Kinetics of Catalytic Ozonation of Oxalic Acid in ... - ACS Publications

The ozonation kinetics of oxalic acid in water in the presence of an activated carbon has been investigated at acid pH. The presence of the activated ...
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Ind. Eng. Chem. Res. 2002, 41, 6510-6517

Kinetics of Catalytic Ozonation of Oxalic Acid in Water with Activated Carbon Fernando J. Beltra´ n,*,† Francisco J. Rivas,† Lidia A. Ferna´ ndez,‡ Pedro M. A Ä lvarez,† and Ramo´ n Montero-de-Espinosa† Departamento de Ingenierı´a Quı´mica y Energe´ tica, Universidad de Extremadura, 06071 Badajoz, Spain, and Centro de Investigaciones del Ozono, Centro Nacional de Investigaciones Cientı´ficas, Ciudad de La Habana, Cuba

The ozonation kinetics of oxalic acid in water in the presence of an activated carbon has been investigated at acid pH. The presence of the activated carbon significantly enhances the degradation rate of oxalic acid if compared to single ozonation and single adsorption. According to total organic carbon measurements, nearly complete mineralization of oxalic acid can be achieved depending on the experimental conditions. The presence of tert-butyl alcohol, which scarcely adsorbs on the carbon surface at the conditions investigated, led to a significant reduction of the oxalic acid removal rate. Consequently, experimental results suggest that the reaction proceeds in the water phase between oxalic acid and oxidant species, likely hydroxyl radicals, coming from the ozone decomposition on the carbon surface. The proposed mechanism yielded a first-order kinetics with respect to ozone, close to the 0.8 order experimentally observed. Also, the energy of activation was found to be approximately 15 kcal mol-1. Introduction Nowadays, increasing attention is being paid to research studies on the catalytic ozone decomposition on different solids to oxidize the matter present in water. Thus, in a recent review1 the present state of the art on this subject has been presented. Also, studies on ozone gas adsorption on alumina, CeO2, activated carbon (AC), etc., indicate the formation of oxygen-active species adsorbed on the catalyst surface.2-4 Although these studies have the ozone destruction in the gas as their main objective, their results can be useful to understand the mechanism of catalytic ozonation in water. Different solid catalystssoxides of iron, titanium, cerium, ruthenium, or AC-based catalystsshave so far been studied5-8 to improve the oxidizing capacity of ozone toward a wide variety of potential pollutants (i.e., phenols, aromatic hydrocarbons, carboxylic acids, dyes, surfactants, etc.). Some of these works report on the significant organic carbon reductions experienced if compared to the noncatalytic ozonations. This paper deals with the kinetics of the AC catalytic ozonation of oxalic acid in water. Oxalic acid, because of its refractory character, is usually found as an end product in ozonation processes (i.e., ozonation of phenols).9,10 For this reason, oxalic acid was chosen as a model refractory compound. In fact, oxalic acid hardly reacts with ozone, with the rate constant of this reaction being lower than 0.04 M-1 s-1.11 Moreover, the reaction of oxalic acid with hydroxyl radicals cannot be catalogued among the fast reactions that these types of radicals undergo. Rate constants of fast reactions of hydroxyl radicals are in the range of 109-1010 M-1 s-1,12 but that with oxalic acid is about 2 orders of magnitude * To whom correspondence should be addressed. E-mail: [email protected]. Fax: 34-924-271304. Phone: 34-924-289387. † Universidad de Extremadura. ‡ Centro Nacional de Investigaciones Cientı´ficas.

lower (approximately 107 M-1 s-1).12 Then, removal of oxalic acid would require a high concentration of hydroxyl radicals or a process where the reaction was due to a different mechanism (i.e., a catalytic surface reaction). The literature also shows that AC can be a useful catalyst to accelerate the ozone reactions in water.5,13,14 Kinetic studies of ozonation in water have only been carried out for reactions conducted in the absence of solid catalysts.11,15 These include ozone-involved advanced oxidation kinetics, the effects of mass transfer, determination of the rate constants of the reactions between ozone and/or the hydroxyl radical with compounds present in water, etc.16 So far, catalytic ozonation studies of dissolved substances in water have usually focused on the discussion of experimental results related to the extent of organic abatement, consumption of ozone, removal of TC, or comparison with other noncatalytic ozonations.1 These kinds of processes present a potential interest because of the improvement observed in the removal rates of some refractory contaminants in water. Then, studies on the kinetics of the catalytic action of AC are of potential interest because this adsorbent material, when ozone is also applied, could act as a catalyst in water treatment facilities. Experimental Part Oxalic acid (dihydrate form) was obtained from Merck, while AC (Gmi 12 × 40 US size) was purchased from Gallaquim, S.L. (Spain). The AC was dried to remove moisture and kept in a water-free atmosphere before use. The main characteristics of this AC were 789 m2 g-1 surface area determined by the BrunauerEmmett-Teller (BET) method with a Quantachrome autosorb 1 automated gas adsorption system, 0.87 g cm-3 apparent density, and 1.49 g cm-3 true density obtained from mercury and helium porosimetry by using Steropicnometer and Autoscan-60 Quantachrome apparatuses, respectively. A total pore volume of 0.66 cm3

