4774
Ind. Eng. Chem. Res. 1997, 36, 4774-4778
Kinetics of Oxalic Acid Ozonation Promoted by Heterogeneous MnO2 Catalysis R. Andreozzi,*,† V. Caprio,‡ A. Insola,† R. Marotta,‡ and V. Tufano§ Istituto di Ricerche sulla Combustione, CNR, p.le V. Tecchio, 80125 Napoli, Italy, Dipartimento di Ingegneria Chimica, Facolta` di Ingegneria, Universita` di Napoli, p.le V. Tecchio, 80125 Napoli, Italy, and Dipartimento di Ingegneria e Fisica dell’Ambiente, Universita` della Basilicata, via della Tecnica 3, 85100 Potenza, Italy
A kinetic model is developed for the ozonation of oxalic acid catalyzed by solid MnO2. The rate of ozonation is limited by the adsorption of oxalic acid on the catalyst surface and by the deactivation of a fraction of the active sites, because of an irreversible reaction with ozone. Moreover, the model includes the ozonation of oxalic acid catalyzed by dissolved manganese. This kinetic model allows for a fairly good correlation of the experimental data when the average size of the catalyst particles is smaller than about 10 µm. When larger particles are employed, noticeable mass-transfer limitations are encountered, mainly deriving from the diffusion of both reactants from the liquid bulk to the solid surface (external and internal). The dependence of the rate constant on pH is experimentally determined and explained in terms of the changing chemical structure of the active sites and of the dissociation equilibrium of oxalic acid. Introduction Catalytic ozonation has been scarcely considered in the literature (Al-Hayek et al., 1989; Paillard et al., 1991, 1992; Andreozzi et al., 1992; Campen, 1993; Thompson et al., 1995), although its development could strongly affect technological advancement of ozonation processes for the treatment of wastewaters. In a previous investigation (Andreozzi et al., 1996a), manganese dioxide has been shown to possess appreciable activity as a heterogeneous catalyst for the oxidation of oxalic acid, which is refractory to ozonation. That study allowed us to assess some fundamental aspects of the catalytic mechanism, related both to the chemical nature of the active superficial sites and to the adsorption of the oxalic acid on the catalyst surface. Unfortunately, univocal information on the mechanistic role exerted by ozone could not be achieved, since alternative reaction pathways were suggested by the experimental results. On the basis of the mechanistic arguments made available by the above investigation, a quantitative description of the catalytic ozonation process is now attempted, also by means of mathematical models, with the aim of proposing a quantitative description of reaction kinetics and of obtaining a deeper insight into the chemical mechanism. The modeling analysis is also based on a wider experimental description of the system behavior, including temperature and pH dependences, which have been evaluated while minimizing the intervention of masstransfer phenomena. Experimental Section All the ozonation experiments were performed in the semibatch apparatus described in detail elsewhere (Andreozzi et al., 1992). For each run, 245 mg of dihydrate oxalic acid (99.8% by weight) was dissolved in 800 mL of an aqueous buffered solution, and a proper amount (mS ) 50 mg) of solid MnO2 was added. * Corresponding author. Telephone: [39](81)7682251. Fax: [39](81)5936936. E-mail:
[email protected]. † CNR. ‡ Universita ` di Napoli. § Universita ` della Basilicata. S0888-5885(96)00159-5 CCC: $14.00
