Wet Oxidation Kinetics of Refractory Low Molecular Mass Carboxylic

P.O. Box 537, SI-1001 Ljubljana, Slovenia. Wet oxidation kinetics of aqueous solutions of formic, acetic, oxalic, and glyoxalic acids was studied in a...
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Ind. Eng. Chem. Res. 1999, 38, 3830-3837

Wet Oxidation Kinetics of Refractory Low Molecular Mass Carboxylic Acids R. V. Shende† and Janez Levec*,†,‡ Laboratory for Catalysis and Chemical Reaction Engineering, National Institute of Chemistry, P.O. Box 30, SI-1001 Ljubljana, Slovenia, and Department of Chemical Engineering, University of Ljubljana, P.O. Box 537, SI-1001 Ljubljana, Slovenia

Wet oxidation kinetics of aqueous solutions of formic, acetic, oxalic, and glyoxalic acids was studied in a titanium autoclave at a temperature range of 150-320 °C and oxygen partial pressures between 8 and 60 bar. Oxidation reactions obeyed a first-order kinetics with respect to concentration of all substrates. On the basis of acid concentration decay, the activation energy for acetic, oxalic, and glyoxalic acid oxidation was 178, 137, and 97 kJ/mol, respectively; whereas on the total organic carbon (TOC) conversion basis, these values were slightly higher, namely 182, 141, and 104 kJ/mol. The activation energy for formic acid took a unique value of 149 kJ/ mol regardless of the type of concentration used. The rate of oxidation was proportional to a square root of oxygen concentration (partial pressure) for acetic, formic, and oxalic acids, whereas it was linearly proportional for glyoxalic acid. When sufficiently high oxygen partial pressure was applied (g22 bar), the individual acid conversion in a mixture of these acids was well predicted by the rate expression derived for that acid. The lumped TOC concentration of mixtures did not obey a first-order kinetic behavior, although underlying TOC kinetics for each individual acid was linear. The oxidation results are also discussed in a view of speculated reaction pathways and the reactor material. Introduction Oxidation of low molecular mass mono- and dibasic acids in aqueous solution has received major attention in the past few years, because they appear as the ultimate organic intermediates in oxidation schemes of many organic pollutants. They are refractory in nature, and their conversion into CO2 and H2O is considered as the rate-limiting step; therefore, their complete oxidation to CO2 may be prohibitively expensive. A group of these acids includes: glyoxalic, oxalic, propionic, acetic, and formic acids. Among these compounds, acetic acid is the most refractory and requires severe oxidation conditions (high temperature and high oxygen partial pressure) for its decomposition. Propionic acid is the second-most refractory compound, which can be produced in oxygen depleting zones of wet reactors during deep oxidation of some organic compounds.1 It further oxidizes partly into acetic and partly into formic acid.2,3 Glyoxalic and oxalic acids oxidize easily to more stable formic acid at relatively mild oxidation conditions. Besides appearing as intermediates in wet oxidation processes, low molecular mass carboxylic acids are valuable reagents in the production of commercial organic products. Therefore, they are frequently found in waste streams from petrochemical, pharmaceutical, textile-dye, caprolactum manufacturing, and synthetic polymer industries. Knowing the oxidation kinetics of these acids is thus inevitable for the optimal design of wet reactors, especially when considering that the construction material may be an expensive titanium alloy.4 The studies on wet oxidation of low molecular mass carboxylic acids have been summarized recently in detail.5 * Corresponding author. E-mail: [email protected]. † National Institute of Chemistry. ‡ University of Ljubljana.

The studies reported in the literature on noncatalytic oxidation of low molecular mass carboxylic acids are meager. The kinetic parameters provided by earlier investigators are summarized in Table 1. As one can see, the kinetic parameters disagree even though the similar material of reactor (SS-316) was used. Although the reported partial orders with respect to the substrates are mainly consistent, quite a large deviation exists with the oxygen order. If the oxidation takes place in the homogeneous (liquid) phase, one would expect a half-order or higher dependence (first order in glyoxalic acid) with respect to oxygen. Because the oxygen orders in Table 1 are generally lower than those expected, one may speculate that the kinetic parameters in Table 1 have been catalytically affected somehow by the reactorwall material. The higher activation energy found for propionic acid in a reactor made of titanium3 further supports the premise above, which also agrees with some other findings.6-8 These differences should not be neglected, particularly when the presence of chlorides in a wastewater (which causes stress-corrosion cracking) dictates the use of titanium as the construction material for the wet reactor. In many cases, the oxidation of parent compounds follows a first-order dependence, whereas for the lumped concentration parameter such as total organic carbon (TOC) or chemical oxygen demand (COD), the order is higher. A recent typical example is oxidation of Orange II, for which the disappearance rate is a first-order process.9,10 Refractory intermediate products such as acetic, formic, glycolic, and oxalic acids, formed during the course of oxidation, apparently slowed the reaction, and consequently mimicked the second-order behavior with respect to the lumped TOC concentration. On the other hand, the presence of acetic acid accelerates the oxidation of formic acid,5 and oxalic acid oxidizes with