10.1021/ie020311d CCC: $22.00 © 2002 American Chemical Society Published on Web 11/08/2002

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g-1 was calculated by means of the aforementioned densities. FT-IR spectra of the nonozonated and ozonated AC in the presence and absence of oxalic acid were also acquired by collection of 40 scans at 2 cm-1 resolution on a Perkin-Elmer spectrophotometer in the 4000-450 cm-1 wavenumber range. The absorption spectra was recorded using a KBr pellet containing 0.1% of carbon. These pellets were dried overnight at 100 °C before being recorded. The spectrum of the pure KBr pellet was used as the background while recording sample spectra. Catalytic experiments of slurry type were carried out in a 500 mL glass-jacketed vessel provided with an agitation impeller. The reactor was filled with 400 mL of ultrapure water (obtained from a Millipore MilliQ system), where a known amount of AC (0.625-1.25 g L-1) was then added. Agitation provided was in all cases high enough to keep the slurry system perfectly mixed. All experiments were carried out by duplicate in the presence (catalytic) and absence (noncatalytic or blank experiment) of AC. Also, adsorption experiments (in the absence of ozone) were completed as a second blank type experiment for comparative purposes. The temperature inside the reactor was kept constant within (0.5 °C by circulating water through the jacket from a thermostatic batch. The aqueous solution contained a 8 × 10-3 or 4 × 10-3 M concentration of oxalic acid. Experiments were conducted at the pH of the oxalic acid aqueous solution, that is, pH 2.5. Once the slurry was charged and the reaction temperature achieved, in ozonation experiments, an ozone-oxygen mixture was fed to the reactor through a diffuser plate situated at the reactor bottom. Samples were steadily withdrawn for analysis. Ozone was produced from pure oxygen in a 301.7 Sander laboratory ozone generator. Thermal decomposition of oxalic acid onto AC in an inert atmosphere was studied by completing thermogravimetric and differential thermogravimetric measurements in a Mettler MTA 3000 system provided with a Mettler TG50 thermobalance. Samples of approximately 110 mg of AC were placed in a platinum crucible and heated in a dynamic nitrogen atmosphere (200 mL min-1) at 5 °C min-1 in the 20-250 °C temperature range. The equilibrium adsorption isotherm for oxalic acid from an aqueous solution on AC was determined at 20 °C using the batch bottle-point method. Different amounts of AC were weighed and suspended into glass bottles containing 20 mL of an organic solution. The bottles were sealed and placed in an orbital shaker, where samples were kept at a constant temperature within (0.1 °C. The concentration of oxalic acid was determined by high-performance liquid chromatography by means of a 1100 series Hewlett-Packard liquid chromatograph provided with an UV-visible detector and a Chemstation software package for data analysis and automatic control of chromatographic conditions. A 4.6 mm i.d., 250 mm length ODS-P column (Sugelabor, Spain) was used. The analysis was carried out isocratically with a 80/20 (v/v) water-acetonitrile mobile phase (phosphoric acid was used to keep the water pH at 2.5). The flow rate was set at 1 mL min-1. 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 ozone concentration in the aqueous phase was followed by the Indigo method17 after filtration of the samples. In some experiments, the total carbon (TC) was also

Figure 1. Evolution of dimensionless remaining oxalic acid concentration with time corresponding to experiments of single adsorption ([), single ozonation (9), and catalytic ozonation with AC (2) of oxalic acid in water. Experimental conditions: initial Coxalic ) 8 × 10-3 M; CO3(gas) ) 30 mg L-1; mass of carbon ) 1.25 g L-1. Carbon size: 1.0-1.6 mm. Gas flow rate: 15 L h-1. Agitation: 200 rpm. T ) 20 °C, pH 2.5.