Aqueous solutions of KH2PO4, Na2HPO4, and H3PO4 were used as buffer mixtures to keep constant the pH during the whole ozonation experiments. The solution was submitted to ozonation by bubbling a 36 L/h ozonized oxygen stream whose ozone content was 3% by volume. Samples withdrawn from the reactor at different ozonation times were analyzed by means of a Hewlett-Packard 1090L HPLC equipped with a spherisorb S5C6 reverse-phase column flowed with 1.0 mL/min of aqueous buffered acidic solution. Manganese dioxide (99.99% purity, by Aldrich) was separated by sieving the particles into different granulometric fractions and directly used without any treatment. The average radius was used to characterize the particle size. All the experiments were performed in the pH range 4.1-6.0, and in the temperature range 10-25 °C. Mass-Transfer Limitations The unsteady operation of the mechanically agitated gas-liquid-solid reactor is affected by different mass transfer phenomena, which occur both in series and in parallel with the chemical reaction. In particular, ozone diffuses from the gaseous bubbles to the liquid bulk, and both reactants (ozone and oxalic acid) diffuse from the liquid bulk to the external solid surface and, subsequently, to the internal solid surface. The gas-liquid mass transfer is described according to the fluid dynamic model previously developed (Andreozzi et al., 1991):
d[O3]L dt
Nr
) koLa([O3]BR - [O3]L) -
∑j rj
(1)
d[O3]B Q ) ([O3]in - [O3]B) dt VB koLa([O3]BR - [O3]L) VL (2) VB d[O3]F Q ) ([O3]B - [O3]F) dt VF
(3)
where the hold-up, VB/VL, the volumetric coefficient of © 1997 American Chemical Society
Ind. Eng. Chem. Res., Vol. 36, No. 11, 1997 4775
Figure 1. First-order kinetic constant, k1, (min-1), as a function of the average particle radius, R (µm). T ) 25 °C, pH ) 4.10.
mass transfer, koLa, and the Ostwald coefficient, R, have been evaluated from previous experiments (Tufano et al., 1994; Andreozzi et al., 1996b). In effect, some ozone is consumed in the liquid film, because of several reactions, namely, the homogeneous reaction catalyzed by dissolved manganese and the heterogeneous reaction catalyzed by very small MnO2 particles possibly entrained in the gas-liquid film. However, a rough evaluation of the Hatta number on the basis of the overall rate constant, kO3, derived from a simplified (steady-state) mass balance:
Q([O3]in - [O3]F) ) VLkoLa([O3]FR - [O3]L) ) VLkO3[O3]L (4) shows that the resulting enhancement factor (Astarita et al., 1983; Danckwerts, 1970) is not significantly different from one. The chemical submodel includes the reaction terms, written as a function of the composition of the liquid bulk. In the presence of mass-transfer limitations, this composition should be corrected with a suitable effectiveness factor (Froment and Bischoff, 1979). In effect, the external resistances (i.e., the liquid-solid mass-transfer limitations) are expected to be markedly larger than the gas-liquid resistances, mainly because of the smaller interfacial area (as evaluated from the average particle radius). Moreover, the relevance of the internal resistances should be carefully evaluated, since the BET measurements (specific area ) 6.9 ( 0.2 m2/g) indicate the presence of some porosity in the solid particles. A quantitative treatment of these phenomena is hindered by the lack of independent measurements of the effective diffusion coefficient and by the widespread distribution of the catalyst size. Thus, it was decided to experimentally determine the domain (if any), in which the chemical reaction is independent of mass transfer limitations deriving from liquid-solid resistances and from the solid porosity. Several experimental runs were performed with fixed amounts of catalyst of different average particle radius, R. The resulting concentration-time data were analyzed by means of a first-order kinetic model (i.e., rA ) k1[A]), in order to evaluate the equivalent first-order rate constant, k1, for the acid consumption. Subsequently, the values of k1 were plotted as a function of R, as shown in Figure 1. The asymptotic behavior shown at the smallest catalyst sizes allows us to
Figure 2. Concentration (mmol dm-3)-time (min) data obtained at pH ) 4.10, T ) 25 °C, and R ) 7.2 µm; the curves have been computed with the first-order kinetic model. (b) Oxalic acid; (O) ozone in the freeboard; (- - -) ozone in the liquid.
