Catalyst Activity. A catalyst induction period was observed for most of the tests. In general, the highest activit y of the catalyst was not obtained until after about 2 hr on stream. The activity then declined with time on stream. A typical conversion-time curve is shown in Figure 1. The data presented in Tables I-IV are averages of several samples collected during a period, generally 2 hr, when the butadiene conversions were the highest. Catalyst Deposits. At the relatively high temperature (500-550°C) used to disproportionate butadiene a t high conversions, a high rate of coke laydown on the catalyst would normally be expected. As shown in Table I, the coke amounted to 1-2% of the converted butadiene; this is equivalent to a rate of coke deposition on the catalyst of 5-10% per hour. Cyclohexene from Butadiene. Cyclohexadiene can be selectively hydrogenated to cyclohexene. This was demonstrated in a 10-hr cycle using a nickel-based selective hydrogenation catalyst. Conversion of cyclohexadiene and selectivity to cyclohexene were both near 10070. Combining disproportionation of butadiene to cyclohexadiene with selective hydrogenation of cyclohexadiene to cyclohexene provides a route for synthesizing cyclohexene from butadiene.
Summary The postulated reaction paths for the synthesis of cyclohexene, cyclohexadiene, and benzene from butadiene are summarized in Figure 2. The relative yields of cyclohexadiene and of benzene are a function of process conditions. At high conversion levels (i.e., >35%), the highest yield of cyclohexadiene obtained was at 515°C with catalyst containing 3% tungsten oxide, and the highest yield of aromatics was a t 550°C with 16% tungsten oxide catalyst (Table V). Cyclohexadiene was hydrogenated a t near 100% selectivity to cyclohexene.
Literature Cited Banks, R . L.. Bailey, G . C., ind. Eng. Chem.. Prod. Res. Deveiop. 3, 170 ( 1964 j . Banks, R . L., Regier, R. B., Ind. Eng. Chem.. Prod. Res. Deveiop., 10, 46 (1971). Banks, R . L., "Topics in Current Chemistry," Band 25. pp 40-67, 1972. Heckelsberg, L. F., Banks, R . L., Bailey, G . C., l n d Eng. Chem.. Prod. Res. Develop.. 7,29 (1968). . , 99 Heckelsberg, L. F., Banks, R. L., Bailey. G. C., J . C ' Q ~ Q ~ 13, (1969). Johnson, M. M . , Nowack, G P., Tabler, D . C., J . Catal.. 27, 397 (1972). Lewis, K . E., Steiner, H., J . Chem. SOC.,3080 (1964).
Receiced f o r ret'iezc July 29, 1974 Accepted December 6, 1974
Catalytic Decomposition of Hydrogen Peroxide by Manganese-Alumina Joseph T. Kohler, Robert E. Altomare, and James R. Kittrell" Department of Chemical Engineering, University of Massachusetts. Amherst. Massachusetts 0 1002
The potential of manganese-alumina for the catalytic decomposition of hydrogen peroxide was investigated in a continuous, tubular, packed bed reactor. First-order kinetics were suggested by the conversion time data. The effect of catalyst particle size on the apparent rate constants was shown to be highly significant. Tests on catalytic activity in the presence of ions such as citrate and phosphate indicated a severe inhibition. The catalytic ability of the manganese was shown to deactivate appreciably over a short time period and then stabilize. Immobilized catalases, in contrast, deactivate severely at identical substrate peroxide concentrations.