Figure 2. Isothermal data of the adsorption of oxalic acid from water into AC at 20 °C and pH 2.5.

followed. In these cases, a DC-190 Dorhmann TOC analyzer, based on carbon dioxide infrared absorption, was used. Because experiments were carried out at pH 2.5, TC values actually corresponded to the remaining total organic carbon (TOC). Results and Discussion To eliminate any possible contribution of the reaction between hydroxyl radicals, generated from the direct ozone decomposition in water, and oxalic acid, experiments on adsorption and catalytic and noncatalytic ozonation of oxalic acid were carried out at pH 2.5. Also, it should be pointed out that the ozonation of high molecular weight compounds such as aromatic hydrocarbons or phenols eventually leads to acidic aqueous solutions where carboxylic acids, such as oxalic acid, are the main end products.9,10 These compounds are usually refractory toward further ozonation and would accumulate in water.10 Consequently, the ozonation at acid pH would represent a situation potentially present in these kinds of processes. In Figure 1 comparative examples of the results obtained from adsorption and catalytic and noncatalytic ozonation of oxalic acid are presented. It is observed that single ozonation does not lead to any significant removal of oxalic acid after 4 h of reaction time. A similar statement applies to single adsorption which only allows for approximately 8% of oxalic acid removal after 4 h of treatment. This represents about 50 mg of oxalic acid adsorbed/g of AC, which is a reduced fraction of the total amount of oxalic acid that the AC used can adsorb at equilibrium, as shown in Figure 2. Thus, from Figure 2 it is deduced that for concentrations of oxalic acid in the range of 4 × 10-3-8 × 10-3 M, used in this study, the corresponding equilibria concentrations on the carbon surface are about 380 and 330 mg of oxalic acid/g

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Figure 3. Evolution of dimensionless remaining TOC with time corresponding to experiments of oxalic acid ozonation and adsorption in the presence and absence of tert-butyl alcohol. General experimental conditions: CB0 ) 720 mg L-1, CO3(gas) ) 30 mg L-1, pH 2.5. Ccat ) 1250 mg L-1. AC particle: 1.0-1.6 mm. Gas flow rate: 15 L h-1. Agitation: 200 rpm. T ) 20 °C. tert-butyl alcohol concentration: 10-3 M. (b) Adsorption of tert-butyl alcohol alone. (+) Adsorption of oxalic acid alone. (2) Adsorption of oxalic acid and tert-butyl alcohol. (×) Single ozonation of oxalic acid. ([) AC catalytic oxalic acid ozonation in the presence of tert-butyl alcohol. (9) AC catalytic oxalic acid ozonation.

of AC, respectively. These results, compared to that found from the kinetics of oxalic acid adsorption, indicate that for the reaction times of this work oxalic acid adsorption is far from the equilibrium and, consequently, most of the active carbon surface is available. As a consequence, oxalic acid adsorption is weak at the conditions investigated (50 mg g-1 , 380-330 mg g-1). Contrarily to single ozonation and adsorption processes (see Figure 1), the simultaneous presence of AC and ozone leads to roughly 55% oxalic acid depletion. From the above experimental facts, the beneficial effect of AC as a catalyst is evident. Thus, as inferred from Figure 2, for the oxalic acid concentration/AC mass ratios used in this work, it is expected that, during catalytic ozonation experiments, most of the oxalic acid used will remain nonadsorbed. Because ozone hardly reacts with oxalic acid, any significant removal of this compound when ozone and AC are present must be attributed to some sort of catalytic oxidation. Furthermore, because the main objective of oxidation is mineralization of the organic matter present in water, TOC measurements were also carried out during the experiments. In Figure 3 the evolution of dimensionless remaining TOC with time corresponding to the experiments shown in Figure 1 (single adsorption, single ozonation, and AC catalytic ozonation) is shown. From the results of Figure 3, TOC removals of approximately 5, 6, and 50% are obtained from single ozonation, single adsorption, and AC catalytic ozonation, respectively. As can be observed, TOC profiles are similar to those found for oxalic acid removal (see Figure 1). This comparison suggests that the process leads to nearly complete mineralization of the parent compound, a hypothesis that was confirmed when TOC measurements were compared to TOC values calculated from the remaining oxalic acid concentration. This comparison is depicted in Figure 4 where the experimentally measured TOC is plotted against TOC values calculated from the remaining oxalic acid concentration. As can be seen, points of Figure 4 are situated around a straight line. From the least-squares fitting, the slope of this straight line was found to be close to 1 (see Figure 4 for the fitting results). Similar fitting results were obtained from experimental values at different working conditions. This represents the major advantage of the combined use of ozone and AC to remove oxalic acid from water. The high degree of mineralization is also