10.1021/ie9902028 CCC: $18.00 © 1999 American Chemical Society Published on Web 09/03/1999

Ind. Eng. Chem. Res., Vol. 38, No. 10, 1999 3831 Table 1. Kinetic Parameters for Noncatalytic Oxidation of Low Molecular Mass Mono- and Dibasic Acids acid and reference acetica Foussard et al.15 propionic Day et al.2 Merchant16 Shende and Levec3 formica Foussard et al.15 Shende and Mahajani14 oxalic Foussard et al.15 Shende and Mahajani11 glyoxalic Shende and Mahajani11 a

initial conc. (g/L)

reactor type and material

temperature (°C)

oxygen pressure (MPa)

activation energy (kJ/mol)

acid order (m)

oxygen order (n)

∼30

batch SS-316

270-320

2.0-20b

167.7

1.0

0.37

7.4-14.8 1.0 (acid) 0.48 (TOC)

batch SS-316 batch SS-316 batch titanium

232-288 250-275 290-310

1.72-5.17 2.5-4.5

135 139 150 158

1.43 1.0 1.0 1.0

0.39 0.0 0.5 0.61

24-43 1.0 (COD)

batch SS-316 batch SS-316

190-313 150-240

2.0-20b 0.35-1.38

143.5 121.3

1.33 1.0

0.46 0.86

0.1-0.2 1.2 (COD)

batch SS-316 batch SS-316

227-288 225-245

2.0-20b 0.69-1.03

133.8 129.4

1.0 1.0

0.31 0.32

1.0-2.5 (acid) 0.43-1.08 (COD)

batch SS-316

150-200

0.35-1.38

53.5c 117d

1.0c 1.0d

0.92c 0.2d

As sodium salt. b Total pressure. c First step.

d

Second step.

appreciably higher rates when glyoxalic acid is present in a solution.11 Hence, in addition to providing a thorough oxidation kinetics of aqueous solutions of formic, acetic, oxalic, and glyoxalic acids, this work also aims to investigate a synergistic effect that may accompany the oxidation of mixtures of these acids. When an acid undergoes oxidation via intermediate products, the oxidation kinetics is also provided with respect to the lumped TOC concentration. Experimental Section Materials. All acids used in this study were obtained from Aldrich Chemical Co., USA, and used as received (purity >98%). Oxygen from a cylinder with a minimum purity of 99.5% was used for oxidation. Nitrogen from a cylinder of 99.99% was used whenever required. Apparatus and Experimental Procedure. All experiments were performed in a 2-L titanium (grade 4) autoclave (model 4572, Parr Instrument Co. Ltd., Illinois, USA), which was equipped with a magnetically driven turbine-type impeller (titanium); temperature, pressure, speed of agitation control units; and a liquidinjection vessel (SS-316; 200-mL capacity) mounted on the top of the autoclave. An upstream electronic mass flow controller (Brooks 5850) and a downstream electronic back pressure controller (Brooks 5866) were used to maintain constant operating pressure and constant delivery of oxygen (1.0 L/min) when the apparatus operated semicontinuously. In all experiments, before sparging oxygen, the reactor content was heated to the set temperature and a sample was withdrawn. This sample was referred as a “zero” time sample for kinetic measurements. However, a detailed experimental procedure can be found elsewhere.3 Mixtures with different acid concentrations (based on their lumped TOC concentration) were prepared and 50 mL of solution injected from the injection vessel directly into the autoclave. Analysis. Concentration of the acids was measured using a reversed-phase high-performance liquid chromatography (HPLC) (Hewlett-Packard 1100/DAD) system equipped with a Rheodyne 7725 injection valve with a 20-µL sample loop. Analysis was carried out using an SS-column of 300 × 8 mm inside diameter (i.d.), which consists of sulfonated cross-linked styrene-divinylbenzene copolymer as a stationary phase (‘Eurokat H’ from Geratebau Saulentechnik Eurochrom, Germany). A mobile phase of 0.01 N H2SO4 with a flow rate of 1 mL/

