Ozone-Enhanced Oxidation of Oxalic Acid in Water with Cobalt

Jun 4, 2003 - Anaid Cano Quiroz , Carlos Barrera-Díaz , Gabriela Roa-Morales , Patricia Balderas Hernández , Rubí Romero , and Reyna Natividad...
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Ind. Eng. Chem. Res. 2003, 42, 3210-3217

Ozone-Enhanced Oxidation of Oxalic Acid in Water with Cobalt Catalysts. 1. Homogeneous Catalytic Ozonation Fernando J. Beltra´ n,* Francisco J. Rivas, and Ramo´ n Montero-de-Espinosa Departamento de Ingenierı´a Quı´mica y Energe´ tica, Universidad de Extremadura, 06071 Badajoz, Spain

Oxalic acid in water has been treated with ozone in the presence of a Co(II) salt at acidic pH. The influence of different variables, including the initial oxalic acid, cobalt(II), and ozone gas concentrations and the temperature, has been investigated. The experimental stoichiometric ratio varied between 0.7 and 1.4 mol of ozone consumed per mole of oxalic acid consumed, while the ozone efficiency reached values as high as 25%. At any conditions applied, nearly total mineralization was achieved. The process developed between the slow and moderate kinetic regimes of ozone absorption. Experiments in the slow kinetic regime allowed for the experimental determination of the reaction kinetics, which was found to be first order with respect to oxalic acid and catalyst. For the case of ozone, nearly one-half-order kinetics was found. A mechanism that involves the participation of cobalt-oxalate complexes is also discussed. 1. Introduction One of the main problems of ozonation in the treatment of wastewater is the appearance and accumulation of refractory compounds, specifically, carboxylic acids, aldehydes, ketones, etc., that interfere with mineralization of the organic matter present in water. This problem is partially overcome by the use of advanced oxidation processes, which implies the combination of oxidant agents such as hydrogen peroxide (in addition to ozone); radiation (UV or visible), and/or catalysts such as Fe(II) (i.e., Fenton reagent); TiO2 (photocatalytic processes), etc.1-3 These processes are based on the generation of hydroxyl radicals that unselectively attack the organic matter in the solution.4 For ozonation, however, another possible way to improve the removal rate of refractory contaminants is the use of metal salts as catalysts. In fact, it is well-known that the decomposition of ozone can be initiated by metal catalysts dissolved in water.5 Thus, the literature provides examples of water and wastewater ozonation in the presence of metal salts of different nature.6,7 Nonetheless, the toxic character of metals has stopped the application of this kind of process in the water treatment field. On the contrary, the use of solid catalysts can be a useful alternative to both take advantage of the advantages of catalytic ozonation and avoid the problem of toxic metals in water.8 Also, in the treatment of some wastewater, catalytic ozonation, even when it is carried out homogeneously, can provide an increase in the TOC and COD removal rates of the wastewater.9 In this work, a cobalt catalyst has been used to improve the removal of oxalic acid, a refractory compound, through ozonation in water. Oxalic acid was chosen because it does not react directly with ozone and because it is one of the end products that accumulate in water during the ozonation of phenols and other aromatic compounds.10 The work was carried out at pH 2.5 because, under these conditions, the formation of * To whom correspondence should be addressed. Tel.: 34924-289387. Fax: 34-924-271304. E-mail: [email protected].

