0

i" ..a t-Butanol \X "~ ~ pH2 COx 10 2, M '~.,p~ "....=.. o=, 01

\ "~ ~"~t, ,e, 5 \ \ ,'-.. , , . . ,.. o N

i i i i , i i i , , , i t i , 4 8 12 16 20 24 28

time, min Bicarbonate ion in

I'~ pH 7 pH 12 C s, M ,,,

~,%~ - - - - n o 0.075 \ ~'\q\\

*~ \ ' k2 , }'s I,, k ...

I o \ \

i t i i i i i i i i i i i i i

0 4 8 12 16 20 24 28

time, rain Fig. 2. Influence of hydroxyl radical scavengers on the ozonation of atrazine. Variation of CA/CAo with time.

Po~= 1050Pa, CA0 = 5 x 10-SM (average values).

General mechanism of the ozonation of atrazine

The results shown in the previous section indicate that the ozonation of atrazine depending on the experimental conditions is carried out through direct reaction with ozone (i.e. at pH 2 and in the presence of t-butanol) and via radicals (hydroxyl radicals) a priori formed from the decomposition of ozone (Staehelin and Hoignr, 1985). On the other hand the presence of important amounts of dissolved ozone (especially at pH < 12) suggests undoubtedly that the chemical reactions develop in the bulk water (Danckwerts, 1970; Charpentier, 1981; Beltrfin et al., 1992). Under these circumstances and assuming that the ozonation of atrazine follows the well- known mechanism of Staehelin and Hoign6 (1985), atrazine is degraded by the following reactions and mechanism:

Direct reaction with ozone:

kd A -t- zO3-----* Interme&ates (1)

where z is the stoichiometric ratio and kd the rate constant of step (1), and by the action of hydroxyl radicals:

A + OH kOH,A, Intermediates kOH,A

=l.8101OM-Js I (2)

(Beltrfin et al., 1993a).

presence of t-butanol to check the importance of the radical contribution to the elimination of atrazine. The direct reaction of ozone and atrazine rep- resented, as shown later, a 24 and 81% of the elimination of atrazine in the absence and presence of t-butanol (0.01 M in concentration and 20C), respectively. Now, it can be observed (see Fig. 3) that an increase of temperature leads to an increase of the degradation rate of atrazine.

At first sight, hydroxyl radicals come from the decomposition of ozone through the following general mechanism (Staehelin and Hoignr, 1985) where the rate constant units are in M -1 s ~ or s-l:

Initiation step:

03 + OH- k~ ~ 70, HOr + 02 . (3)

Propagation steps:

1.0 ,.,

~'~ T *C PO31 Pa I \~ , - . . . . re , , _ I-X * ;~ '~ '~ o 20 1100 l '~ '~.~=. m, 10 840 I- \ \ " ~_~ff-,.~. Q 3 670

" ~ *" . . ' -2"* ' -

OO31' " 1100 Pae z~'~ ~. .o T,*C ~ ,,LU~ ,,~

2010 ~ - - .

t II i 31 = = n i a a I I I

0 , 8 12 18 20 2 ' , '~ '8 time, min

Fig. 3. Influence of temperature on the ozonation of atrazine. Variation of CA/CAo with time at pH 2, CA0 = 5 x 10-SM (in the presence of 0.01 M t-butanol at

different Po3i).

pK = 4 .8

He2.' , O~. + H + (4)

k I = 1.9 109

O~. + 03 , O;. + 02 (5)

k 2 = 5 x 1010

O~. + H + , HO3. (6)

k3 = 1.4 105

He 3. , OH. + 02 (7)

k4 = 2 x 109

OH. + 03 , HOt + 02. (8)

Termination steps [in addition to reaction (2)]:

ksi

OH. + Si , final products? (9)


being S~ a hydroxyl radical scavenger like t-butanol, and

OH. + Intermediates kp, final

or more intermediate products? (10)

In the above mechanism some features have been considered. First, ozone is a promotor of its own decomposition through reactions (5) and (8). Second, in the presence of scavengers of hydroxyl radicals, S, the termination step could be the reactions of hydroxyl radicals with atrazine and intermediates. Third, intermediates or ozonation byproducts can compete with atrazine for the ozone available through their direct reactions (not considered). And fourth, reactions between molecular compounds (atrazine and intermediates) are considerred as termination steps of the main radical mechanism (Glaze et al., 1992).