min was used for the analysis. The column was set at a temperature of 75 °C. Maximum absorption was attained at 214, 200, 209, 204 nm for glyoxalic, oxalic, formic, and acetic acids, respectively. By the analytical method used it was possible to evaluate quantitatively 1.0 mg of an acid in a liter of their aqueous solution mixture with a high degree of reproducibility. The initial concentration of acids in all oxidation experiments was 1000 mg/L; oxalic acid with 500 mg/L was an exception. Because the concentrations of acids were very low, the oxygen concentration in the liquid phase was assumed equal to the solubility of oxygen in water.12 The amount of TOC in samples collected during the oxidation experiments was measured by means of an advanced HTCO Rosemount/Dohrmann DC-190 TOC analyzer equipped with a nondispersive infrared CO2 detector. TOC was determined by subtracting inorganic carbon (IC) from total carbon (TC). The relative standard deviation of three different measurements never exceeded 1.5% for the range of TOC concentrations measured. Results and Discussion Wet oxidation is a heterogeneous reaction and consists of various steps taking place in a series at the macroscopic level: oxygen transfer across the gasliquid interface and chemical reaction in the liquid phase. The gas-side mass transfer resistance was estimated to be negligible because of the very high diffusivity of oxygen in the gas phase and its low solubility in water.12 To verify the absence of the liquid-side mass transfer resistance, the effect of the impeller speed on the rate of TOC reduction was studied separately for each of the acids. However, it was found that the rate of oxidation was independent of the impeller speed in a range of 1000-1500 rpm, indicating there were no limitations associated with the oxygen transfer. Low molecular mass carboxylic acids undergo at least four types of reactions: (i) reactions involving cleavage of the O-H bond, (ii) attack by a nucleophile (like oxygen-rich species) on the C-O bond, (iii) decarboxylation, and (iv) attack on R-carbon. In monocarboxylic acids, the attack of O2 and O• free radical on the C-O bond and also at the R-carbon can be considered as major reactions breaking the parent acid into a compound(s) with lower molecular mass and further into CO2 and water.

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Figure 1. First-order kinetic plot for formic acid oxidation based on TOC concentration measurements.

Figure 2. Arrhenius plot for oxidation of formic and acetic acids.

Oxidation of Formic Acid. In the present investigation only CO2 and H2O products were observed during oxidation of formic acid, even though the decarboxylation reaction (CO appearance) may occur under hydrothemal conditions.8 This finding is in agreement with others13 who reported that the decarboxylation of formic acid is totally negligible when oxygen concentration is very low (0.56 × 10-4 mol/L). Bjerre and Sorensen6 have observed CO during decomposition of formic acid at 260 °C, but the yield of CO2 was much higher than the yield of CO. However, we believe that in our case formic acid was oxidized directly into CO2 and water. Kinetics of formic acid oxidation was studied in a temperature range of 240-270 °C and with use of oxygen partial pressures between 8 and 20 bar. At less than 240 °C the conversion of formic acid was low, but it increased very rapidly with temperature. For instance, in a 2-h experiment at 240 °C and 8 bar of oxygen pressure, the conversion was 24%, whereas at 270 °C it reached 90%. However, the reaction was found to obey a first-order kinetics with respect to both the acid and the lumped TOC concentration. A first-order kinetics plot is shown in Figure 1 for the lumped TOC concentration. The Arrhenius plot is shown in Figure 2. The activation energy evaluated from the plot for formic acid is 149.2 kJ/mol, which is higher than that reported by earlier investigators (see Table 1). The effect of oxygen on the disappearance rate of formic acid was studied at 255 °C in a pressure range of 8-20 bar. It does not seem to be very significant; for

Figure 3. Comparison between experimental and calculated acid and TOC conversions for formic and acetic acids.

example, when the oxygen partial pressure was increased from 8 to 20 bar, the TOC reduction increased for only 31% in 2 h. The order with respect to oxygen was 0.5, which confirms the speculation above on the acid direct oxidation to CO2. The oxygen order reported by earlier investigators is obviously different (Table 1). Shende and Mahajani14 explained their finding by a hypothesis in which formic acid is attacked by both O• and OH• radicals. Because wet oxidation is a freeradical-mediated reaction, the order found in this investigation merely coincides with the oxygen stoichiometric coefficient for the complete oxidation of formic acid. Because of the value of the activation energy higher than that reported previously (Table 1), it is also speculated that oxidation of formic acid has not been influenced catalytically by the titanium reactor-wall surface. The kinetic data obtained can be correlated by the following rate expression.