hydroxyl radicals from the hydroxyl-ion-catalyzed decomposition of ozone is negligible.5 The work is divided into two parts separately considering homogeneous and heterogeneous catalysis. In this first part, cobalt(II) is used to check the catalytic action of this metal in the ozonation process. Cobalt ions are reported to decompose ozone into hydroxyl radicals, and as a consequence, their application with ozone could be considered a potential advanced oxidation approach.11,12 Accordingly, cobalt ions could be use to increase the removal of oxalic acid or other refractory compounds from water. In addition, the literature also reports that Co(II) forms complexes with oxalic acid that can react with ozone at appreciable rates.13 There are a few works dealing with the use of Co(II) in water during ozonation processes. To the best of our knowledge, only three works on this subject have been reported so far. In two of them, Co(II), among other soluble metal catalysts, was used to treat an industrial wastewater and a river water with ozone.6,7 The third work appeared recently in the literature.13 Similarly to this paper, it considers the treatment of oxalic acid with ozone-Co(II), but in it, the pH is between 5.3 and 6.7 and the process is carried out by mixing aqueous solutions of ozone and oxalic acid. That work13 explains the experimental results on the basis of direct reactions between ozone and Co-oxalic acid complexes without the participation of free radicals. The lack of more information about this type of reaction is undoubtedly due to the potentially toxic character of cobalt ions that prevents their use in water. However, it should be highlighted that the aim of the study presented here is not to confirm cobalt ions as useful homogeneous catalysts in water ozonation processes but to check the magnitude of its catalytic action to further prepare solid cobalt-supported catalysts, as is shown in the second part of this work.14 Also, it is the objective of this part to determine a kinetic rate expression for the homogeneous catalytic ozonation process to establish any possible contribution of this catalysis on the heterogeneous process presented in the second part of this work.14

10.1021/ie0209982 CCC: $25.00 © 2003 American Chemical Society Published on Web 06/04/2003

Ind. Eng. Chem. Res., Vol. 42, No. 14, 2003 3211

Figure 1. Evolution of the dimensionless concentration of oxalic acid with time for catalytic ozonation experiments at different initial oxalic acid concentrations. Conditions: gas flow rate ) 24 L h-1, agitation speed ) 300 rpm, T ) 20 °C, pH ) 2.5, catalyst concentration ) 0.8 mg L-1, ozone gas concentration ) 30 mg L-1, initial oxalic acid concentration (× 10-3 M) ) (2) 8, ([) 4, (9) 1.

2. Experimental Part Oxalic acid was obtained from Aldrich and used as received. Ozone was produced from pure oxygen in a laboratory ozonator. Co(NO3)2 (Merck) was used as the catalyst. Experiments were carried out in a 1000-mL jacketed glass vessel provided with an agitation bar. For each run, 800 mL of an oxalic acid aqueous solution was charged into the reactor. The temperature inside the reactor was kept constant by circulating water from a thermostatic batch in which the reactor was submerged. The oxalic acid concentration was determined by liquid chromatography as is presented in detail in a previous publication.15 Ozone in the gas phase was analyzed by means of a GM109 Anseros ozomat analyzer based on the absorption of ozone at 254 nm. The concentration of dissolved ozone was determined by the indigo method.16 The total carbon (TC) dissolved in water was also followed with a DC-190 Dorhmann TOC analyzer, which measures the carbon dioxide infrared absorption. Because the experiments were carried out at pH 2.5, the TC values correspond, in fact, to the remaining dissolved organic carbon (TOC). 3. Results and Discussion Experiments on the ozonation of oxalic acid at pH 2.5 were carried out at different conditions of oxalic acid concentration [(1-8) × 10-3 M], ozone gas concentration (7-45 mg L-1), metal ion concentration (0.4-1.6 mg L-1), and temperature (10-40 °C). In all cases, the gas flow rate and agitation speed were kept constant at 24 L h-1 and 300 rpm, respectively. Influence of Variables. The effects of the four variables studied are shown in Figures 1-4 where the evolution of the remaining dimensionless concentration of oxalic acid with time under different conditions has been plotted. In Figure 1, the influence of the initial oxalic acid concentration is shown. As can be seen from this figure, the higher the initial oxalic acid concentration, the higher the time needed to reach a given oxalic acid conversion. It is also seen that the total removal of oxalic acid can be achieved in about 75 min when the initial concentration of oxalic acid is 10-3 M. In Figure 2, the effect of ozone concentration in the gas fed to the reactor is presented. In this case, for the higher ozone concentration applied (45 mg L-1), a nearly 50% oxalic acid conversion is reached in 90 min starting with 8 × 10-3 M of oxalic acid. The positive effect of this variable is a

Figure 2. Evolution of the dimensionless concentration of oxalic acid with time for catalytic ozonation experiments at different ozone gas concentrations. Conditions: gas flow rate ) 24 L h-1, agitation speed ) 300 rpm, T ) 20 °C, pH ) 2.5, catalyst concentration ) 0.8 mg L-1, initial oxalic acid concentration ) 8 × 10-3 M, ozone gas concentration (mg L-1) ) (9) 7, ([) 15, (2) 30, (b) 45, (4) 30 with 0.01 M tert-butyl alcohol.