determine the kinetic domain (about R < 10 µm), in which the system behavior only depends on chemical kinetics, whereas the decreasing values of k1 measured at the largest values of R demonstrate the relevance of mass transfer limitations in this domain. Further evidence for the presence of mass-transfer limitations is given by the observation that a first order kinetics adequately fits the concentration-time curves only at the largest values of R. This result is consistent with the behavior expected for a reaction limited by the diffusion of oxalic acid, since mass-transfer phenomena linearly depend on the concentration of the diffusing species. On the contrary, at the smallest values of R, i.e., in the kinetic domain, this simple model produces nonrandom errors for both ozone and oxalic acid, as shown in Figure 2 for R ) 7.2 µm. In fact, the oxalic acid consumption is underestimated by the model for t < 15 min and overestimated for t > 15 min. Moreover, the model predicts a larger ozone concentration in the exit stream, in the early stages of the reaction. These errors are not so large to invalidate the procedure which allowed us to single out the kinetic domain. Nevertheless, it appeared advisable to perform a more detailed analysis of the system behavior inside the kinetic domain, in order to assess a more satisfactory kinetic model. Kinetic Analysis (a) The Kinetic Model. Kinetic modeling is performed on the basis of the following hypotheses. It is assumed that a fraction, fd, of the Mn(III) surface sites, responsible for the MnO2 catalytic activity toward oxalic acid ozonation (Andreozzi et al., 1996a), can undergo oxidative attack by ozone, leading to their irreversible deactivation. This reaction can account for the larger rate of consumption experimentally determined for ozone in the early stages of reaction (Figure 2). Moreover, this reaction also results in an increased reactivity of oxalic acid in the early stages of reaction (i.e., when most of the catalytic sites are still active) and in a decreased reactivity in the late stages (i.e., when a fraction of the sites has been deactivated). Further evidence for the presence of a deactivation mechanism is given by the results of some experimental runs, carried out with MnO2 particles previously taken
4776 Ind. Eng. Chem. Res., Vol. 36, No. 11, 1997 Table 1. Kinetic Parameters Determined from Experimental Runs Performed at Different Catalyst Sizesa R ) 2.5 µm model a 102k1, min-1 102kA, min-1 kd, dm3 mmol-1 min-1 fr σO3, % σA, %
R ) 7.2 µm
model b
model a
3.71 ( 0.13
model a
3.40 ( 0.15 5.50 ( 0.18 1.06 ( 0.06 0.505 ( 0.029 2.6 2.2
5.7 7.5
R ) 22 µm
model b
model b
3.14 ( 0.10 5.26 ( 0.23 0.685 ( 0.057 0.435 ( 0.036 2.9 4.7
6.1 11.0
4.60 ( 0.16 0.671 ( 0.050 0.514 ( 0.030 2.4 3.7
7.0 6.9
a m ) 50 mg, pH ) 4.10, T ) 25 °C, and k ) 6.57 × 10-3 min-1. Model a ) first-order kinetic model; model b ) “deactivable sites” S h kinetic model.
Table 2. Kinetic Parameters Determined from Experimental Runs Performed at Different pH Valuesa pH 102kA, min-1 kd, dm3 mmol-1 min-1 fr σO3, % σA, % 103kh, min-1 a
4.10
4.50
4.65
4.80
5.00
5.50
6.00
5.26 ( 0.23 0.685 ( 0.057 0.435 ( 0.036 2.9 4.7 6.57
3.70 ( 0.20 0.293 ( 0.059 0.584 ( 0.071 3.8 2.9 6.33
2.46 ( 0.13 0.191 ( 0.014 0.873 ( 0.099 2.9 2.5 6.17
2.04 ( 0.03
1.59 ( 0.03
1.22 ( 0.03
0.55 ( 0.03
1 1.9 3.2 6.01
1 2.4 3.0 5.86
1 2.8 2.9 5.42
1 4.7 2.9 2.42
mS ) 50 mg, R ) 7.2 µm, T ) 25 °C. “Deactivable sites” kinetic model.