-rTOC ) -rFA ) 4.48 × 1011 0.5(0.01 exp(-17949/T)C1.0 (1) FACO2

Figure 3 represents a parity plot of the TOC concentrations calculated by eq 1 and those obtained experimentally. The data points lay well inside the confidence limits of (10%. Oxidation of Acetic Acid. The kinetics of acetic acid oxidation was studied at temperatures between 300 and 320 °C with oxygen partial pressures from 25 to 60 bar. In a 3-h experiment, the concentration of acid was reduced 9% at 300 °C and 60 bar of oxygen partial pressure, whereas at 320 °C and 30 bar of oxygen pressure the reduction was 20%. Foussard et al.15 have studied wet oxidation of acetic acid (as sodium salt) in a SS-316 reactor between 270 and 320 °C and concluded that higher oxidation rates are expected above 320 °C. Merchant16 has observed that after a 5-h treatment at 275 °C, there was only 5% reduction in COD (corresponding to acetic acid). High stability of acetic acid is caused by the methyl group, which is an electron donor to the carbon having dO and -OH groups. During the course of acetic acid oxidation no intermediates (e.g., formic and oxalic acids) were detected by the HPLC analysis, which suggests that acetic acid oxidizes directly into CO2 and water. Throughout the entire temperature range, the rate of acetic acid oxidation was linearly proportional to the acid concentration.

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Figure 4. First-order kinetic plot for acetic acid oxidation based on acid concentration measurements.

A first-order kinetics plot is shown in Figure 4, and the Arrhenius plot is inserted in Figure 2. The activation energy was evaluated to be 178 kJ/mol, which is higher than that reported,15 where the reactor-wall material (SS-316) probably acted as a catalyst. On the other hand, much higher activation energy (197 kJ/mol) has also been reported for the initiation of acetic acid by molecular oxygen.17 The effect of oxygen partial pressure was studied across a range of 25-55 bar at 310 °C. A half-order dependence resulted, although the overall oxygen stoichiometry for the complete oxidation of acetic acid is 2. It is thus believed that relatively high oxidation temperatures lead to the formation of O• radicals,18 which further react with acetic acid molecules resulting in a half-order dependence with respect to oxygen. The experimental data for the disappearance rates of acetic acid can be presented in the following form: 0.5(0.04 -rAA ) 7.59 × 1010 exp(-21374/T)C1.0 AACO2

(2)

Although no stable intermediates were observed, slightly higher activation energy (182 kJ/mol) was found for TOC reduction (Figure 2). The experimental data for TOC reduction are well described by the kinetic equation 0.52(0.03 -rTOC ) 1.72 × 1011 exp(-21889/T)C1.0 TOCCO2 (3)

Comparison of the calculated concentrations (acid and TOC) with those measured experimentally is shown in Figure 3. All data points for acetic acid are also well inside the (10% confidence limits. It should be pointed out that no data exist in the literature for the activation energy for lumped parameters such as TOC and COD for noncatalytic oxidation of acetic acid. Oxidation of Oxalic Acid. Formic acid was observed as a major intermediate during oxidation of oxalic acid. It is known, however, that the decarboxylation temperature for pure acid is 180 °C.19 In this work, at 250 °C the thermal decomposition of the aqueous solution of oxalic acid yielded 52% of formic acid in 1 h, whereas at 300 °C with an initial concentration of 0.1 g/L it was almost completely decomposed to formic acid in 5 min. On the basis of these observations one can propose an overall reaction scheme such as represented by eq 4.

Thus, the decarboxylation and the attack of the C-O group are considered two major reaction routes in the decomposition of oxalic acid. Although the decarboxylation seems to be a very temperature-sensitive reaction, the stability and the reactivity of this compound may also depend on the chemical environment. For example, oxalic acid decomposes with higher rates in the presence of glyoxalic acid.11 According to the proposed reaction scheme (eq 4), oxalic acid undergoes thermal and oxidative decomposition simultaneously. Typical acid concentration profiles for thermal and combined decompositions are depicted in Figure 5 for different temperatures. The TOC reduction during oxidation of oxalic acid is not as fast as its concentration decay because of formic acid formation. The formation of formic acid at different temperatures and two oxygen partial pressures is depicted in Figure 6. Between 190 and 220 °C this formation is a first-order process. As one can see, the conversion of oxalic acid into formic acid depends also on the oxygen partial pressure. An experiment at 22 bar of oxygen partial pressure yielded less formic acid than that performed at 15 bar. This implies that high oxygen pressure favors oxidation of oxalic acid directly to CO2. Shende and Mahajani11 found the direct oxidation of oxalic acid to CO2 as the main reaction route, but they studied the oxidation in a SS autoclave and at slightly higher temperatures (210-225 °C). This difference in oxidation routes, however, can be ascribed to the reactor-wall material. Because in the titanium reactor formic acid oxidizes very slowly at 240 °C, one can always find it accumulated in solution below this temperature. The disappearance rate of oxalic acid was studied at temperatures between 220 and 250 °C and oxygen partial pressures between 8 and 22 bar. Although the acid disappearance rate obeyed a first-order kinetics with respect to the acid concentration in both decomposition routes, the order with respect to oxygen (studied at 220 °C) was 0.50. This indicates that oxygen was used for the oxidation of oxalic (and formic) acid according to its stoichiometric equation for direct oxidation to CO2 and water. The oxalic acid disappearance rate can be written as a sum of thermal and oxidative contributions, thus