Figure 3. Evolution of the dimensionless concentration of oxalic acid with time for catalytic ozonation experiments at different catalyst concentrations. Conditions: gas flow rate ) 24 L h-1, agitation speed ) 300 rpm, T ) 20 °C, pH ) 2.5, ozone gas concentration ) 30 mg L-1, initial oxalic acid concentration ) 8 × 10-3 M, catalyst concentration (mg L-1) ) (*) 0.2, (9) 0.4, ([) 0.8, (2) 1.2, (b) 1.6.

logical consequence of the increase in the ozone driving force from the gas to the water as a result of the increase of ozone partial pressure. Figure 3 presents the influence of catalyst concentration on the oxalic acid removal rate. Again, the effect is positive, with a decrease in the concentration of oxalic acid remaining after reaction with increasing catalyst concentration applied. Under the conditions investigated, the presence of 1.6 mg L-1 of Co(II) in solution allows about 70% removal of oxalic acid in 90 min. Finally, the effect of temperature on the Co(II) homogeneous catalytic ozonation of oxalic acid is shown in Figure 4. As can be seen, there is a definitively positive influence of temperature, with oxalic acid conversions of between about 20 and 65% in 90 min at temperatures between 10 and 40 °C, respectively. From Figures 1-4, it is evident that Co(II) is an active homogeneous catalyst of the ozonation of oxalic acid. Oxalic Acid Mineralization. Mineralization is an important aspect to study in this kind of process. The experimental results also showed high levels of mineralization achieved. The mineralization process was nearly quantitative regarding the oxalic acid consumption because the remaining TOC calculated from the oxalic acid concentration at any time was slightly lower than that found with the TOC analyzer, as shown in Figure 5. Particularly at advanced reaction times, the difference between the two TOC values reaches approximately 10%. In any case, Co(II) seems to be an appropriate catalyst for the mineralization of oxalic acid in an aqueous ozonation process. It should be noted that

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Figure 4. Evolution of the dimensionless concentration of oxalic acid with time for catalytic ozonation experiments at different temperatures. Conditions: gas flow rate ) 24 L h-1, agitation speed ) 300 rpm, pH ) 2.5, catalyst concentration ) 0.8 mg L-1, initial oxalic acid concentration ) 8 × 10-3 M, ozone gas concentration ) 30 mg L-1, T (°C) ) (9) 10, ([) 20, (2) 30, (b) 40.

where vg is the gas flow rate; V is the reaction volume; CB0 and CB are the molar concentrations of oxalic acid at the start of ozonation and after a given period of time t, respectively. The experimental results show that, at a given reaction time (i.e., 60 min), the ozone efficiency slightly increases with increasing ozone concentration in the gas fed to the reactor, increasing mass of catalyst used, and increasing temperature applied, reaching values ranging between 7 and 25%. The overall stoichiometry of the ozone-oxalic acid reaction presented values varying between 0.7 and 1.4 mol of ozone consumed per mole of oxalic acid consumed. According to these results, the actual value of the stoichiometric ratio was likely close to 1. This value, as will be shown later, agrees with the mechanism proposed. Although the TOC analysis suggests the presence of a low fraction of intermediates (see Figure 5), the resulting stoichiometric reaction of the homogeneous oxalic acid ozonation for total mineralization at pH 2.5 could be

HC2O4- + O3 f 2CO2 + OH- + O2

Figure 5. TOC from carbon analyzer (TOC/TOC0a, vertical axis) versus TOC from remaining oxalic acid concentration (TOC/TOC0b, horizontal axis).

nearly total mineralization was also found in the presence of a powered suspension of activated carbon17 and TiO2.15 This supports the fact that catalytic ozonation of oxalic acid yields carbon dioxide and water with a low fraction of intermediates formed. Ozone Efficiency and Stoichiometry. The ozone efficiency and stoichiometry of the process were also determined. The ozone efficiency is defined as the ratio between the amount of ozone consumed and the total amount of ozone fed for a given period of time. Because of the changing value of the ozone concentration at the reactor outlet during the first few minutes of ozonation, the ozone efficiency was calculated using eq 1

