Ind. Eng. Chem. Res. 1987,26, 2401-2403
2401
Catalyzed Oxidation Kinetics of Anthracene with Oxygen in Ethylene Glycol Emilio A. Cepeda Departamento Quimica Tgcnica, Colegio Universitario de Alaua, Vitoria, Alava, S p a i n
Mario Diaz* Departamento Quimica Tgcnica, Universidad del Pais Vasco, Lejona, Vizcaya, S p a i n
The kinetics of liquid-phase oxidation of anthracene to anthraquinone by gaseous oxygen, with CuBrz as a catalyst and ethylene glycol as the solvent, has been studied. The oxidation kinetics, without an induction period, is interpreted by considering an analyzed intermediate compound, and the kinetic constants are evaluated over the temperature range 120-160 O C . One mechanism for the reaction which conforms with the experimental results is also proposed. In recent years, the interest in the development of new processes using crude coal tar as feedstock has received new support. In some cases, new specific compounds have been looked for, while in others classical products and processes previous to the development of petroleum technology have been reinvestigated more thoroughly with a view to their scientific and technical possibilities in modern markets. One of these products, anthraquinone, is very important in dye-stuffs manufacture, more recently in the wood pulp industry which has good possibilities for future demand. Moreover, anthraquinone cam be obtained in a feasible form by oxidation from purifed cake anthracene (Collin and Mildenberg, 1978). Studies of this oxidation have been carried out in the gas phase with several catalysts, e.g., with vanadium (Ghosh and Nair, 1980; Sen et al., 1979). The possible economic advantages of mild conditions and lower temperatures working in the liquid-phase catalyst have promoted different studies and processes, as with chromates (Henglein, 1963; Weissermel and Arpe, 1978), with nitric acid (Kumar and Sekhar, 1982), and with metallic salt catalysts Ce(1V) (Rindone and Scolastico, 1971), Cu(I1) (Brossard et al., 1977; Koshitani et al., 19821, or metallic complexes (Muller and Bobillier, 1981). Some metals like copper seem to be good catalysts in the liquid phase, but there are no practical studies on these reactions. Brossard et al. (1977) identified an intermediate of the reaction but gave little experimental data and made no kinetic study. In this paper, we study the Br2Cu-catalyzed oxidation of anthracene by air and ethylene glycol as a solvent, to clarify the nature and regime of the reaction from the kinetic data of the overall process, and on this basis we propose a mechanism of the reaction.
Experimental Section Materials Specifications. Anthracene (95%) obtained from Fluka was purified by the method of Takeuchi and Furusawa (1965) to 99+% purity. CuBr2,ethylene glycol, and other chemicals were pure reagents from a reputable firm. Cylinder (99.95%) oxygen was employed. Apparatus. The reaction was carried out in a jacketed cylindrical glass reactor containing 250 cm3 of solution. This 7-cm-diameter reactor was fitted with a six-bladed turbine impeller of 3.5-cm diameter placed one-third from the bottom to the gas-liquid interface and four 7-mm symmetrically placed baffles. The gas was fed to the reactor through a tube sparger located beneath the impeller. Temperature in the reactor was controlled with 0.1 "C nominal precision by means of polyethyleneglycol 400
Scheme I
(5')
circulating in the jacket from a bath. The gas was previously saturated in ethyleneglycol at the same temperature. Reaction Procedure. Anthracene in ethylene glycol was placed in the reactor and thermostated to the reaction temperature while stirred. The catalyst was added, the gas was fed in, and the reaction time was set initially to zero. Samples of 1 mL were periodically withdrawn for analysis. In the analytical method that we have developed, the sample was dissolved in 50 mL of ethyl acetate and analyzed by LLC with a column of Polyglosil C8, 10 pm, with a UV detector (540 nm) and methanol-water (65/35) mobile phase.
Results and Discussion The final reaction product is anthraquinone (AQ) and the intermediate is 9-oxo-9,lO-dehydroanthracene-10-
0888-5885/87/2626-24Ql$Q1.5O/Q 0 1987 American Chemical Society
2402 Ind. Eng. Chem. Res., Vol. 26, No. 12, 1987 T.140'C
/
~
0
10
3
!(hr)
Figure 1. Products and kinetic curves during reaction.