The catalytic decomposition of hydrogen peroxide is of interest in many areas of food processing. One suitable use of hydrogen peroxide is the sterilization of food products, for it destroys harmful organisms at low temperatures. Hence, hydrogen peroxide must be removed after its use in pasteurization of milk prior to cheese-making, (Rosell, 1961; Roundy, 1958) and after desugaring of egg whites (Scott, 1956). In addition, hydrogen peroxide occurs as a by-product which must be removed following the enzymatic production of gluconic acid, the enzymatic desugar36
Ind. Eng. Chem., Prod. Res. Dev., Vol. 14, No. 1, 1975
ing of eggs, and the enzymatic oxygen removal from fruit juices. In all these cases, an efficient method of decomposition of hydrogen peroxide into oxygen and water is required. One method of hydrogen peroxide decomposition employs the enzyme catalase, either in the soluble form (Scott, 1956) or immobilized on a solid support (Traher and Kittrell, 1974). However, both beef liver catalase and, to a lesser degree, fungal catalase (Scott and Hammer, 1960) are rapidly inactivated by their substrate, hydrogen
peroxide. Furthermore, this deactivation becomes especially severe at the high peroxide levels desirable for many food applications. It is nevertheless desirable to economically convert hydrogen peroxide into water and oxygen H,Oz
-
H,O
f
i/zO,
(1)
a t the high levels of hydrogen peroxide useful for these food applications. Hence a large number of inorganic metal catalysts were evaluated for their ability to promote the above reaction. Even though such inorganic catalysts would not be expected to be as active as catalase, they would be expected to be robust to high hydrogen peroxide levels. Hence, they would have considerable utility for use in combination with immobilized catalase. For example, the inorganic catalyst could be used a t the top of a tubular flow reactor in the presence of the high peroxide levels of an incoming food product stream; the highly active immobilized catalase could then be used as the bottom portion of the catalyst bed in the same reactor to remove all traces of hydrogen peroxide leaving the inorganic catalyst bed. Thus the range of utility of immobilized catalase could be greatly extended. One superior inorganic catalyst for the decomposition of hydrogen peroxide is a commercially available manganese dioxide supported on activated alumina (Harshaw MN0201). This catalyst was selected from tests in our laboratory of approximately 50 inorganic metal catalysts supported on alumina or silica-alumina. The activity of all of the metals of the group including Ni, Cu, Co, Ni-Co, Fe, and Cr was generally similar and much lower than that of manganese dioxide. Hence, only a representative of this group, Cu, is discussed further in this paper. Ruthenium is also reported to be a superior hydrogen peroxide decomposition catalyst, but no tests were run on this metal. The purpose of the present paper is to summarize the activity and deactivation characteristics of the manganese dioxide catalyst in tubular flow reactors. These results provide incentive for an evaluation of combinations of manganese dioxide and catalase to achieve the decomposition of hydrogen peroxide a t high concentrations of the peroxide. Experimental Section All of the experiments reported in this paper were carried out in a continuous flow tubular reactor similar to that used by Rovito and Kittrell (1973). The feed to the reactor was solutions of known concentrations of Fischer certified H-325 hydrogen peroxide in either distilled deionized water (DDW) or in pH 7 citrate phosphate buffer (prepared by adding about 2 g/l. dibasic sodium phosphate and sufficient sodium citrate to adjust the pH to 7). Normal hydrogen peroxide concentration is 0.01 M , though solutions of peroxide concentration as high as 0.1 M and as low as 0.001 M were sometimes employed. In an experimental run, samples of catalyst as received from the Harshaw Chemical Co. were ground in a hammer mill, screened to yield the appropriate range of particle sizes, washed with DDW, dried, weighed, and charged to, in all cases but one, a glass reactor. Catalysts used were either manganese dioxide (Harshaw MN-0201) or copper oxide (Harshaw Cu-1710); charges ranged from 3.25 to 40 g, depending on activity. The feed tanks were then filled with one of the above mentioned substrate solutions and pressurized to 60 psi. The solution was metered through a Fischer-Porter rotameter at flow rates ranging from 0.1 to 10 ml/sec, passed through a coil in the Haake E-51 controlled constant temperature bath, and finally through a
reactor filled with catalyst. Coolant in the bath was circulated through the jacketed reactor to control the reactor temperature to *0.5"C of the set value. Under the above experimental conditions, the effects of film diffusion on the reactor performance were considered small. Generally, the run-to-run reproducibility was within 5% on conversion (Altomare, 1974). The per cent hydrogen peroxide decomposed (conversion) was determined from measurements on the stream entering and leaving the column. A sample of measured volume taken from the reactor effluent was added to a beaker containing 50 ml of 1% KI solution and a few drops of starch solution. The resulting blue solution was titrated until clear with sodium thiosulfate of known molarity and the concentration of peroxide was determined as
In this study, it was necessary to define contact time in the reactor as the ratio of weight of catalyst charged to the reactor to the volumetric flow rate. A weight basis was chosen because of significant differences in bulk density for different supports and particle sizes; data may be converted to time units by dividing by the bulk density of the bed, which is approximately 1.13 g/cm3. The Results section of this paper is broken into two parts. The first studies the effect of process variables on activity. The second examines the deactivation characteristics of manganese. Results and Discussion Activity Studies: Catalyst Comparison. To compare the activity of several hydrogen peroxide decomposition catalysts, a 0.01 M HzOz solution, buffered with the addition of 2 g/l. of dibasic sodium phosphate and sufficient sodium citrate to bring the p H to 7.0, was passed over 25 g of the supported manganese dioxide. A similar run was conducted with 40 g of supported copper oxide. The particle size range of each catalyst was 500-1000 k . As shown in Figure 1, the manganese dioxide catalyst was 100-fold more active than the copper oxide catalyst, which was in turn representative of the many other catalysts tested, summarized earlier. Note also that the linear dependence of In (1 - x ) on contact time in Figure 1 suggests that the decomposition rate is first order in hydrogen peroxide concentration for the conditions summarized in Figure 1. For reference purposes, 500-1000-p size range particles of catalase immobilized on nickel silica alumina (Altomare, e t al., 1974a,b) would be expected to have an activity about 15-20 times higher than the manganese dioxide. Hence, although the manganese catalyst is not as active as immobilized catalase, its superior resistance to high concentrations of HzOz to be discussed later, combined with its good activity, makes it very suitable for use in conjunction with immobilized catalase. Activity Studies: Particle Size. Figure 2 represents the variation of peroxide decomposition with space time as a function of particle size. Such data were deemed necessary for future design considerations. It is quite evident that catalytic activity, as indicated by the apparent firstorder rate constant h, has a drastic deendence on particle size. In the range studied, 250-1000 g, there is an eightfold difference in k . As expected, the rate of reaction varies inversely with particle size. A more critical look a t the data indicates some slight curvature. This curvature could be the result of catalyst inactivation, which is discussed later in this paper. or possibly the effects of film diffusion a t various flow rates (Rovito and Kittrell, 1973). Normally from the data of Figure 2, one can construct I n d . Eng. Chem., P r o d . Res. Dev..