(

EffO3 ) 1 -

)

∫0tCO g dt 3

CO3git

× 100

(1)

where CO3gi and CO3g are the molar concentrations of ozone in the gas at the reactor inlet and outlet, respectively, at a given time t. Notice that the numerator in eq 1, which rigorously represents the amount of ozone absorbed in water during a period of time t, is, in fact, the ozone consumed by reactions because the accumulated ozone was found to be negligible. On the other hand, the stoichiometric ratio z is defined as the amount of ozone consumed per unit of oxalic acid consumed during a period of time t. For similar reasons as before, z was calculated as follows

z)

vg(CO3git -

∫0tCO g dt) 3

V(CB0 - CB)

(2)

(3)

Kinetic Study. The main objective of this work is to gain information on the reaction rate of oxalic acid ozonation to be applied later in the heterogeneous catalytic ozonation, as discussed in the second part of this work.14 The first step in this sense is the establishment of the kinetic regime of ozone absorption. Kinetic Regime of Ozonation. In the absence of rigorous reaction rate data, the kinetic regime of ozonation can be established through a determination of the reaction factor, E, defined by

E)

NO 3 kLaCO3*

(4)

where NO3, kLa, and CO3* represent the actual absorption rate of ozone, the volumetric mass-transfer coefficient, and the ozone concentration at the gas-water interface, respectively. Notice that, in eq 4, the reaction factor is defined as a function of the maximum physical ozone absorption rate (denominator of the equation). Therefore, for slow and even for some moderate reactions, E takes values less than unity.18 When ozone is present in the water, the actual absorption rate can be written as follows

NO3 ) kLa(CO3* - CO3)

(5)

where CO3 represents the dissolved ozone concentration. Substitution of eq 5 into eq 4 leads to

E)

CO3* - CO3 CO3*

(6)

Equation 6 allows the reaction factor be determined from experimental data. Because perfect mixing conditions are assumed for both the gas and the water phases in the reactor used, CO3* can be calculated from the experimental ozone partial pressure in the gas at the reactor outlet, PO3s, once Henry’s law has been taken into account

CO3* )

PO3s He

(7)

Ind. Eng. Chem. Res., Vol. 42, No. 14, 2003 3213 Table 1. Experimental Dissolved Ozone Concentration, Reaction Factor, and Apparent Hatta Numbers in the Co(II) Homogeneous Catalytic Ozonation of Oxalic Acid in Watera CB0 CO3gi CCo T CO3 (×103 M) (mg L-1) (mg L-1) (°C) (×105 M) Eexpb Hac 1 4 8 8 8 8 8 8 8 8 8 8

30 30 30 7 15 45 30 30 30 30 30 30

0.8 0.8 0.8 0.8 0.8 0.8 0.4 1.2 1.6 0.8 0.8 0.8

20 20 20 20 20 20 20 20 20 10 30 40

13.2 8.2 2.6 0.7 2.8 11.4 6.6 0.9 1.1 11.4 3.6 1.0

0.12 0.35 0.75 0.80 0.58 0.41 0.47 0.93 0.91 0.20 0.89 0.89

0.03 0.06 0.14 0.16 0.10 0.07 0.01 0.26 0.24 0.04 0.22 0.22

kinetic regime slow slow slow slow slow slow slow moderate moderate slow slow slow

a Other experimental conditions: gas flow rate ) 24 L h-1, agitation speed ) 300 rpm. b Determined using eq 6 with CO3* values calculated from the literature19 and eq 7. c Determined using eq 8.

where He is the Henry constant.19 Values of the reaction factor were then calculated from eqs at the experimental conditions investigated in this work (see Table 1). It can be seen that E is especially close to unity in experiments carried out at catalyst concentrations higher than 1 mg L-1, which is a clear indication of a moderate reaction.18,20 However, this conclusion should be confirmed through a determination of the Hatta number. The Hatta number indicates the relative importance of the chemical reaction and mass-transfer steps in a gasliquid reaction.20 For reaction factors less than unity, an apparent Hatta number can be calculated by applying film theory concepts,18 to give