I
LAN1.5
518 1 6 3 g - m o l / L
LCuBr21=1 78 1639-moi/L
9
0 A.'", 0
,
2
,
,
,
,
5
4
, -A 8
10
12
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Figure 3. Fitting of kinetic results to first-order kinetics for reaction with several initial anthracene concentrations.
i = 140'C 02=51$L/5
16Bg-mol/ L s
30 1 20
,
-
. I
"I
1
111
I
-
I
Y !I
01 0
// 0
43
1
,
I
16
24
32
1 40 16$-mol/i lCu(ll)l
15
20
25
rps
Figure 2. Effect of agitation rates on the first and second reaction rates.
spiro-Pdioxalane (I) (Scheme I, eq 9) as we have identified by IR and RMN and which has also been described by Brossard et al. (1977). No other products were detected. The type of kinetic curves being shown are in Figure 1. The small variation in the visible spectrum after the beginning of the reaction seems to show the absence of changes in the form and amount of species; copper(II) and its complexes should be the main components of the catalyst. Kinetic Studies. Reaction regimes in gas-liquid reactions in stirred tanks are frequently identified from experiments with different agitation rates. In this way, the kLa values continuously increase with agitation up to 60 rps at least, and the kinetic regime is deduced if no dependence in the observed rate is obtained. In Figure 2 the observed rates for I
AN-I-AQ
H
are shown for different yields and agitation rates. We carried out the experiments at 20 rps, thus avoiding the diffusional resistances, with the organic phase saturated with oxygen. The observed rates could be adjusted based on the concentration of the reagents in the liquid phase: anthracene and the intermediate product for the first and second reactions, respectively. Different experiments with initial concentrations of anthracene from 1to 4 g/L were carried out at 140 "C to test the equation validity range and the possibility of an induction period. The results in Figure 3 show the correlation for the first-order dependence in IAN1 of the first reaction, KI = 6.9 X s-l in the complete range of anthracene concentrations. In the same way, the values of the rate constant for the second
Figure 4. Catalyst concentration effects on the kinetic rate constants.
reaction, kH = 3.0 X s-l, are computed by the Himmelblau-Jones-Bischoff method (1967), being also constant in the AN concentration range studied. No kinetic data are available in the literature, but we have results (Brossard et al., 1977; Muller and Bobillier, 1981; Koshitani et al., 1982) of the yield at certain temperatures and times of the reaction for anthracene oxidation in several solvents (dioxane, benzene/TBHP, ethylene glycol, diglyme) and catalysts (Br2Cu(OX), RhC1(PPh3), RhCO(PPh,),Mo(CO)J. From these data, we have calculated in an approximate way the first-order rate constants. These values are of the same order or are one order of magnitude smaller than the results from our experiments. In some of these, diffusional resistances might play an important role, giving smaller kinetic constants; Brossard et al. give with Br2Cu catalyst a value of kI = 1.46 X lo4 s-l, Koshitani et al. a value of 3 X s-l, and Muller and Bobillier values in the range 10-5-104 s-l for several solvents and catalysts. No data are available for the second reaction. The dependence on the observed reaction rate by the catalyst concentration was studied with copper(II) bromide concentrations ranging from 8 to 40 X lo4 g-mol/L. The effects on kI are shown in Figure 4. The catalyst has no influence on the second reaction, while for the first K,' = KI/Qn, n = 1. We carried out experiments with oxygen instead of air in the gas phase to find out the kinetic order in O2 and to discuss the mechanism of the reaction. The lack of effect in the kinetic constants shows that there is no dependence of O2 concentration in reactions rI and rH. Finally, experiments in the temperature range 120-160 "C were carried out, and the rate constants kIand kH were fitted to the Arrhenius equation: kI = 2.11 X lo8 x exp(-18235/R7') L/(s.mol) and KH = 1.76 X lo9 x
Ind. E n g . Chem. Res. exp(-25943/RT) s-l. The high apparent activation energies common in chemical reactions confirm the assumption of a kinetic regime in this gas-liquid reaction under our experimental conditions. The complete solubility of the reagents and products in the solvent was assured. The high yield and selectivity for this reaction, e.g., 98% at 158 OC, make its application possible for industrial use. The low solubility of reagents and products is nevertheless a problem that must be treated. Solubility data in this system and with other solvents as well as kinetic data in the presence of the solid phase must be obtained.
Mechanism Approximation Several authon have suggested mechanisms for this type of oxidation; in every case they look complex and not strictly decided. Koshitani et al. (1982) gave a mechanism through anthrone formation that resulted in impure product. No intermediate acetal is obtained in this system (acetic acid solvent), so it cannot be extrapolated to our experiments. Brossard et al. (1977) obtained the acetal with possible evolution, but no detailed steps for the reaction were given. Rindone and Scolastico (1971) in reactions with cerium catalyst suggested the existence of a carbonium ion that could be easily attacked by the solvent. Thus, from these contributions and on the basis of the kinetic results obtained, for the dependence of reagents and catalyst, we propose a possible mechanism (Scheme 1). Mechanism. This is based on two steps (reactions 1 and 2) to form 9-anthryl radical by Cu2+catalyst attack (Wilk et al., 1966;Andrulis et al., 1966),as has been shown in some of these types of reactions. Then by typical propagation reactions 3-7 (Koshitani et al., 1982; Lyons, 1977), a 9-anthronyl radical is produced, where Cu(I), Cu(II), and AN are involved. Step 5' might be viewed as a resonance hybrid. On the basis of other works (Gill, 1970),we thought that 9-anthronyl radical might react with Cu(II) to give an 9-anthronyl cation (eq 8). This compound could react with the solvent (eq 9) to obtain the intermediate product I. The solvent has then an important role in the mechanism, as it is obtained. Finally the intermediate gives rise to anthraquinone by means of a hydrolysis equation.