Figure 4. Comparison between experimental and calculated dimensionless remaining TOC corresponding to the catalytic ozonation of oxalic acid in water. Fitting results of the straight line: (TOC/TOC0)exp ) 1.06(TOC/TOC0)calc - 0.06 (R2 ) 0.98).

supported by the fact that the total consumption of ozone and oxalic acid for the experimental results of catalytic ozonation gives an empirical stoichiometric ratio between 1.3 and 2.4 mmol of oxalic acid consumed/ mmol of ozone consumed. These values are close to 2, the theoretical stoichiometric ratio for total carboxylic acid mineralization. Another interesting question related to ozonation reactions in water is the possible involvement of hydroxyl radicals. It is known that when using ozone, in the absence of catalysts, oxalic acid can only be removed by hydroxyl radicals coming from ozone decomposition. Thus, some previous works have reported the elimination of oxalic acid from water with ozone and hydrogen peroxide.18 The ozone-hydrogen peroxide combination, at pH > 7, leads to a high hydroxyl radical concentration, mainly responsible for organic oxidation. In ozonation systems, the high reactivity of hydroxyl radicals always brings the question about their participation in the process. As far as AC catalytic ozonation is concerned, Jans and Hoigne´13 have reported that the decomposition of ozone in water on an AC and a carbon black yields hydroxyl radicals in the water phase. These authors confirmed their results with the use of pchlorobenzoate, which does not react with ozone and, at the conditions they reported, does not adsorb on the carbon surface. Here, some experiments have been conducted in the presence of tert-butyl alcohol, a known hydroxyl radical scavenger.19 Figure 3 also shows the evolution of the dimensionless remaining TOC with time corresponding to adsorption and catalytic ozonation experiments of oxalic acid in the presence of tert-butyl alcohol. It is observed from Figure 3 that the presence of tert-butyl alcohol significantly inhibits the oxidation rate of oxalic acid compared to the AC catalytic ozonation without tert-butyl alcohol. This suggests that the process goes through formation of hydroxyl radicals that are scavenged in the presence of tert-butyl alcohol, with the result being a decrease of the ozonation rate of oxalic acid. However, this inhibition could also be due to the adsorption of tert-butyl alcohol on the carbon surface. If a strong adsorption of tert-butyl alcohol occurs, the number of active centers on the carbon surface available for adsorption of ozone and/or oxalic acid would diminish. This would reduce the adsorbed reactant species concentration and then the oxalic acid oxidation rate. Thus, some adsorption experiments of tert-butyl alcohol with and without the presence of oxalic acid were also carried out to ascertain whether tert-butyl alcohol is adsorbed by competing with the adsorption of oxalic acid. In Figure 3, the results of these adsorption experiments are shown. These results, expressed as the remaining TOC in solution, indicate that tert-butyl

Ind. Eng. Chem. Res., Vol. 41, No. 25, 2002 6513 Table 1. Thermogravimetric Analysis: Amount of Oxalic Acid Adsorbed on AC in Single Adsorption and Catalytic Ozonation Processes

sample type SAb catalytic ozonationc catalytic ozonationd catalytic ozonatione d

oxalic initial water acid % oxalic mass, mg mass, mg mass, mg acid massa 113.8 108.0 106.3 108.0

3.1 2.6 2.7 2.7

6.3 6.3 6.1 6.0

5.7 6.0 5.9 5.7

a Dry basis. b SA: single adsorption. c After 1 h of reaction time. After 2 h of reaction time. e After 3 h of reaction time.