Determination of the rate constant of the direct reaction between atrazine and ozone, kd

According to this mechanism the rate of atrazine degradation has two contributions through reactions (1) (direct way) and (2) (radical way)

dCA kd CA Co 3 + kon.A CA Coll. (11)

dt z

where CA, Co3 and Co, represent the concentrations of atrazine, dissolved ozone and hydroxyl radicals, respectively. In equation (1 l) there is no volatilization term since atrazine is a non-volatile compound, at least for the short experimental reaction times of this work.

In order to determine the direct rate constant, ko, it is necessary to eliminate the contribution of hydroxyl radicals. In so doing, a series of experiments was performed at pH 2 in the presence of t-butanol to decrease the action of radicals. This compound reacts fast with hydroxyl radicals, the rate constant of this reaction is 7.3 108M-ls- I (Scholes and Willson, 1967), and practically does not react with ozone, the rate constant of its direct reaction with ozone has been reported to be lower than 3 10-3M I s - - I (Hoign6 and Bader, 1983a). A concentration of 0.05 M of t-butanol was applied to ensure that all hydroxyl radicals formed are inhibited by the scavenger, the initial atrazine concentration being 4.8 x l0 -5 M. Under these conditions, the rate of atrazine oxidation is

dCA =kdCACo ' (12) dt z

Another particular feature of the ozonation of atrazine, especially at pH 2, is that the dissolved ozone concentration reaches its stationary and maximum value in hardly three minutes of reaction (which is a strong indication of very slow gas-liquid reaction) as shown in Fig. 4. If equation (12) is applied to reaction times equal to or higher than

15

~; 10

o

x t

o :, 0 5

C03 s - 1.1 x 10 -4M

** ~ a * ~-~ . . . . . . - ! - - - -~ . - - _ 7_ -_ ~ .

T*C Po3 i , Pa

= 20 1100

* 10 840

o 3 695

I = I j , I I I I I I I I 1 I

0 4 8 12 16 20 24 28

t ime, rain

Fig. 4. Changes of dissolved ozone concentration with time during the ozonation of atrazine at pH 2 in the presence of

0.01 M t-butanol at different temperatures.

that necessary for ozone to reach its stationary concentration, Co3s, then the oxidation of atrazine follows a pseudo first order kinetics

where

For t t> ts dCa dt = k'd CA (13)

k~ = k~ co. (14) Z

Equation (13) can be integrated under the following boundary conditions

t=ts CA=Cas (15)

t=t CA=CA (16)

CAs being the concentration of atrazine when the ozone concentration reaches its stationary value, Co3 s. After integration, the resulting equation is

In CA = -- k'd(t -- ts) (17)

According to equation (17) a plot of ln(CA/CAs ) vs ( t - ts) should yield a straight line whose slope is

-0.1

~'~ -0.2

* 3 695 "~ -0 .3 10 840

o 20 1100 -.,,,.. i i i i i i i B , = i , , ,

0 4 8 12 16 20 24 28

t - t s, rain

Fig. 5. Verification of equation (17). Variation of ln(CA/CAo ) with (t - t~) at different temperature, pH 2 and in the presence of 0.05 M t-butanol. CA0 = 4.8 x l0 -5 M.

2158 FERNANDO J. BELTR~.N et al.

-k~. Figure 5 shows this plot corresponding to the experiments carried out at different temperature (3, 10 and 20C), pH 2 and in the presence of 0.05 M of t-butanol. It can be observed that points are situated around straight lines whose slopes are negative and increase (in absolute value) with the increasing temperature. By taking into account the values of Co3s and z [the latter was found to be 1, calculated following a procedure already published (Sotelo et al., 1990)], the rate constant of the direct reaction of atrazine and ozone was determined to be 4.5, 3.5 and 2.8 M -j s -~ at 20, 10 and 3C, respectively. From these values the following Arrhenius function was obtained

kd=l.01 x 104exp(--18.8/RT), M- is i (18)

where the activation energy is in kJ.mol -~. Since atrazine is a dissociating compound there are

in fact two direct reactions with ozone to consider; those with its protonated and unprotonated forms which are in equilibrium in water