-rOA ) 7.08 × 108 exp(-14313/T)C1.0 OA + 1.31 × 0.5(0.01 (5) 1011 exp(-16478/T)C1.0 OACO2

The first right term was obtained separately by fitting the thermal decomposition C versus time data (dashed curves in Figure 5). Equation 5 is represented in Figure 5 by solid curves. The Arrhenius plot with the activation energy of 137 kJ/mol for the oxalic acid oxidation is illustrated in Figure 8. Because formic acid was observed as an intermediate product during the oxidation of oxalic acid, the TOC versus time measurements were highly beneficial. For example, in a 2-h experiment at 220 °C and 8 bar of oxygen partial pressure, the TOC conversion was 67.7%, but 42% of carbon was found in formic acid. Because of accumulation of formic acid, the TOC conversions were always lower than those for oxalic acid. However, across

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Figure 5. Effect of temperature on oxalic acid thermal and oxidative decompositions.

Figure 6. Effect of temperature and oxygen partial pressure on formic acid formation during oxidation of oxalic acid.

Figure 8. Arrhenius plot for oxidation of oxalic and glyoxalic acids.

Figure 9. Comparison between experimental and calculated acid and TOC conversions for oxalic and glyoxalic acids.

on the TOC conversion is shown in Figure 7, and the Arrhenius plot with the activation energy of 141 kJ/mol for the TOC reduction is presented in Figure 8. The rate expression for the reduction of lumped TOC concentration can now be written as 0.47(0.03 -rTOC ) 4.37 × 1011 exp(-16999/T)C1.0 (6) OACO2

Figure 7. First-order kinetic plot for oxalic acid oxidation based on TOC concentration measurements.

the entire temperature range the oxidation obeyed a first-order kinetics with respect to the total organic carbon. When the oxygen partial pressure was increased from 8 to 15 bar, the TOC reduction increased from 58 to 72% after 2 h of treatment. The oxygen order was found to be slightly lower (0.47) than that for the acid oxidation; one can speculate that it may be because of the thermal decomposition. (One should remember that during the decarboxylation reaction TOC is also reduced because of the release of CO2.) A first-order plot based

A comparison of the experimental results with those calculated using eqs 5 and 6 is presented in Figure 9. Good agreement can be seen between the calculated and measured values. The same conversions for the acid and TOC were found when the oxidation was carried out in the semicontinuous mode of operation. Oxidation of Glyoxalic Acid. Among the low molecular mass acids, glyoxalic acid is the least refractory acid, which oxidizes well below 200 °C. Our experiments revealed that during oxidation of glyoxalic acid, formic and oxalic acids are major intermediates. The normalized concentration profiles of these intermediates and mother compound are shown in Figure 10 for a typical experiment performed at 165 °C and 8 bar of oxygen partial pressure. It can be seen that formic acid concentration increases almost linearly with time, whereas oxalic acid concentration is kept fairly constant but at a very low value. At higher oxygen partial pressure (15 bar) the concentration of oxalic acid was increased.