alcohol scarcely adsorbs on the AC (during the first 150 min, less than 2% tert-butyl alcohol was adsorbed). Furthermore, if one compares the adsorption results of oxalic acid in the absence and presence of tert-butyl alcohol, it is observed that there is no difference at all between them. Thus, in absolute TOC values, removals of oxalic acid and tert-butyl alcohol, after 4 h, are approximately 1.5 and 1 mg L-1, respectively, from their single adsorption and 2.3 mg L-1 from the simultaneous adsorption of both tert-butyl alcohol and oxalic acid. The TOC removal in the simultaneous adsorption is then practically the sum of the individual adsorption processes, as can be deduced from these figures, but still it represents a negligible fraction of the TOC initially charged in the reactor (about 3%). Hence, at the conditions here investigated (for 4 h of reaction time), tertbutyl alcohol practically remains nonadsorbed as well. Because nearly all tert-butyl alcohol remains in solution, it is evident that the inhibition of oxalic acid oxidation in the presence of this compound is due to some phenomena occurring in the water phase. These phenomena are likely due to the inhibition of hydroxyl radicals that are generated from ozone-adsorbed species. This is in good agreement with the results of Jans and Hoigne´,13 who reported that ozone adsorption on the AC initiates a radical chain that finally proceeds in aqueous solution with the reaction of hydroxyl radicals formed. Adsorption of Oxalic Acid on AC. In Figure 1 the kinetic results of an adsorption experiment of oxalic acid in water are presented. As observed, after a few minutes from the start of adsorption, the variation of the oxalic acid concentration is so low that, at first sight, a plateau value in the concentration seems to be reached. This suggests that the system is at equilibrium, but the plateau concentration does not correspond to the equilibrium one, as can be inferred from the results of Figure 2. In fact, if experimental points of oxalic acid single adsorption are fitted through linear least-squares analysis, the slope of the resulting straight line equation has a negative value (-1.6 × 10-6 M min-1), which means that oxalic acid is being slowly adsorbed. This conclusion is also confirmed from the results shown in Figure 3, where adsorption of oxalic acid is measured as TOC. Here, it is clearly observed that TOC continuously and slowly decreases along the adsorption time considered (240 min). The results of the poor adsorption of oxalic acid on AC were also confirmed from thermogravimetric analysis. Table 1 summarizes some of the results obtained. As a consequence, given the equilibrium data shown in Figure 2, a weak oxalic acid adsorption took place that allowed a high fraction of the active carbon surface be available. As can also be seen from Table 1, the amount of oxalic acid desorbed from the AC was not dependent on the type of process where the AC was used. This means that the catalytic ozonation rate was

Figure 5. FT-IR spectra data corresponding to different processes: (A) untreated AC; (B) ozonated AC at pH 2.5; (C) ozonated AC at pH 7.5; (D) AC from oxalic acid adsorption at pH 2.5; (E) AC from catalytic ozonation of oxalic acid at pH 2.5.

not dependent on the oxalic acid adsorption. Also, thermogravimetric, single adsorption, and catalytic ozonation results suggest that byproducts of oxidation likely represented a small fraction of the total mass adsorbed. Adsorption of Ozone on AC. Studies of the decomposition of gaseous ozone on AC and other catalysts have shown that ozone is adsorbed to yield different species (ozonide, superoxide ion, etc.).2-4 Because of the unstable ozone-adsorbed species, in one of these works3 IR detection of ozonide-adsorbed species was directly carried out inside the reactor cell. In water, infrared analysis of AC samples subject to ozonation at pH higher than 2.5 clearly showed that some kind of chemisorption also occurs on the solid surface14 (see also Figure 5). In this study, infrared analyses have also been carried out on AC samples in the presence and absence of ozone and oxalic acid. Thus, Figure 5 shows that the FT-IR spectrum A of the initial fresh AC only exhibits a broad band at about 1100 cm-1, which can be attributed to ν(C-O) absorption of ether-type surface groups. Also, it is noticeable that the peak at about 1590 cm-1 is likely due to the CdC stretching mode for aromatics. The spectrum B corresponds to the AC after being subjected to ozonation in water at pH 2.5 for 1 h. At first sight, this spectrum is shaped similarly to spectrum A prepared in the absence of ozone. This means that no significant changes on the carbon surface were produced as a result of ozonation at this pH. However, at about 700-800 cm-1, some small peaks are also observed in spectrum B. Bands of different forms of adsorbed ozone are reported to appear at about 1100, 1035, and 705 cm-1.4 Thus, it is likely that the small