AH + ~ A + H +. (19)

Therefore, k d is expressed as a combination of the rate constants of these reactions as follows

kd = kAH+ Ct + kA(1 -- Ct) (20)

where ct is the degree of protonation which is a function of pH and pKa of equilibrium (19)

10-Prl = 10_PH + 10_P~ a. (21)

However, Hoign6 and Bader (1983b) have reported that protonated amino compounds do not react directly with ozone, so that equation (20) simplifies to

kd = ka(1 - ~). (22)

At pH 2 the degree of protonation of atrazine is 0.28, pKa = 1.6 as reported by Yao and Haag (1991), and from equations (18) and (22), the actual rate constant of the direct reaction between atrazine and ozone at 20C is 6.3 M- i s ~. In fact, for pH >/4,

~ 0 and k d ~ kg, which is in accordance with the kd value reported by Yao and Haag (1991) at pH 4.1 (see Table 1). Comparing with other literature data, the value of k d obtained in this work is something higher than that reported by Xiong and Graham (1992) which was obtained from competitive homogeneous ozonation of atrazine and MCPA as reference compound. Deviations could probably be due to the fact that the latter depends on the rate constant value of the reaction between ozone and the reference compound. On the other hand, Xiong and Graham (1992) also reports kd to be 12.2 M -~ s -~ at pH 7.5 obtained in the presence of 10-4M of carbonate ion, a concentration which seems to be too low to inhibit all hydroxyl radical formed at this pH. Finally, the extremely high values reported by Adams et al. (1990) and Adams and Randtke (1992) at any pH in the

absence of scavengers are undoubtedly due to the contributions of radical reactions which were not inhibited during their experiments.

Modelling the ozonation of atrazine in water

From the mechanism formed by reactions (1)-(10) combined with mol balance equations of atrazine and ozone (for the latter applied to the water and gas phases) the concentration of these species can be predicted for any set of experimental conditions. Thus, the kinetic model equations are:

Mol balance of atrazine:

dCA =kdCACo3+koHACACoH. (23) dt

Mol balance of ozone in water (assuming that gas and liquid phases are well mixed in the reactor):

dCo3 dt = KGa(Co3gRT - HCo3) - ,~ri (24)

where KGa, Co3g, R, H and T are the overall gas phase mass transfer coefficient, referred to the liquid volume, the concentration of ozone in the gas leaving the reactor, the gas and Henry law constants and temperature, respectively, and/~ri the decomposition term of ozone due to chemical reactions, defined, neglecting the contribution of its direct reactions with intermediates, as follows

T, ri = (kdCA + ki 10 (on - 14) .~_ kl Co2. + k4Cou. )Co3. (25)

Mol balance of ozone in the gas phase:

dC3~= v V s \RT Co3g) - KGa' (Co,~RT-HCo3)

(26)

where v0, Po3i and V G are the gas volumetric flow rate, ozone partial pressure at the reactor inlet and the accumulated volume of gas within the reactor (50 cm3), respectively, and KGa' is the overall mass transfer coefficient referred to the gas volume which is related to KGa as follows

vo KGa = -~L K~a" (27)

V e being the liquid volume in reactor (800 cm3). It has to be noticed that mol balance equations of

intermediates should be included in the kinetic model since these compounds also consume hydroxyl radicals and ozone. However, as atrazine and intermediates have similar nature (Legube et al., 1987; Adams and Randtke, 1992) their direct reactions will be neglected because of their expected low contributions to consume ozone. Consumption of hydroxyl radical by intermediates, on the other hand, will be accounted for assuming a constant hydroxyl radical scavenger contribution for atrazine and intermediates during the reaction period as explained later. As


a consequence, tool balance equations for intermediates were not considered.