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When temperature was increased to 180 °C (at oxygen pressure of 8 bar) more oxalic acid was formed, but above 180 °C its amount decreased again because the rate of oxalic acid oxidation became appreciable. The rate of formic acid formation increased with temperature. At 200 °C and 8 bar of oxygen, for example, the concentration of formic acid reached a maximum at 0.62 g/L (starting with 1.0 g/L of glyoxalic acid) and remain unchanged with time after glyoxalic acid approached the zero value. In these conditions thermal decomposition of oxalic acid into formic acid is not a very fast reaction (see the previous section), therefore it is not likely that, appreciable formic acid arrived via oxalic acid. Because the concentration of oxalic acid was certainly below the limits of detection, one can conclude that in these oxidation conditions glyoxalic acid was transformed mainly into formic acid. A good agreement between the experimental TOC measurements and TOC calculated by adding up the carbon present in all three acids (GA, OA, FA) shown in Figure 10 also favors the conclusion above. In an attempt to reduce the concentration of intermediates at 200 °C, oxygen partial pressure was increased to 22 bar. As a result of this pressure increase, the amount of formic acid was reduced about two thirds, whereas oxalic acid was totally absent. In these conditions, oxalic acid also oxidizes very rapidly to CO2. Based on this fact, one can speculate that at high temperatures and pressures a major portion of glyoxalic acid oxidizes directly into carbon dioxide and less into formic and glyoxalic acids. It is not possible to discriminate quantitatively among the last two routes on the basis of the present data. However, eq 7 represents possible reaction routes of glyoxalic acid as well as intermediates formed during the oxidation. In the reaction pathway thermal decomposition of oxalic acid into formic acid is also accounted for, whereas the thermal decomposition of glyoxalic acid into formic acid (and CO) is ignored because it was confirmed experimentally to be negligible at the conditions used. At high oxygen partial pressures it is thus believed that direct oxidation into carbon dioxide is the prevailing reaction of the glyoxalic acid oxidation.

The oxidation kinetics of glyoxalic acid was studied between 150 and 200 °C and oxygen partial pressures between 8 and 18 bar. The fact that glyoxalic acid almost completely disappeared in about 10 min at 200 °C and 8 bar of oxygen pressure indicates that its oxidative decomposition is relatively fast; it is a firstorder reaction with respect to the acid concentration. The effect of temperature on the glyoxalic acid decay is shown in Figure 11. The Arrhenius plot with the activation energy of 97.4 kJ/mol is depicted in Figure 8. The effect of oxygen partial pressure was studied across a range of 8-18 bar at 165 °C; it also resulted in a first-order dependence. Hence, the kinetics for the

Figure 10. Concentration of intermediates observed during oxidation of glyoxalic acid.

Figure 11. Concentration of glyoxalic acid as a function of time at different temperatures.

oxidative decomposition of glyoxalic acid can be presented in the form 1.0 -rGA ) 2.25 × 109 exp(-11719/T)C1.0 GACO2

(8)

where the rate constant accounts for all three oxidative routes indicated by eq 7. Because glyoxalic acid undergoes oxidation into CO2 via some intermediates, for design purposes it is more convenient to present the decomposition kinetics in terms of its lumped TOC parameter. It was found, however, that the rate of total organic carbon disappearance (TOC) is also a first-order process in the entire temperature range investigated (Figure 12). Even though the conversion of glyoxalic acid at 200 °C was almost complete, because of its conversion into formic and oxalic acid TOC was reduced for 44% only. The activation energy based on the TOC conversions was 104 kJ/mol (Figure 8). The kinetic data obtained with the TOC measurements are well correlated by eq 9. 0.92(0.02 (9) -rTOC ) 6.25 × 109 exp(-12549/T)C1.0 TOCCO2

The order with respect to the oxygen concentration is again slightly different from that obtained with the acid concentration measurements (eq 8), as with oxalic acid. The difference is caused because some intermediates may also decompose thermally (i.e., releasing CO2 but not consuming oxygen). The acid and TOC

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Figure 12. First-order kinetic plot for glyoxalic acid based on TOC concentration measurements.

conversions calculated by eqs 8 and 9, respectively, are compared with those obtained experimentally in Figure 9. No difference in the conversion of glyoxalic acid was observed whether the oxidation was carried out in batch or semicontinuous mode of operation. Nevertheless, the yield of oxalic acid was slightly higher when oxidation was performed semicontinuously. For example, in a 2-h semicontinuous experiment at 165 °C and 8 bar of oxygen pressure, the oxalic acid concentration was 0.113 g/L, whereas it was only 0.075 g/L in a batch experiment. There is no rational explanation for this difference, however. Synergistic Effect. To explore the synergy among the acids, a few oxidation experiments were made with aqueous solutions in which acids were mixed in different ratios. For example, a 50% total carbon-based solution, prepared from equimolar solutions of acetic and formic acids, respectively, was oxidized at 265 °C and 8 bar of oxygen pressure. The conversion of formic acid was about 11% higher than the conversion of formic acid being oxidized alone in the same operating conditions. On the other hand, no difference in the conversion of formic acid was found when oxygen pressure was increased up to 22 bar. Similar synergy was obtained for a mixture of glyoxalic and oxalic acids; again, when oxygen partial pressure of 22 bar was applied (at 200 °C) no rate enhancement for oxalic acid was observed. To shed more light on the synergy, the total carbonbased equimolar mixture of formic, acetic, oxalic, and glyoxalic acids was prepared and oxidized at four different temperatures, namely 170, 220, 265, and 310 °C. Oxygen partial pressure was between 22 and 55 bar. It should be pointed out that different temperatures were dictated by a wide range of the oxidation rate constants. The experimental conversions for each individual acid in the mixture are plotted in Figure 13 as data points, and the solid curves represent their concentration profiles predicted by the oxidation rate expressions for each acid alone. Because a good agreement is demonstrated, one can conclude that there is no synergistic effect as long as the oxygen partial pressure exceeds 22 bar. As demonstrated in this work for oxalic and glyoxalic acid, the complete oxidation of all intermediates is accomplished a long time after the original compound disappears. Therefore the rate equations that predict the total carbon reduction are much more valuable than