in contact with bubbling ozone, in the absence of oxalic acid. In fact, this “ozonized” catalyst shows a smaller reactivity toward oxalic acid and a simple first-order kinetics for the acid consumption. When assuming a first-order dependence on ozone, the kinetics of sites deactivation kd
zO3 + Sd 98 deactivated site can be described by the following equation:
NS dfd/dt ) -kd[O3]LNSfd
(5)
At t ) 0, fd ) 1 - fr. In the equation, NS is the total number of active sites and the fraction of reversible sites (nondeactivable), fr, is treated as an adjustable parameter. A two step reaction mechanism is proposed for the oxalic acid oxidation slow
resulting from both chemical kinetics and MnO2 solubility thermodynamics:
kh ) k′h[Mn]sat
8 AS A+S9 k′ A
(7)
where k′h is the kinetic constant of the catalytic step and [Mn]sat is the saturation concentration of the catalytic Mn species. Therefore, the overall ozone consumption rate is given by
fast
AS + O3 98 products + S which leads to the following rate equation:
d[A] ) -k′A[A]NS(fr + fd) - kh[A] dt
Figure 3. Effect of reaction temperature on the rate constant, kh (min-1), for homogeneous ozonation of oxalic acid.
(6)
The rate constant kA ) k′ANS is determined by a comparison with the experimental data, whereas the rate constant kh, which accounts for the homogeneous catalytic ozonation of oxalic acid caused by dissolved Mn, has been evaluated from independent experiments performed at different temperatures and pH values. In those experiments, oxalic acid is submitted to ozonation in homogeneous aqueuos solutions previously kept into contact with MnO2 (at the selected temperature and pH), up to their saturation. The results show a first-order dependence of the reaction kinetics with respect to oxalic acid. The apparent rate constant, kh (whose values are reported in Tables 1 and 2), depends on pH and shows an Arrhenius-type dependence on temperature (Figure 3, Ea ) 24.5 kcal/mol), possibly
d[O3]L ) -kd[O3]Lfdβ - kA[A](fr + fd) - kh[A] (8) dt where the adjustable parameter β accounts for the number of active sites, NS, and for the stoichiometric coefficient, z, of ozone in the deactivation reaction. (b) Results and Discussion. The adequacy of the proposed kinetic model is firstly checked by the identification of the kinetic parameters pertaining to ozonation runs performed at constant values of pH, temperature, and catalyst load, while varying the catalyst particle size. The best-fit value β ) 0.74 mmol dm-3 was taken to be constant for all the analyzed data. As shown in Table 1, the kinetic model gives a very close correspondence between computed and experimental developments of the catalytic ozonation processes, at the smallest values of R. For the same runs, the first-order
Ind. Eng. Chem. Res., Vol. 36, No. 11, 1997 4777
active sites on the catalyst surface is regulated by the following equilibrium:
MOH2+ h MO- + 2H+
(9)
which can be quantitatively expressed in terms of the pH of zero charge for the catalyst, pHpzc. If one makes the likely assumption that only positive sites can be exploited for the adsorption of oxalic acid, then,
NS ) [MOH2+] )
Figure 4. Concentration (mmol dm-3)-time (min) data obtained at pH ) 4.10, T ) 25 °C, and R ) 7.2 µm; the curves have been computed with the “deactivable sites” kinetic model. (b) Oxalic acid; (O) ozone in the freeboard; (- - -) ozone in the liquid.
Nt[H+]2 [H+]2 + Keq
(10)
where Nt is the total number of charged surface sites (Nt ) [MOH2+] + [MO-]). According to the pHpzc definition (Noh and Schwarz, 1990), pKeq ) 2 × pHpzc. Thus, from the value (pHpzc ) 5.60) measured with the method of Subramanian et al. (1988) one obtains Keq ) 10-11.2 M2. On the other hand, the kinetic experiments show a remarkable increase of the system reactivity when pH is decreased well below the value pH ) 5.60. This suggests that oxalic acid adsorption on the catalyst surface mainly concerns HC2O4- ions. Since, in the investigated range of pH (4.1-6.0), the total (measured) concentration of oxalic acid [A] can be written as
[A] ) [HC2O4-] + [C2O42-]
(11)
the concentration [HC2O4-] of the effective reactant must be computed as
[H+] [HC2O4-] ) [A] + [H ] + K2 Figure 5. Effect of reaction temperature on the rate constants at pH ) 4.10. Particle size R ) 7.2 µm. (b) Catalyst deactivation, k ) kd; (O) adsorption, k ) kA.