Vol. 14, No. 1, 1975
37
Feed: 0.01M H202 in pH 7 Citrate-Phosphate Buffer Temperature' 2 5 O C
7 m m Diameter
Reactor
Particle Size: 5 0 0 / 1 0 0 0 ~ 0 Harshow
Cu 1710, 4 0 p m s
k = 0,0008m l I q m - s e c
0 Harrhaw Mn0201, 25qms k : 0.072 mllprn-sec
0.001 0.002 0.003 0.004 INVERSE AVERAGE PARTICLE DIAMETER (MICRONS-')
Figure 3. Variation of apparent first-order rate constant, k , with particle diameter.
\ .011
I
I
I
20
0
I
40
,
I
60
I
I
I
BO
I
,
100
I
I20
I
SPACE TIME (gm-sec/ml)
Figure 1. Activity comparison of manganese dioxide and copper oxide for hydrogen peroxide decomposition.
A 4 2 0 - 5 9 5 ~ ( 5 0 7 & a v g ) k=O36rmlgrn
9%
0 5 9 5 - 8 4 1 l710&oval ~ k:O.I7mllgrn sec
,011 0
'
I
' 6
8
10
12
SPACE TIME (grn-sec/ml)
Figure 2. Activity of MY-0201 as a function of particle size.
an effectiveness factor-Thiele modulus curve. Good correlation has been found in the past (Rovito and Kittrell, 1973), and the effectiveness factor was quite low. However, in this case the attempt was unsatisfactory for a number of reasons. It was felt primarily that the threephase nature of the system did not lend itself to characterization by an adequate kinetic model. Secondly, the large amounts of oxygen gas generated from the decompo38
Ind. E n g . Chern., Prod. Res. D e v . , Vol. 14, No. 1, 1975
sition of the peroxide caused varying amounts of gas bubbles to form, which in turn probably led to an alteration of the mass transfer resistance in a manner not taken into account by theory. Additionally, for high Thiele modulus, i.e. as > 5, there is little difference in effectiveness factor curves regardless of the pellet shape and reaction kinetics (Smith, 1970). For these reasons only the variation of k with particle size is presented in Figure 3. With large Thiele modulus, the apparent rate constant should be proportional to the inverse particle diameter (Smith, 1970). Hence a plot of h us. l/&, should be linear and pass through the origin. The linear dependence is found; however, the curve does not pass through the origin, again indicating breakdown of the Thiele modulus assumptions. Nevertheless certain trends are discernable. Most noticeable is the slight dependence of h with the larger particle sizes. This behavior is quite expected and is generally attributable to intrapellet diffusion resistance limitations in large particles. No studies were conducted below the 250-297-p range because of pressure drop considerations. Activity Studies: Buffer Effect. In the course of early screening experiments with MY-0201 catalyst, it was observed that the activity of the catalyst depended on the nature of the feed solution. The manganese catalyst appeared significantly less active in the presence of citratephosphate buffer as compared to DDW. To confirm this observation, space time-conversion data were obtained for MN 0201 in a flow reactor with and without buffer present in the feed. Two 7-mm reactors were charged with 13 g of 841/1000 p MN-0201. The feed solution passed over the first reactor consisted of DDW with peroxide; that for the second reactor was pH 7 citrate-phosphate buffer and peroxide. Both solutions were 0.01 M in peroxide. The apparent first-order rate constant in both cases is shown in Figure 4. The rate constant is reduced by a factor of about 3 when the feed contains citrate-phosphate ions. An interesting phenomenon was observed when peroxide and DDW was passed over the previously used second reactor. The rate constant was very nearly equal to the rate constant observed in the first bed. Thus it appears that there is a reversible, competitive inhibition of the manganese catalyst in the presence of citrate-phosphate buffer. The effect of buffer concentration on inhibition was not examined. An additional study was performed under similar conditions to test the effects of temperature and bed dilution. Two identical 20-g charges of MN-0201 were used. The first was placed in a 10 mm diameter glass reactor. The latter was diluted with 180 g of Harshaw catalyst 0704,
g
V 0
)
:
.