1987,26, 2403-2408
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This mechanism can be used according to a mechanistic approach and the steady-state principle, for the intermediate production rate, considering reaction 10 as irreversible,
This equation agrees with experimental data; there is no dependence on O2 concentration, first order on Cu2+ and AN in the first reaction, but zero order in the second reaction. Registry No. I, 64420-66-2; C a r 2 ,7789-45-9; HO(CH2)20H, 107-21-1; anthracene, 120-12-7; anthraquinone, 84-65-1; 9H-antacenyl radical, 58194-37-9; (anthracen-9-y1)peroxy radical, 109528-86-1; (anthracen-9-yl)hydrogenperoxide,109528-87-2; (anthracen-9-y1)oxy radical, 24690-75-3; bis( 1,2-ethanediolatoO,O')copper(II), 62011-09-0.
Literature Cited Andrulis, P. J.; Dewar, M. J. S.; Dietz, R.; Hunt, R. L. J. Am. Chem. SOC.1966,88, 5473. Brossard, J.; Janin,R.; Krumenaker, L.; Varagnat J. J. Tetrahedron Lett. 1977, 26. 2273. Collin, G.; Mildenberg, R. Chem. Ind. 1978, 567. Ghosh, A. K.; Nair, G. S. B. Indian J. Technol. 1980,18, 181. Gill, G. B. In Modern Reactions i n Organic Synthesis; Timmons, C. J., Ed.; Van Nostrand Reinholtd: London, 1970. Himmelblau, D. M.; Jones, C. R.; Bischoff, F. B. Ind. Eng. Chem. Process Des. Dev. 1967, 6, 536. Henglein, R. A. Grundrissder Chemischen Tecnik; Verlag Chemie: Weinheim, 1963. Koshitani, J.; Kado, T.; Ueno, Y.; Yoshida, T. J. Org. Chem. 1982, 47, 2879. Kumar, D. C.; Sekhar, D. N. J.Chem. Technol. BiotechnoL 1982,32, 643. Lyons, J. E. In Fundamental Research i n Homogeneous Catalysis; Tsutsui, M., Ugo, R., Eds.; Plenum: New York, 1977. Muller, P.; Bobillier, C. Tetrahedron Lett. 1981,22, 5157. Rindone, B.; Scolastico, C. J. Chem. SOC.1971, 2238. Sen, D. K.; Prassard, D. B.; Nair, C. S. R. J.Indian Chem. SOC.1979, 56, 898. Takeuchi, T.; Furusawa, M. Kogio Kagaku Zasshi 1965, 68, 474. Weissermel, K.; Arpe, H. J. Industrial Organic Chemistry; Verlag Chemie: Weinhein, 1978. Wilk, M.; Bez, W.; Rochlitz, J. Tetrahedron Lett. 1966, 22, 2599.
Received for review February 28, 1986 Revised manuscript received March 5, 1987 Accepted June 20,1987
Optimization of the Preparation of a Catalyst under Deactivation. 1. Control of Its Kinetic Behavior by Electing the Preparation Conditions Andrbs T. Aguayo,* Jose M. Arandes, Arturo Romero, and Javier Bilbao Departamento Qulmica Tgcnica, Universidad del Pais Vasco, A p d o . 644, 48080 Bilbao, S p a i n
The relation of the physical properties and surface acidity of silica-alumina catalysts with their preparation conditions by impregnation is studied. It is observed that the catalysts properties are the result of choosing the values of A1,03 content, temperature, and drying pressure of the silica gel support and calcination temperature of silica-alumina. From a study of the kinetic behavior of 61 silica-alumina catalysts for 2-ethylhexanol dehydration, the relation of conversion at zero time and of a function that represents the catalyst deactivation with the physical properties and surface acidity of the catalysts has been determined. To optimize the preparation of a catalyst being used in an industrial process that is under deactivation, the fol-
lowing steps must be studied (1) the relationship between the physical and chemical properties and the preparation
0888-588518712626-2403$01.50/00 1987 American Chemical Society