kinetic model gives larger values of the average residual error between experimental and computed data, as shown by the relevant standard deviations for both the oxalic acid (σA) and ozone (σO3). As an example, Figure 4 shows the comparison between the experimental data plotted in Figure 2 and the concentration-time curves computed with the “deactivable sites” model. Data from catalytic ozonation experiments at varying reaction temperatures have also been submitted to the modeling procedure. From the Arrhenius plot of the kinetic constants referring to the adsorption step of oxalic acid on the catalyst surface and to the deactivation by ozone of catalyst active sites (Figure 5), the activation energies can be evaluated respectively as Ea ) 13.3 and 15.1 kcal/mol. These values give further confidence in the reliability of the proposed kinetic model. Mechanistic details concerning the adsorption step of oxalic acid on the catalyst surface are disclosed by the results obtained when applying the model to data from ozonation experiments performed at varying pH values. The results, reported in Table 2, show a marked influence of pH upon the adsorption step of oxalic acid, whereas the fraction of active sites stable toward ozone increases up to 1 for pH increasing up to 4.80. Consequently, parametric identification at pH > 4.80 is performed by excluding the catalyst deactivation term. The dependence on pH of kA can be tentatively explained by considering that the distribution of charged
(12)
where K2 is the second dissociation constant of oxalic acid (K2 ) 10-4.3 M). In conclusion, the term kA[A], which accounts for the heterogeneous catalysis in the rate eqs 6 and 8, depends on [H+]. From eqs 10 and 12, one obtains
kA[A] ) k′′A[HC2O4-]NS ) k′′A[A]NtFH+
(13)
where k′′A is a true kinetic constant, independent of pH, and
F H+ )
[H+]
[H+]2
[H+] + 10-4.3 [H+]2 + 10-11.2
(14)
The fairly good linear correlation between kA and FH+, shown in Figure 6, gives substantial support to the above data treatment. The residual deviations from the linearity might be tentatively attributed to the presence of neutral sites (MOH) on the catalyst surface, which obey the equilibrium
MOH h MO- + H+
(15)
whose effect is not included in eqs 13 and 14. Conclusions The ozonation of oxalic acid under the promoting influence of heterogeneous MnO2 catalyst is quantitatively described by a mathematical model which includes a multistep kinetic model. The rate of catalytic ozonation is controlled by the slow reaction step involving HC2O4- adsorption on catalytic active sites. A
4778 Ind. Eng. Chem. Res., Vol. 36, No. 11, 1997 in ) input L ) liquid 1 ) first order Greek Symbols R ) Ostwald coefficient β ) adjustable parameter σ ) percent standard deviation
Literature Cited
Figure 6. Apparent rate constant, kA, as a function of ln FH+ (eq 14).