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bands in spectrum B represent ozone-adsorbed species. The fact that these bands are so small could be attributable to the fact that adsorption of ozone is mainly important on basic points4 that are not present in an acidic environment like the aqueous solution of oxalic acid. To support this fact, spectrum C, corresponding to an AC sample treated with ozone at pH 7.5, is also presented. In this case a peak at 790 cm-1 is clearly distinguished. Another possible explanation of the nearly peak absence in spectrum B is that physically adsorbed ozone rapidly decomposes at room temperature.4 In Figure 5 spectrum D of an AC after oxalic acid adsorption is also presented. The development of infrared adsorption bands occurs as a result of oxalic acid adsorption. Significant bands appear at 3450, 2900, and 1700 cm-1. These bands can be attributed to infrared active functionalities of dehydrated oxalic acid.20 Spectrum E, finally, corresponds to an AC sample from the catalytic ozonation of oxalic acid at pH 2.5. Again, the main bands noticed correspond to the adsorbed oxalic acid. In this case, bands at 1700 and 3450 cm-1 are clearly developed likely as a consequence of the formation of some oxidation products that remain adsorbed onto carbon. Notice, however, that the fraction of oxidation products that remain adsorbed should be very low because thermogravimetric analysis of AC from oxalic acid adsorption and catalytic ozonation leads to a similar loss of mass attributable to adsorbed oxalic acid. The results derived from oxalic acid adsorption and catalytic oxidation with and without tert-butyl alcohol and in the absence of ozone-adsorbed species from the IR analysis strongly suggest a fast decomposition or desorption step of ozone-adsorbed species to yield hydroxyl radicals in water, which then react with oxalic acid. Kinetics and Mechanism. Irreversible heterogeneous gas-liquid and gas-liquid-solid (porous catalyst) reactions develop through steps in series processes:21 external diffusion or mass transfer through the gas and liquid film layers close to the gas-water interface, reaction in the bulk liquid, external diffusion or mass transfer through the liquid film surrounding the catalyst particle, and internal diffusion accompanied with surface reaction inside the catalyst pores. When masstransfer resistances (external and internal) are considered negligible, chemical reactions (both in the bulk liquid or at the catalyst surface) are the limiting step and control the process rate. Under this chemical kinetic regime, the rate constant of the catalytic reaction can be directly determined from experimental data.22 For doing so, it is advisable to establish the conditions for this kinetic regime. This is usually made by first studying the effect of hydrodynamic variables such as the gas flow rate, agitation speed (the case of agitating contactors), and catalyst particle size (the case of solid catalytic reactions) on the reaction rate. Influence of Mass-Transfer-Related Variables. Table 2 presents the mass-transfer-related variables investigated and their corresponding interval ranges. During ozonation processes, gas-side resistance to mass transfer can be neglected because ozone is a scarcely soluble gas in water.23 In this work, no effect at all of these variables (gas flow rate, agitation speed, and particle size) was noticed. Then, following these results, it seemed that the process rate was controlled by

Figure 6. Evolution of dimensionless remaining oxalic acid concentration with time corresponding to experiments of different temperatures. General experimental conditions: CB0 ) 720 mg L-1, CO3(gas) ) 30 mg L-1, pH 2.5. Ccat ) 1250 mg L-1. AC particle: 1.0-1.6 mm. Gas flow rate: 25 L h-1. Agitation: 200 rpm. T (°C): (9) 10; ([) 20; (2) 30.

Figure 7. Evolution of dimensionless remaining oxalic acid concentration with time corresponding to experiments of different initial oxalic acid concentrations. General experimental conditions: CO3(gas) ) 30 mg L-1, pH 2.5. Ccat ) 1250 mg L-1. AC particle: 1.0-1.6 mm. Gas flow rate: 15 L h-1. Agitation: 200 rpm. T ) 20 °C. CB0 (mg L-1): (9) 360; ([) 720. Table 2. Variables Investigated and the Range of Application variable

range studied

units

gas flow rate, vG agitation speed, N AC particle size, dp temperature, T ozone gas concn, CO3(gas) AC concn, Ccat oxalic acid concn, CB

7-25 100-300 (1-1.6) to (2.0-3.15) 10-30 15-52 625-2500 360-720

L h-1 rpm mm °C mg L-1 mg L-1 mg L-1

chemical reactions in the water or by surface phenomena. Influence of Chemical-Reaction-Related Variables. These include the temperature, concentration of reactants (oxalic acid and ozone), and catalyst. Table 2 also lists interval ranges investigated for these variables. Figures 6 and 7 show as examples the evolution of the oxalic acid dimensionless remaining concentration with time corresponding to experiments completed at different temperatures and initial oxalic acid concentrations, respectively. From Figure 6, it can be seen that, at a given time, the increase of the temperature leads to an increase of the oxalic acid conversion. It can also be observed that, at the highest temperature investigated, 30 °C, total conversion of the acid is achieved in less than 4 h. Also, the effect of the ozone concentration in the gas fed to the reactor and mass of catalyst on the process rate (not shown) was positive. On the contrary,