E)

Ha 1 cosh Ha sinh Ha cosh Ha + HaD sinh Ha

(

) (8)

where D is a dimensionless number representing the relative importance of the liquid volume and the film liquid volume, defined as

β D) aδL

(9)

with β, a, and δL defined as the liquid hold-up, the specific interfacial area, and the film length, respectively. The liquid hold-up was determined to be 0.95, while the specific interfacial area, a, was determined as the ratio between the volumetric and individual mass-transfer coefficients, i.e., kLa/kL. The volumetric mass-transfer coefficient was determined, as reported in the literature,19 from the absorption of ozone in organic free water. It was found to be 0.05 s-1 at pH 2.5. On the other hand, the individual mass-transfer coefficient, kL, was taken as 10-4 m s-1, which represents a typical value in semibatch reactors.20 The film length was estimated from the equation of film theory

δL )

DO3 kL

(10)

where DO3 is the ozone diffusivity, which was also taken from the literature.21 With this information, Ha was determined by a trial-and-error procedure from eq 8 and the experimental values of E reported in Table 1. This

table also presents the calculated Ha values. As can be seen, the Ha numbers corresponding to experiments carried out at catalyst concentrations higher than 1 mg L-1 are very close to 0.3, the limit at which the moderate kinetic regime is assumed to start. For the rest of the conditions investigated, Ha is much lower than this limit, which confirms the slow kinetic regime. According to these values, chemical reactions only controlled the ozone absorption rate when the Co(II) concentration in water was lower than 1 mg L-1. Mechanism of Reactions. Although the establishment of a mechanism of reactions is beyond the scope of this paper, some comments on this matter based on the experimental results obtained are presented here. As reported in a previous work,17 ozone barely reacts with oxalic acid in water. At acidic pH, as in this work, free radicals, or, more precisely, hydroxyl radicals, cannot be generated from the hydroxyl ion decomposition of ozone as is the general case in the water ozonation of micropollutants.5 Thus, it is evident that the only way that oxalic acid ozonation can occur is through other free-radical reactions or direct mechanism pathways such as have recently reported (i.e., through reactions involving ozone-cobalt(II)-oxalic acid complexes13). It is also known that the reaction of Co(II) with ozone at acidic conditions (pH < 1.6) leads to the formation of hydroxyl radicals. Hydroxyl radicals, once formed, could react with oxalic acid and provide another route to mineralization, as was recently shown in the presence of an activated carbon.17 The reactions that would imply the participation of hydroxyl radicals are as follows10,11

Co2+ + O3 + H2O f Co(OH)2 + + HO• + O2

(11)

HO• + O3 f HO2• + O2

(12)

HO• + B f products (CO2 + H2O)

(13)

HO2• + Co(OH)2+ f Co2+ + H2O + O2

(14)

where B represents oxalic acid. Because the rate constants of these reactions are known,4,10,11 a kinetic model was prepared from mass balance equations of ozone and oxalic acid by assuming that only reactions 11-14 occurred. The model was solved using Runge-Kutta methods to yield negligible destruction of oxalic acid by this route. According to these results, it seems evident that oxalic acid is removed by a different mechanism pathway. Also, one of the ways to ascertain the presence of hydroxyl radicals is the use of scavengers such as tertbutyl alcohol.5 Thus, some catalytic ozonation experiments were carried out in the presence of tert-butyl alcohol. As also shown in Figure 2, however, under the conditions investigated here, tert-butyl alcohol does not affect the oxalic acid removal rate. This means that direct reactions of ozone, likely with cobalt-oxalate complexes, are responsible for the disappearance of oxalic acid. These reactions should be those started by the formation of cobalt(II)-oxalic acid complexes, as have already been reported at near-neutral pH conditions.13 At pH 2.5, oxalic acid is mainly present in monoacid oxalate form (HC2O4-), and four different cobalt(II)-oxalate complexes can form. In Table 2, values of the pK’s of oxalic acid and the stability constants of these complexes are listed.22,23 According to these values, if the total concentrations of Co(II) and oxalic

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Table 2. pK Values of Oxalic Acid and Stability Constants (Ks) of Cobalt(II)-Oxalic Acid Complexes in Watera species

pK

ref

oxalic acid

pK1 ) 1.2 pK2 ) 4.2

22

log KS Co(HC2O4)2 Co(HC2O4)+ Co(C2O4) Co(C2O4)22a

10.6 5.5 3.5 5.8

23 23 23 23

T ) 25 °C.