COH'

and

Determination of radical concentrations

Applying the hypothesis of the stationary state to the net formation rates of radicals in the proposed mechanism, the hydroxyl and superoxide ion radical concentrations are finally expressed by the following equation

2k i 10~PH- 14)Co 3 (28) kOH.A CA + ~kpi Cpi + Zk~i C~i

2k i 10(PH - 141 _[._ k4 Con. Co~. - kl (29)

Determination of the kinetic regime of absorption

It should be first noticed that the system formed by equations (23)-(25) only holds if the reactions through which ozone decomposes are slow relative to the ozone mass transfer rate (Charpentier, 1981). This fact can be checked by calculating the Hatta number of these reactions, which relates the chemical reaction and physical absorption rates within the diffusive film in the proximity of the gas-water interface. For an irreversible m,n order gas-liquid reaction, the Hatta number is defined as follows

Ha = (30) kL

where m and n are the reaction orders with respect to the gas being dissolved (ozone in our case) and a compound or species B already present in the liquid phase (atrazine, hydroxyl ion or a radical in our case), k is the rate constant of the reaction considered, Do3 the diffusivity of the dissolved gas in the liquid and k L is the individual liquid phase mass transfer coefficient.

The decomposition of ozone develops through four different reactions [see equation (25)]. However, the contributions of these reactions in equation (25), after considering the stationary state hypothesis for hydroxyl and superoxide ion radicals [see equations (28) and (29)], eventually become

~r i = [kdC A -4- 3k i 10 (pH - 14) l

4k4ki 10ton - 14)C3 I f - - - - - - 03 + ko..ACg + ZkpiCoi + Sks, C~iJ (31)

Equation (31) indicates that ozone decomposes through three hypothetical reactions: an 1,1 order direct reaction with atrazine (direct reactions between ozone and intermediates have been neglected), a pseudo first order hydroxyl ion catalysed reaction and, finally, a pseudo second order radical reaction,

Table 3. Hatta numbers of ozone decomposition reactions in the ozonation of atrazine a

Reaction b pH 2 pH 7 pH 12 Direct 5.410 -3 6.4x10 3 6.4xl0 3 Hydroxyl ion

initiation 5 x 10 -6 1.6 x 10 3 5.2 x 10 Free radical 2.4 x 10 6 7.1 x 10 4 2.1 x 10

aHatta numbers calculated from equation (30). Conditions: 20C, CA0 = 5 x 10 ~ M; kL = 10 4 ms- ~. bContribution of terms of equation (31).

whose rate constants are, respectively, kd, 3k~ 10 ~pH - 14) and

4k4k i 10(PH - 14)

koH. A CA + 2;kpi Cpi + Sksi Csi

= (klCo2. +k4Co,.)/Co~ (32)

Applying equation (30) to these reactions their Hatta numbers were calculated at different pH levels. First, k L was taken to be 10 -4 ms - l , which is a typical value for the type of contactor used in this work, and Do3 was obtained from bibliography (Matrozov et al., 1976). Table 3 presents the results. It can be observed that in all of the cases but one (at pH 12) Ha is well below 0.3, condition for the reactions to be slow, and hence the proposed mol equation system is appli- cable. At pH 12, however, the hydroxyl ion catalysed decomposition reaction of ozone (initiation step in the mechanism) is a moderate reaction, that is, part of the dissolved ozone decomposes in the film layer. As a consequence the kinetic model does not hold at pH 12 and we have focused the attention on the ozonation at pH 2 and 7.

Determination of K6a

In order to solve equations (23)-(26), it is necessary to first calculate the overall mass transfer coefficient, KGa, and express the concentration of hydroxyl and superoxide ion radicals as a functions of the concentrations of stable species, such as atrazine, ozone and scavengers. In an ozonation experiment when the concentrations of ozone in the water and in the gas leaving the reactor have reached their stationary values, Co3s and Co3gs, respectively, K~a' can be obtained from equation (26) where the ozone ac- cumulation term is now zero due to the steady state condition

t~ FPO3i ] KGa 0 [_~-~ -- Co3~

' = (33) VG ( Co, g~ R T - H Co3s ) "

Table 4. Rate constants of reactions between the hydroxyl radical and scavengers, k, used in this work

ksa Species (M - ) s i ) Reference

t-Butanol 7.3 x l0 s Scholes and Willson (1967) Carbonate ion 4.2 x l0 s Weeks and Rabani (1966) Bicarbonate ion 1.5 x l0 T Weeks and Rabani (1966) Phosphoric acid 2 x 106 Grabner et al. (1973) Diacid phosphate ion 2.2 x 106 Grabner et al. (1973) Monoacid phosphate ion 8 105 Grabner et al. 0973)

2160

Table 5. Calculated ~ and experimental values of COB and ?o~

CAo x 10 5 C~ Experimental Calculated a (M) pH (M) cOH b ~'O3 COB 703

FERNANDO J. BELTR./~N et al.