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as observed from Figure 7, increasing the initial concentration of oxalic yields a decrease of its conversion. From Figures 6 and 7, a zero-order kinetics with respect to oxalic acid can also be deduced, a situation also observed in the rest of the experiments. Also, regarding the kinetics, the significant influence of temperature on the oxalic acid conversion, at constant values of the rest of variables, confirms the hypothesis of a chemical phenomenon controlled process. Moreover, according to the results obtained, the positive influence of the ozone concentration and AC mass or surface suggests the oxidation kinetics to be a function of these variables as shown later. The zero-order kinetics with respect to the oxalic acid concentration, the slight adsorption of this compound on the carbon surface, and the effect of tertbutyl alcohol (which is also slightly adsorbed for reaction times higher than 150 min) support the conclusion on the reaction in the bulk water, likely between hydroxyl radicals and oxalic acid. According to precedent comments, a possible simple reaction mechanism for the AC catalytic ozonation of oxalic acid would consist of the following steps:

Surface phenomena: Adsorption steps: k1d

O3 + S y\ z O3-S k 1i

k2d

z B-S B + S y\ k 2i

(1) (2)

Reaction and desorption steps: k3d

z O-S + O2 O3-S y\ k 3i

k4d

O-S + H2O y\ z HO• + OH- + S k 4i

(3)

Reaction in bulk water: k5

(5)

Notice that steps (1) and (3) are well documented in the literature.2 Step (2) is here proposed to account for the low contribution of physical adsorption to remove oxalic acid (as much as 8% at the conditions investigated), while step (4) is proposed to account for the formation of oxidant species in water. According to works on ozone decomposition, it seems reasonable to admit that these entities are hydroxyl radicals. Furthermore, as commented on before, the inhibition of oxidation due to tertbutyl alcohol reinforces the conclusion about hydroxyl radical participation. Step (5) finally accounts for the bulk water oxidation reaction. Following this mechanism, the oxalic acid removal rate, in the semibatch reactor used, will be given by eq 6, where -rad would represent the contribution of

-dCB/dt ) k5CHOCB + (-rad)

-dCB/dt ) k1dCO3Cv

(7)

where Cv represents the concentration of vacant active sites at the carbon surface. As commented on before, most of the active surface should be available so that Cv can also be a constant that would depend on the catalyst surface. Then, eq 7 can finally be as follows:

-dCB/dt ) kTCO3

(8)

kT ) k1dRCncat ) kwCncat

(9)

where

where n is the reaction order with respect to the catalyst concentration. From the experimental results, after a few minutes from the start of ozonation, a constant value of the concentration of dissolved ozone in water was always observed CO3s (not shown). Then eq 8 represents a zeroorder kinetics:

-dCB/dt ) k′

(10)

k′ ) kTCO3s

(11)

where

which should be integrated within the following limiting conditions:

t ) tsCB ) CBs (4)

where S and B stand for any free active center on the AC surface and oxalic acid, respectively.

HO• + B 98 2CO2 + H2O + H+

steps (1), (2), and (4) as irreversible, at the conditions applied (no ozone-adsorbed species are detected likely because of the fast rates of these steps), and the steadystate situation for the hydroxyl radical concentration (eq 6) reduces to the following one:

(6)

adsorption to the removal rate. However, taking into account the low contribution of adsorption, considering

t ) tCB ) CB

(12)

where ts is the time needed by ozone to reach the plateau concentration in water and CBs the concentration of oxalic acid at this time. Equation 10, however, can be integrated, without error, from the start of ozonation (t ) 0) because ts is negligible (a few minutes) compared to the total reaction period (4 h). Then, integration of eq 10, after variable separation, yields eq 13. As can be

CB k′ )1t CB0 CB0

(13)

seen in Figures 6 and 7, eq 13 is confirmed from empirical results that show the linear relationship between the oxalic acid concentration and time. From the least-squares analysis of these lines, values of k′ were obtained at different conditions from the slopes of the lines (k′ ) slope × CB0). The results are presented in Table 3. At 20 °C, from experiments completed at different concentrations of ozone in the feeding gas (not shown), the apparent rate constant kT and order of reaction with respect to ozone were also obtained from a plot of the ln(k′) against ln(CO3s) (concentrations of dissolved ozone, CO3s, were experimentally obtained). The results are shown in Figure 8. The least-squares analysis of the straight line (R2 ) 0.99) yields a value of 0.8 (close to 1) for the ozone reaction order and 0.0207 M-0.8 min-1 for the pseudo rate constant (kT).