fraction of the active surface sites can be destroyed by irreversible ozone attack at pH values < 4.8, whereas at pH > 4.8, no catalyst deactivation is observed. The influence of pH upon reaction kinetics indicates that positively charged active sites and HC2O4- ions are involved in the rate controlling adsorption step. In fact, the observed values can be coherently explained when taking into account the pHzpc value of the catalyst and the equilibrium constant of oxalic acid dissociation. Nomenclature A ) oxalic acid As ) adsorbed oxalic acid Ea ) activation energy, kcal mol-1 fd ) fraction of deactivable surface sites fr ) fraction of nondeactivable sites FH+ ) function defined in eq 14 k, k′, k′′ ) rate constants koLa ) volumetric gas-liquid coefficient of mass transfer, min-1 K2 ) second dissociation constant of oxalic acid, M Keq ) constant of equilibrium (9) mS ) amount of catalyst, mg Nr ) number of reactions included in the model NS ) number of active sites, mmol dm-3 Nt ) total number of charged sites, mmol dm-3 Q ) gas flow rate, dm3 min-1 r ) reaction rate, mmol dm-3 min-1 R) average particle radius, µm S ) catalytic site T ) temperature, K t ) time, min V ) volume, dm3 z ) stoichiometric coefficient Subscripts A ) oxalic acid O3 ) ozone B )bubbles d ) deactivation, deactivable F ) freeboard h ) homogeneous reaction
Al-Hayek, N.; Legube, B.; Dore’, M. Catalytic ozonationsFe(III)/ Al2O3sof phenol and its ozonation by products. Environ. Technol. Lett. 1989, 10, 415-426. Andreozzi, R.; Caprio, V.; D’Amore, M. G.; Insola, A.; Tufano, V. Analysis of Complex Reaction Networks in Gas-Liquid Systems: the Ozonation of 2- Hydroxypyridine in Aqueous Solutions. Ind. Eng. Chem. Res. 1991, 30, 2098-2104. Andreozzi, R.; Insola, A.; Caprio, V.; D’Amore, M. G. Quinoline Ozonation in Aqueous Solutions. Water Res. 1992, 26 (5), 639643. Andreozzi, R.; Caprio, V.; Insola, A.; Marotta, R.; Tufano, V. The Use of Manganese Dioxide as a Heterogeneous Catalyst for Oxalic Acid Ozonation in Aqueous Solution Appl. Catal. A: Gen. 1996a, 138, 75-81. Andreozzi, R.; Caprio, V.; Ermellino, I.; Insola, A.; Tufano, V. Ozone Solubility in Phosphate-Buffered Aqueous Solutions: Effect of Temperature, tert-Butyl Alcohol, and pH. Ind. Eng. Chem. Res. 1996b, 35 (4), 1467-1471. Astarita, G.; Savage, D. W.; Bisio, A. Mass Transfer with Chemical Reaction, John Wiley & Sons: New York, 1983. Campen, J. P. A Catalytic Ozonation Process, Which is Applicable to a Wide Variety of Mildly Polluted Waste Water Stream. Contribution to the 11th World Ozone Congress, S-16-1 to S-1612, San Francisco, Aug 29-Sept 3, 1993. Danckwerts, P. V. Gas-Liquid Reactions; McGraw-Hill: New York, 1970. Froment, G. F.; Bischoff, K. B. Chemical Reactor Analysis and Design; John Wiley & Sons: New York, 1979. Noh, J. S.; Schwarz, J. A. Estimation of Surface Ionization Constants for Amphoteric Solids. J. Colloid Interface Sci. 1990, 139, 139-148. Paillard, H.; Dore`, M.; Bourbigot, M. M. Prospects Concerning Applications of Catalytic Ozonation in Drinking Water Treatment. IOA 10th Ozone World Congress, Monaco, 19-21 March, 1991, Vol. 1, pp 314-329. Paillard, H.; Allemane, H.; Delouane, B.; Legube, B. Comparative Efficiency of Three Systems (O3, O3/H2O2 and O3/TiO2) for the Oxidation of Natural Organic Matter in Water. IOA Ozonews 1992, 20, 24-37. Subramanian, S.; Noh, J. S.; Schwarz, J. A. Determination of the Point of Zero Charge of Composite Oxides. J. Catal. 1988, 114, 433-439. Thompson, F. E.; Sharratt, P. N.; Hutchinson, J. Heterogeneous Catalytic Oxidation of Organic Pollutants in Aqueous Solutions. The 1995 ICHEME Research Event First European Conference, pp 297-299. Tufano, V.; Andreozzi, R.; Insola, A.; Caprio, V.; D’Amore, M. G. Optimal Operating Conditions for Lab-Scale Ozonation Reactors. Ozone: Sci. Eng. 1994, 16, 181-195.
Received for review March 11, 1996 Revised manuscript received June 3, 1997 Accepted June 19, 1997X IE960159D
X Abstract published in Advance ACS Abstracts, August 15, 1997.