Table 3. Distribution of Concentration of Oxalic Acid Species and Cobalt(II)-Oxalic Acid Complexes at the Start of One Ozone-Cobalt(II)-Oxalic Acid Experimenta species

concentration (M)

species

concentration (M)

H2C2O4 HC2O4C2O42Co(HC2O4)2

3.7 × 10-4 7.44 × 10-3 1.48 × 10-4 1.69 × 10-5

Co(HC2O4)+ Co(C2O4) Co(C2O4)2) Co(II)

1.8 × 10-8 3.6 × 10-12 1.1 × 10-13 7.7 × 10-12

a Experimental conditions: total Co(II) concentration ) 1 mg L-1 (1.7 × 10-5 M), total initial oxalic acid concentration (CB0) ) 8 × 10-3 M, pH ) 2.5.

should be the reaction responsible for the removal of oxalic acid. Also, notice that, according to the values of the stability and rate constants of complex formation and decomposition, the time needed to form these complexes can be estimated. Thus, the rate constants for the formation of Ni(II) and Co(II) complexes are between 104-107 M-1 s-1, whereas those of complex decomposition are between 102-103 s-1.24-28 For example, by assuming 100 s-1 for the rate constants of the inverse reactions in the cobalt(II)-oxalic acid complex equilibrium, all complexes are formed in less than 10 s. Because the addition of ozone to the aqueous solution containing oxalic acid and Co(II) was made after a much longer period of time, when ozonation started, the concentrations of oxalic acid and the Co(II) species were likely as indicated in Table 3. Therefore, it can be accepted that the reaction of ozone was mainly with Co(HC2O4)2 under the conditions investigated. Again, using Runge-Kutta methods, an estimation of the rate constant of the reaction between ozone and Co(HC2O4)2 was made. Figure 6 also presents the influence of this rate constant value on the concentration profile of oxalic acid in water once it was assumed the reaction of ozone and Co(HC2O4)2 was responsible for the removal of oxalic acid. For the experimental conditions indicated above, a value of the rate constant between 1000 and 1200 M-1 s-1 leads to the best fit between the experimental and calculated results. As reported previously,13,29 the partial donation of electron density from oxalate to Co(II) during complex formation might also increase the reactivity of Co(HC2O4)2 with ozone. After reaction, a cobalt(III)-oxalate complex is likely formed that decomposes to regenerate Co(II), giving rise to the appearance of superoxide ion radical and carbon dioxide.13 A possible mechanism in this case would be

Co(HC2O4)2 + O3 f Co(HC2O4)2+ + O3-• (15) Figure 6. Comparison between experimental and calculated profiles of dimensionless concentration of oxalic acid remaining with time for one homogeneous catalytic ozonation experiment. Conditions: gas flow rate ) 24 L h-1, agitation speed ) 300 rpm, T ) 20 °C, pH ) 2.5, catalyst concentration ) 0.8 mg L-1, initial oxalic acid concentration ) 8 × 10-3 M, ozone gas concentration ) 30 mg L-1. Symbols correspond to experimental points. Lines are results calculated (1) by assuming only ozone reactions with Co(C2O4) and Co(C2O4)22- and (2-7) by considering only the ozone-Co(HC2O4)2 reaction with the following assumed values for the rate constant (M-1 s-1): (2) 400, (3) 600, (4) 800, (5) 1000, (6) 1200, (7) 1400.

acid are 1 mg L-1 and 8 × 10-3 M, respectively (typical conditions applied in this work), then Co(HC2O4)2 is the main complex formed in water with a rather low presence of the three other complexes (see Table 3). The rate constants of reactions between ozone and Co(C2O4) and Co(C2O4)22- were reported at 25 °C in a previous paper13 as 30 and 4000 M-1 s-1, respectively. Using these values and mass balances of ozone and oxalic acid and assuming that these reactions were the only ones responsible for oxalic acid ozonation, the theoretical concentration of oxalic acid remaining was determined as a function of. As shown in Figure 6, for the conditions investigated, the concentration of oxalic acid remaining in water was still high after 90 min of reaction, likely because of the very low concentrations of these two Co complexes. Because Co(HC2O4)+ is also present at very low concentration, the reaction of ozone with Co(HC2O4)2