Figure 13. Comparison between individual experimental acid concentrations in mixture of formic, acetic, oxalic, and glyoxalic acids and those predicted by respective kinetic expressions.

Figure 14. Comparison between experimental TOC concentrations in mixture of formic (F), acetic (A), oxalic (O), glyoxalic (G), and propoinic (P) acids and those calculated by adding up TOC from respective kinetic expressions.

those for mother compounds. To see how well one can predict the TOC reduction in a mixture, two different carbon-based equimolar solutions of formic, acetic, propionic, oxalic, and glyoxalic acids were prepared and treated at 300 ˚C and 50 bar of oxygen pressure. The TOC reductions during oxidation of these two mixtures are plotted in Figure 14 as a function of time. Two important conclusions can be drawn from these results: (i) a slight disagreement exists between the experimental TOC and that obtained from adding up the TOC predicted by the TOC rate expressions for each acid, and (ii) the best fits to the experimental as well as to the predicted points exhibit a second-order kinetic behavior, although the underlying TOC kinetics for each of the acids was a first-order rate process. It seems that more refractory acids, such as acetic, propionic, and formic acids, slowed the oxidation reactions and consequently mimicked a second-order behavior with respect to the lumped TOC concentration. It is quite obvious that no difference between the measured and calculated TOC would be observed if all acids oxidize directly to carbon dioxide and water. Thus, this difference accounts for the partial oxidation and thermal decomposition (resulting in CO2 release) of all species present in the mixture. The small difference found in this work further persuades us to conclude that the acids mainly undergo

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oxidation directly to carbon dioxide and water. On the other hand, it is quite possible that the difference would not be the same with a different ratio of acids, nor when oxidation is taking place in a SS-316 reactor. Conclusion On the basis of this study and the previous one for propionic acid,3 the oxidative refractoriness of low molecular weight aliphatic acids can be arranged in the following order: acetic acid > propoinic acid . formic acid > oxalic acid > glyoxalic acid. Oxidation of formic and acetic acids has not yielded stable intermediates, whereas formic acid is a major intermediate product during oxidation of oxalic and glyoxalic acids. Therefore, the disappearance rates of formic and acetic acids coincide with their TOC reduction rates, whereas the rates for oxalic and glyoxalic acids differ notably. The disappearance rate for the lumped organic carbon (TOC) concentration in a mixture of acids obeys a second-order behavior, although the underlying kinetics is linear. Because the activation energies found in this work are higher than those obtained in SS autoclaves, one can conclude that titanium is a fairly inert material. Although the observed oxygen order matches well with the stoichiometric coefficients for complete oxidation of formic, oxalic, and glyoxalic acids, its value may differ for oxidation of other compounds, because wet oxidation is a nonelementary reaction involving the free-radical mechanism. Thus, the oxygen order of 0.5 found for acetic acid might be attributed to the formation and reaction of O• radicals. The oxidation selectivity toward the formation of carbon dioxide increases with oxygen partial pressure. In a titanium autoclave no synergistic effect is exhibited in a mixture of acids if prepared on the equimolar basis as far as oxygen partial pressure exceeds 22 bar. Because titanium-made reactors have been requested more frequently in treating a variety of industrial aqueous streams, the kinetic expressions provided in this work for the lumped TOC parameter will assist in their efficient design. Acknowledgment The authors acknowledge support from the Slovenian Ministry of Science and Technology under Grant J20783. R. V. Shende thanks the National Institute of Chemistry for the fellowship aid. Nomenclature CA ) concentration of acid, mg/L CO2 ) oxygen concentration, mol L-1 CTOC ) total organic carbon (TOC) concentration, mg L-1 E ) energy of activation, kJ/mol k ) rate constant, L(m+n)-1 mg(1-m) mol-n s-1 m ) order with respect to acid or TOC n ) order with respect to oxygen P(O2) ) oxygen partial pressure, bar -ri ) rate of oxidation, mg L-1 s-1 R ) gas constant, ) 8.314 J mol-1 K-1

t ) real time, s T ) temperature, K Subscripts AA ) acetic acid FA ) formic acid GA ) glyoxalic acid OA ) oxalic acid PA ) propionic acid