4.84 2 - - 4.4 x 10-14 24 1.6 x 10-2 100 4.24 2 0.01 d 5.9 x 10 15 81 1.9 x 10 2t 100 4.80 2 0.05 d 3.5 x l0 n 100 4.1 x l0 -22 100 5.31 7 - - 1.4 x 10-13 18 1.3 x 10-15 96 4.76 7 0.075 c 4 .3x10-~4 37 5x 10-16 98 5.07f 7 - - 4.6 x 10- ~4 35 6.6 x 10 - ~6 96 4.98 7 - - 1.2 x l0 -13 23 1.3 x l0 -15 96 5.34 g 7 - - 6.2 x 10 is 10 2.5 x 10 -15 96

~Obtained from the basic mechanism of reactions (1)-(10); average value of Po3~ is 1050 Pa unless indicated, bFrom equation (41). CFrom equation (39). at-Butanol concentration. ~rotal bicarbonate ion concentration, rPo3~ = 500 Pa. gPo3~ = 2000 Pa.

The application of equation (33) to the ozonation experiments yielded an average value for KGa' of (4.7 _ 2.1) 10-3 mol - atm -~ (1 of gas) -I s -~. Now, from equation (27) KGa was found to be (2.9_+ 1.3) x 10 4M-atm- ls - l .

Prediction of concentrations of atrazine and ozone

At this stage, the first order differential equations (23), (24) and (26) can be solved numerically to obtain the concentration profiles of atrazine and ozone. In this work, the fourth order Runge-Kutta method was applied, solving the system with the initial conditions

t=0 CA=C ~ Co~=0 and Co3g=0. (34)

Due to the unknown composition of the ozonated atrazine solution at any time, in the denominator of equation (28) the scavenger term of atrazine and intermediates [reactions (2) and (10)] was assumed to be constant and expressed as

kOH, AC A q- ,~kpiCpi = kOH, A CA0. (35)

This is a reasonable hypothesis if the following circumstances are considered or assumed:

1. The rate constants, kOH, A and kpi, of hydroxyl radical reactions can be approximately equal because the unselective character of hydroxyl radicals [the rate constants of these reactions are of similar magni- tude, Farhataziz and Ross (1977)] and the intermediates formed during the ozonation of atrazine, as reported by Legube et aL (1987) and Adams and Randtke (1992), are also similar in nature to their parent compound, atrazine.

2. Total mol balance of atrazine and intermediates holds at any time during the reaction period considered in these experiments. It has to be noticed that Adams and Randtke (1992) reports that mineralization is negligible which supports the above comment.

For solving the system, the values of ksi applied are given in Table 5 and the Henry constant was obtained from a previous work (Sotelo et al., 1989). Under these conditions and hypothesis, however, the model overestimated the experimental concentrations except when a high concentration of scavenger was present.

For this case, Fig. 6 presents the variation of calculated and experimental concentrations of atrazine and ozone in water and in the exiting gas for the latter versus reaction time. This suggests that the model works well when hydroxyl radical reactions are sup- pressed. The results obtained at other conditions, however, can be surprising at first sight, since mol balances of intermediates were not included. Under these circumstances, it should be expected the calculated concentrations of atrazine to be lower than the actual values because intermediates also consume ozone and hydroxyl radicals and their contribution would tend to stabilize the concentration of atrazine. It could seem that the hypothesis given by equation (35) is the reason of the high values calculated for the concentrations of atrazine. However, solution of differential equations (23), (24) and (26), considering only the role of atrazine as hydroxyl radical scavenger, that is, kOH, AC A instead of kOH, ACAo in equation (35), leads to high calculated concentrations as well. There can be different explanations to this behaviour. Thus, atrazine and some ozonated by- products could be promoters of hydroxyl radicals if their reaction with these radicals eventually release the superoxide ion radical, as reported by Staehelin and Hoign6 (1985) for some compounds. However, the promoting character of intermediates have also been considered when equation (35) was not applied. Furthermore, if atrazine is also a promotor of hydroxyl radicals, in the absence of scavengers, the only possible termination step would be the reaction between hydroxyl radicals (Staehelin et al., 1984)

kt=Sxl09 M I s I OH" + OH" , H,O2 (36)