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Table 3. Experimental Apparent Zero-Order Rate Constant Values of the AC Catalytic Ozonation of Oxalic Acida CO3(gas), mg 15 30 52 30 30 30 30b,c 30b,d 30b,e 30f

L-1

Ccat, mg

L-1

k′ ×

105,

1250 1250 1250 625 1870 2500 1250 1250 1250 1250

mol

L-1

min-1

1.12 1.84 2.64 1.28 2.40 3.28 1.20 1.92 3.60 1.94

a General conditions for surface reaction control experiments: 15 L h-1, 200 rpm, 1-1.6 mm, 20 °C, 720 mg L-1 oxalic acid unless indicated. b 25 L h-1. c 10 °C. d 20 °C. e 30 °C. f 360 mg L-1 oxalic acid.

Figure 8. Verification of eq 11. Determination of the rate constant kT and ozone reaction order.

Figure 9. Verification of eq 9. Determination of kw and the catalyst reaction order at 20 °C.

k′w ) 6.13 × 106 exp(-14420/RT) L1.56 mol-0.8 (m2 of catalyst)-0.76 min-1 (15) where the activation energy (14 420) is in cal mol-1. Conclusions The kinetic study of the AC catalytic ozonation of oxalic acid in water has been carried out at conditions where chemical reaction controls the process rate. The process implies decomposition of ozone on the carbon surface to generate oxidant species in the water, likely hydroxyl radicals, that then react with nonadsorbed oxalic acid. This is confirmed from the experimental results obtained in the presence of tert-butyl alcohol that is scarcely adsorbed onto the carbon surface, at least during the first 150 min of the process. Adsorption of oxalic acid is also low at the conditions investigated. Thus, in 4 h, it represents as much as 8% of the oxalic acid removed from water, approximately 50 mg g-1, while at equilibrium a maximum amount of oxalic acid between 300 and 400 mg can be adsorbed per g of AC. IR analysis indicates the adsorption of some end products, but this represents a negligible fraction of the mass adsorbed as deduced from TOC balance results. The AC catalytic ozonation of oxalic acid leads, therefore, to nearly complete mineralization. The experimental kinetics has been found to be of zero order with respect to oxalic acid and approximately first order with respect to ozone. The results can be explained from a mechanism involving surface phenomena (adsorption, reaction, and desorption) and bulk water reactions between oxidant species (hydroxyl radicals) and oxalic acid. The experimental reaction order with respect to ozone and the activation energy of the catalytic ozonation were found to be 0.8 and 14 420 cal mol-1, respectively. Because of the importance of AC in water treatment and the use of ozone as an energetic oxidant, more studies are now being undertaken to confirm the presence of hydroxyl radicals and their role on this catalytic ozonation. Acknowledgment

The results of the kinetics of catalytic reactions are usually expressed per unit of catalyst mass or catalyst surface. From eq 9, a plot of ln(kT) against ln(Ccat) should yield a straight line and values of n and kw could be obtained from the slope and intercept. In Figure 9, this plot is presented corresponding to the experiments carried out at different catalyst concentrations and at 20 °C. The least-squares analysis of the straight line (with R2 ) 0.99) leads to the values of n ) 0.76 and rate constant, kw ) 0.017 L0.76 (g of catalyst)-0.76 min-1 or k′w ) 1.07 × 10-4 L0.76 (m2 of catalyst)-0.76 min-1 (with k′w ) (SBET)-0.76kw). Finally, from the rate constants given in Table 3 corresponding to experiments completed at different temperatures (see Figure 6), an Arrehnius plot (not shown) yields the equation (R2 ) 0.99)

kw ) 9.75 × 108 exp(-14420/RT) L1.56 mol-0.8 (g of catalyst)-0.76 min-1 (14) or in terms of the surface of the catalyst

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Received for review April 25, 2002 Revised manuscript received September 17, 2002 Accepted October 1, 2002 IE020311D