Literature Cited (1) Devlin, H. R.; Harris, I. J., Mechanism of Oxidation of Aqueous Phenol With Dissolved Oxygen. Ind. Eng. Chem. Fundam. 1984, 23, 387-392. (2) Day, D. C.; Hudgins, R. R.; Silveston, P. L. Oxidation of Propionic Acid Solutions. Can. J. Chem. Eng. 1973, 51, 733-740. (3) Shende, R. V.; Levec, J. Kinetics of Wet Oxidation of Propionic and 3-Hydroxypropionic Acids. Ind. Eng. Chem. Res. 1999, 38, 2557-2563. (4) Levec, J. Wet Oxidation Processes for Treating Industrial Wastewaters. Chem. Biochem. Eng. Q. 1997, 11, 47-58. (5) Mishra, V. S.; Mahajani, V. V.; Joshi, J. B. Wet Air Oxidation. Ind. Eng. Chem. Res. 1995, 34, 2-48. (6) Bjerre, A. B.; Sorensen, E. Thermal Decomposition of Dilute Aqueous Formic Acid Solutions. Ind. Eng. Chem. Res. 1996, 31, 1574-1577. (7) Brill, T. B.; Schoppelrei, J. W.; Maiella, P. G.; Belsky, A. A Spectrokinetics of Solvo-Thermal Reactions. Proc. Int. Conf. Solvotherm. Reactions, 2nd. 1996, 5-8. (8) Yu, J.; Savage, P. E. Decomposition of Formic Acid under Hydrothermal Conditions. Ind. Eng. Chem. Res. 1998, 37, 2-10. (9) Donlagic, J.; Levec, J. Oxidation of an Azo Dye in Subcritical Aqueous Solutions. Ind. Eng. Chem. Res. 1997, 36, 3480-3486. (10) Donlagic, J.; Levec, J. Comparison of Catalyzed and Noncatalyzed Oxidation of Azo Dye and Effect on Biodegradability. Environ. Sci. Technol. 1998, 32, 1294-1302. (11) Shende, R. V.; Mahajani, V. V. Kinetics of Wet Oxidation of Glyoxalic Acid and Oxalic Acid. Ind. Eng. Chem. Res. 1994, 33, 3125-3130. (12) Crammer, S. D. The Solubility of Oxygen in Brines from 0 to 300 °C. Ind. Eng. Process Des. Dev. 1980, 19, 300-305. (13) Baldi, G.; Goto, S.; Chow, C. K.; Smith, J. M. Catalytic Oxidation of Formic Acid in Water-Intraparticle Diffusion in Liquid Filled Pores. Ind. Eng. Chem. Process Des. Dev. 1974, 13, 447-452. (14) Shende, R. V.; Mahajani, V. V. Kinetics of Wet Oxidation of Formic Acid and Acetic Acid. Ind. Eng. Chem. Res. 1997, 36, 4809-4814. (15) Foussard, J. N.; Debellefontaine, H.; Besombes, V. J. Efficient Elimination of Organic Liquid Wastes: Wet Air Oxidation. J. Environ. Eng. 1989, 115, 367-385. (16) Merchant, K. P. Studies in Heterogeneous Reactions. Ph.D. Thesis, Department of Chemical Technology, University of Bombay, India, 1992. (17) Boock, L. T.; Klein, M. T. Lumping Strategy for Modeling the Oxidation of C1-C3 Alcohols and Acetic Acid in HighTemperature Water. Ind. Eng. Chem. Res. 1993, 32, 2464-2473. (18) Joglekar, H. S.; Samant, S. D.; Joshi, J. B. Kinetics of Wet Air Oxidation of Phenol and Substituted Phenols. Water Res. 1991, 25, 135-145. (19) Cornils, B.; Lappe, P. Dicarboxylic acids, Aliphatic. In Ullmann’s Encyclopaedia of Industrial Chemistry; VCH: Weinheim, 1989; Vol. A8.

Received for review March 19, 1999 Revised manuscript received July 6, 1999 Accepted July 17, 1999 IE9902028