Note that reaction between ozone and hydroxyl radical is not a termination step but a propagation reaction. According to all this, the concentration of hydroxyl radicals could be given by equation (37)

[2k~lO C~"- ~4~C 705 Co. = L kt 3 / (37)

However, for an average concentration of ozone of 10-SM and pH 7 equation (37) leads to a concentration of hydroxyl radicals of 1.7 x 10 -j M which is too high compared to the experimental values obtained as shown later.

Another possible explanation to the differences between predicted and experimental concentrations of atrazine can be due to the presence of traces of impurities in the starting material that, acting as initiators, could have promoted the decomposition of ozone into radicals, a role that has also been attributed to some substances (Staehelin and Hoign6, 1985). In this case, the following reaction should be considered

hi I + O3---+ O3. -+ OH"

While solving the simultaneous differential equations of the kinetic model )'o3 was obtained to be 100% at pH 2 and 96% at pH 7. These results together with the predicted concentration of hydroxyl radicals obtained from equation (28) are given in Table 5.

On the other hand, the experimental values of Vo3 and Con. were obtained as follows. If it is assumed a constant concentration of dissolved ozone, Co3m, for the whole reaction period (which is true after a few minutes of the start of ozonation) and the stationary hypothesis for the hydroxyl radical holds (thus, Con. is constant), the rate of disappearance of atrazine is

so that the concentration of hydroxyl radicals becomes

(2ki 10(PH- 14) + kl Cl)Co3 COH" ~---

koH, A CA, + ~kp~ Cpi + Y,k~ C~i

(2k~ 10 CpH - 14)C03 + # (38)

koH, A CAo + Z'ksi Csi

where the term fl is a hydroxyl radical initiating rate factor that would involve other contributions to the formation of hydroxyl radicals apart from the initiation step [reaction (3)]. This term would depend on the experimental conditions and composition of the starting aqueous solution.

Direct and radical contributions to the oxidation of atrazine

An important feature of this process is to know the contribution of the direct and radical ways to the elimination of atrazine that, on the other hand, is related to the initiating rate factor, ft. Thus, the percentage elimination of atrazine through its direct reaction with ozone is

kdC3 x 100. (39) ~o3 = kd Co 3 + kor~, A Con.

dCA dt - (kdC3m + kOH'A CoH')CA = kTCA (40)

Integration of equation (40) leads to the determination of k v and hence to the experimental hydroxyl radical concentration

kT -- kd Co3m (41) CoHexp = koH, A

1.0~

I: T B u o

b o

0.5

o t


1(~

: f

' ~ ' h ' t~ ' lh ' ~o ' ~4 ' 18 ' o , 8 ,2 16 2o 24 28

t ime, ra in t l r r~ , m~

(a) (b)

0 2

o o o

o

o l Ca lcu la ted data

O~ i i i i i , , I , , I i I I I

4 8 12 16 20 24 28

t ime, ra in

(c)

Fig. 6. Checking of the ozonation kinetic model. Variation of experimental and calculated concentrations with time: (a) atrazine as CA/CAo; (b) dissolved ozone; (c) ozone in the gas. Conditions: pH 2,

CA0 = 5 X 10 -5 M, 0.05 M t-butanol, Po~ = 1100 Pa.

2162 FERNANDO J. BELTRAN et al.

Table 6. Values of the hydroxyl radical initiating rate factor, #, at different experimental conditions

Po~, CA0 10s C, #~ (Pa) pH (M) (M) (M s - i )

1020 2 4.8 - - 4.05 x 10 - s 1100 2 4.2 0.01 b 4.8 x 10 - I I100 2 4.8 0.05 b 1.3 x 10 - 1020 7 5.3 - - 1.32 10 _7 1020 7 4.8 0.075 7.75 x 10 -~ 500 7 5.1 - - 4.00 x 10 - s

1000 7 5.0 - - 1.06 x 10 -7 2000 7 5.3 - - 5.50 x 10 -7

~Calculated from equation (42). bt-Butanol concentration. ~l'otal bicarbonate ion concentration.