Co(HC2O4)2+ f Co(HC2O4)+ + 2CO2

(16)

Co(HC2O4)+ + HC2O4- f Co(HC2O4)2

(17)

and

According to these reactions, the theoretical overall stoichiometric ratio of the ozone-oxalic acid reaction is 1, which is in agreement with many of the results observed experimentally in this work. Notice, also, that reaction 15 generates the ozonide ion radical, O3-•, which eventually forms the hydroxyl radical, HO•.5 This free radical could attack both ozone and oxalic acid. In fact, some of the stoichiometric results lead to values greater than 1 for the overall stoichiometric coefficient of the ozone-oxalic acid reaction. Thus, it is possible that, for some conditions, the consumption of ozone is also due to the appearance of hydroxyl radicals. This would agree with a stoichiometry of more than 1 mol of ozone consumed per mole of oxalic acid consumed. However, previous literature information13 does not support this pathway for ozone consumption, as reported for oxalic acid oxidation in the presence of Co(II) at pH 7, where hydroxyl radicals were also noticed, but not their reaction with ozone. On the other hand, the consumption of oxalic acid by hydroxyl radicals can be disregarded from the Co(II) catalytic ozonation experiment in the presence of tert-butyl alcohol (see Figure 2).

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Figure 7. Determination of apparent ozone reaction order. Conditions as in Figure 2.

Empirical Rate Equation. A more rigorous study on the mechanism and kinetics of this system will be the subject of future work. Here, however, one of the aims of the current work was to obtain an empirical reaction rate equation to later measure any possible contribution of homogeneous catalysis to the total oxalic acid ozonation in the presence of a Co solid catalyst (see the second part of this work14). To this end, the empirical rate equation presented below was finally adopted

-rB ) kCBnCO3mCCop

(18)

where n, m, and p are the apparent reaction orders with respect to oxalic acid, ozone, and catalyst, respectively, and CCo is the Co(II) concentration. In a well-mixed semibatch reactor, the mass balance of oxalic acid is

-

dCB ) kCBnCO3mCCop dt

(19)

It was observed from experimental results (not shown) that the concentration of dissolved ozone increased sharply during the first 3 min of reaction to reach a constant value. According to the concentrations used in this work, Co(II) is the limiting species, although it is continuously regenerated (see reactions 15-17), so its concentration can also be considered constant. Therefore, after a very short initial period (about 5 min) needed for the ozone dissolved concentration to reach a constant value as well, eq 19 becomes

dCB ) khomCBn dt

(20)

khom ) kCO3mCCop

(21)

where

Experimental data on the oxalic acid concentration with time were fitted to different kinetic orders between 0 and 1, with first-order leading to the best-fit results. Then, apparent first-order kinetics was assumed (n ) 1 in eqs 18-20). Values of khom corresponding to experiments carried out at different ozone partial pressures or ozone gas concentrations were then plotted, in logarithmic coordinates, against the logarithm of the corresponding dissolved ozone concentrations reached in these experiments, as shown in Figure 7. As can be seen from this figure, a good linear relationship was observed. From the least-squares fitting of the experimental data, the slope of the straight line was found to be 0.4, which resulted as the apparent reaction order

Figure 8. Determination of apparent catalyst reaction order. Conditions as in Figure 3.

with respect to ozone. Then, eq 21 becomes

khom ) kCO30.4CCop

(22)

Next, from experiments carried out at different catalyst concentrations, the value of p (apparent reaction order with respect to the catalyst) was also determined. Although CO3 varied from one experiment to the next, it remained constant for nearly the whole reaction period within each experiment. Thus, according to these observations and eq 22, a plot of the logarithm of the ratio between khom and CO30.4 against the logarithm of the cobalt concentration should lead to a straight line with a slope equal to p. This plot is shown in Figure 8, corresponding to the experiments at different catalyst concentrations. As can be seen, the experimental points corresponding to mass concentrations of catalyst lower than 1 mg L-1 are situated about a straight line, whereas points of higher catalyst concentrations deviate from this linearity. The least-squares analysis of the straight-line portion for low catalyst concentrations (R2 ) 0.99) has a slope of 1.13. If all experimental points are considered, the least-squares analysis is worse (R2 ) 0.95), with a slope of 1.7, which would be the apparent reaction order for the catalyst valid for the whole range of concentrations investigated. The deviation from linearity of the points at higher catalyst concentrations can be explained if one considers the kinetic regime of ozone absorption. Thus, for catalyst concentrations above 1 mg L-1, the dissolved ozone concentration was very low, about 10-5 M or lower, which suggests a change in the kinetic regime (see Table 1). For low or negligible ozone concentrations, a moderate kinetic regime can be assumed to hold, as the data on the apparent Hatta number confirmed (see kinetic regime section). For low catalyst concentrations, in contrast, the kinetic regime can be classified as slow because of the significant dissolved ozone concentrations measured and lower apparent Hatta numbers found. Thus, the deviation of experimental points corresponding to catalyst concentrations above 1 mg L-1 was likely due to the change in kinetic regime from slow to moderate, where mass transfer also controls the rate of the process. As a consequence, to determine the apparent activation energy of this reaction system, only the experiments with catalyst concentrations lower than 1 mg L-1 were considered, as they developed in the slow kinetic regime where chemical reactions controlled the rate of the process. Thus, following the empirical kinetic study, the ratio between khom and the product CO304CCo1.1 corresponding to experiments at different temperatures

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The results obtained in this work confirm that Co species can be useful as the starting material in the preparation of solid catalysts for heterogeneous catalytic ozonation, as is shown in the second part of this work.14 Acknowledgment This work was supported by the CICYT of Spain and The European Region Development Funds of the European Commission (Project PPQ2000/0412). Literature Cited Figure 9. Determination of apparent activation energy of the Co(II) homogeneous catalytic ozonation of oxalic acid. Conditions as in Figure 4.

represents the apparent reaction rate constant k in eq 22. The logarithm of this constant was plotted against the inverse of temperature (see Figure 9) to determine the following Arrhenius equation

k ) 4.76 × 1015 exp(-7207/T)

R2 ) 0.99

(23)

From eq 23, an apparent activation energy of about 15 kcalmol-1 was deduced for the homogeneous catalytic ozonation of oxalic acid in the presence of Co(II). The rate equation describing the homogeneous catalytic ozonation of oxalic acid by Co(II) under chemical reaction control is finally given by eq 24

-r ) 4.76 × 1015 exp(-7207/T)CBCO30.4CCo1.1 M min-1 (24) Equation 24 was finally used to account for the contribution of homogeneous catalysis during ozonation of oxalic acid in the presence of a solid cobalt oxide/ alumina catalyst (see second part of the work).14 4. Conclusions The presence of Co(II) in water significantly enhances the ozonation rate of oxalic acid at acidic pH. The stoichiometry of the process was found to vary between 0.7 and 1.4 mol of ozone consumed per mole of oxalic acid consumed, and the ozone efficiency reached a maximum value of 25% in several cases. The reaction leads to nearly total mineralization of oxalic acid, which demonstrates that this kind of process could be very appropriate for the removal of refractory organic matter from water. The ozonation develops in the slow kinetic regime of absorption for catalyst concentrations lower than 1 mg L-1. For higher catalyst concentrations, the kinetic regime approaches moderate, which means that masstransfer limitations control the rate of the process to some extent. In the slow kinetic regime, an apparent first-order reaction with respect to oxalic acid and catalyst and nearly one-half reaction order with respect to ozone were determined from the experimental data. The apparent activation energy was found to be 15 kcal mol-1. The experimental results suggest that the mechanism of Co(II) catalytic ozonation of oxalic acid goes through the reaction of ozone with Co(HC2O4)2, although further studies on this matter are still needed.

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Received for review December 6, 2002 Revised manuscript received March 20, 2003 Accepted May 1, 2003 IE0209982