Table 5 also shows the experimental values of Coll. calculated from equation (41). Comparing with the predicted values one can observe important differences. Also, shown in Table 5 are the experimental values of Yo3 obtained from equations (41) and (39). These results show the importance of radical reactions even at pH 2 in the absence of scavengers. According to them, in addition to reaction (3), there have to be other routes to form hydroxyl radicals. These can be due to the promoting character of some intermediates and/or to the initiating effect

of some impurities at trace concentration. In fact, both factors are accounted for in the term fl that can be calculated once Co.. is known, by the following equation

fl = (koH, ACA0 + Sks iCs i )Cou . - 2ki 10


u

1.0

~ pH 7 u 0.075

4 8 12 16 20 24 28

x t' z

0

o

o

I w,,.T_ u' u o D

a o

Bf~rbonoto ion

F7 C~. M I mm pH 7 0

| a 0.07S

0 4 8 12 16 20 24 28

IO

3"

Calculeted data with sottvenger + time, rain

(a)

0 O 0

Bicarbonate Ion C=, M

0 pH 7

u 0.075

4 12 16 0 24 28 time. rain

(e)

time. rain (b)

Fig. 8. Checking of the ozonation kinetic model with the inclusion of ft. Variation of experimental and calculated concentrations with time at pH 7: (a) atrazine as CA/CAo; (b) dissolved ozone; (c) ozone in the

gas. Po3i = 1020 Pa, CA0 = 5.3 x l0 -5 M.

deviations lower than 15% except those of dissolved ozone in the absence of scavengers. Deviations between experimental and calculated values of dissolved ozone concentration could be due to the consumption of ozone in direct reactions with some intermediates that were not accounted for. These reactions do not likely develop in the presence of scavengers since the nature of intermediates can be different (Legube et al., 1987; Adams and Randtke, 1992).

CONCLUSIONS

The ozonation of atrazine is a process that at neutral pH develops mainly through radical reactions. The total rate constant of the direct reaction with ozone is very low (6.3 M -l s -I at 20C) so that even at pH 2 the contribution of the radical way can be important. In addition, the process is highly dependent on the presence of impurities that pro- motes the decomposition of ozone and/or the promoting character of some intermediates able to release superoxide ion radicals and hence hydroxyl radicals. The lack of knowledge of these contributions can be grouped in a term called the hydroxyl radical initiating rate factor. Inclusion of this factor in a kinetic model formed by mol balance equations

of atrazine and ozone, once the decomposition of the latter due to chemical reactions is obtained from applying the mechanism of Staehelin and Hoign6 (1985), considerably improves the concordance between predicted and experimental concentrations especially when the presence of a strong hydroxyl radical scavenger is considered in the model. The more important discrepancies are observed when comparing the experimental and calculated concentrations of dissolved ozone in the absence of scavengers. This could be attributed to the appearance of some reactive intermediates towards ozone. As a consequence further studies to elucidate these reactions and competitivity of intermediates especially for the ozone available are needed to complete the full mechanism of the ozonation of atrazine.

It should be noticed that in a real situation with a surface water, contaminated by atrazine or a given pollutant, carbonate/bicarbonate ions would be likely present, possibly at appreciable concentration (> 10 -4 M). Since the full composition of the aqueous ozonated solution is not possible to know and ozonation is going to be applied, the model used in this work could be adopted. Thus, once the initiating rate factor is calculated as shown in this work and intro- duced in the kinetic model, the concentration of the

WR 28/1(~-I

2164 FERNANDO J. BELTK.AN et al.

remaining atrazine or a given pollutant, for an ozone dose applied, can be predicted.

Acknowledgement--Authors thank the CICYT of Spain for its support through grant AMB93/654.

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Advanced Oxidation of Atrazine in Water—I. Ozonation

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