-
s 0 6 0
5 -
I3pms Harshaw Mn-0201,841/1000LL
0.4
p H 7 DDW,0025M H202,10mm Reaclor
21
-
.
e
LL
2
0
6
4
10
8
I2
16
14
20
18
S p o c e Time (gm-sec/mC
Figure 3 . Space time-conversion curve for H202 decomposed by MS-0201 at 2 5 T . 100
.
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0 p H 7 , DDW, 25'C,O.01MH2O2 0 pH 7, Citrate-Phosphate, 25'C,0.01 M H P 2 A pH 7, Cilrate-Phosphate, 1O.C ,0.1 M H202
0 p H 7 , Citrate-Phosphate, IO'C,O.I M H 2 0 2 , Diluted with l8Opms Ni-Kieselquhr 841/1000# Catalyst E ! L
'
Harshaw M n O Z O l 84111000fl
.o2-
0
4
8
I2
16
20
24
ELAPSED TIME
28
32
36
40
44
48
(HOURS)
Figure 6 . Deactivation studies for Harshaw Mh--0201; feed rate 1 ml/sec, temperature 25°C.
Figure 4 . Examination of the effects of buffer and temperature on Harshaw MN-0201.
Ni-Kieselguhr. The purpose of diluting the bed was to simulate the performance of MN-0201 for eventual use with immobilized glucose oxidase on Kieselguhr. These data are also shown in Figure 4. The peroxide concentration was 0.10 M As expected, the data follow apparent first-order kinetics. Results for both beds were identical, indicating that dilution has little or no consequence on MX-0201 performance. An indication of the activation energy for peroxide decomposition by the manganese catalyst can be otained from the apparent rate constants observed in pH 7 buffer at 25 and 10°C. Using the Arrhenius equation, an approximate activation energy of 8650 cal/mol was calculated from the two available data points. This value is quite reasonable when compared to the activation energies of 18.0 kcal/mol for spontaneous peroxide decomposition, 13.5 for I- ions, 6.4 kcal/mol for soluble catalase (Whitaker, 1972) and 4.7 for immobilized fungal catalase (Altomare, et al , 1974a). However, in regions of high intrapellet diffusion resistance the reaction rate is proportional to the square root of h for first-order kinetics (Smith, 1970). If this is the case, as we believe, the calculated activation energy is only one-half the true activation energy. Hence the true activation energy could be as high as 17.3 kcal/ mol for manganese-catalyzed peroxide decomposition. Since many possible food applications of the manganese catalyst may have glucose in the feed stream, 0.4% glucose was added to a feed solution along with 0.023 M hydrogen peroxide. The result is compared in Figure 5 to a case with no glucose and lower HzOz. Evidently both glucose level and peroxide level have no effect on the conversion obtained, as would be expected for first-order kinetics. Deactivation Studies. In Figure 6 are summarized several observations regarding the rate of deactivation of manganese during hydrogen peroxide decomposition. The most notable characteristlc of this catalyst is shown by the open points of Figure 6, demonstrating the effect of
peroxide level on performance. First, there is little or no effect on initial activity, as expected from Figure 5 and the first-order kinetics. The high level of peroxide, 0.01 M,does not increase the rate of deactivation quite in contrast to immobilized beef liver catalase (Altomare, et al., 1974b) or fungal catalase (Altomare, et al., 1974a). Indeed, high peroxide levels may enhance the stability of the manganese. For comparison, at the 0.01 M H202 level, immobilized beef liver catalase deactivates a t a rate of 12% per hour (Altomare, et al., 1974b) immobilized fungal catalase deactivates a t a rate of 1.3% per hour (Altomare, et al., 1974a), and the manganese of Figure 6 deactivates a t a rate of 0.15% per hour. Since manganese is very stable at high peroxide concentrations and immobilized catalase is highly active, an optimal combination of a manganese bed followed by an immobilized catalase bed is obvious for such applications as the peroxide pasteurization of milk. Also shown in Figure 6 are the effects of buffer ion concentration (open circles us solid circles) and catalyst particle size (solid circle us solid squares). The effect of these variables in initial activity is similar to that reported earlier. The presence of the buffer ions causes a severe loss of activity; only 3.25 g of 250-297-~particles provides about the same activity as 13 g of 841-1000-~particles. Although the test of Figure 6 is a not very sensitive measure of the effect of these variables on deactivation rate, it is apparent that none of the variables studied in Figure 6 greatly affects deactivation rate. Finally, the steep initial decline in activity for the first 6 hr of the data of Figure 6 appears to be only an initial transient. For any applications of manganese to peroxide decomposition, the longevity of the active catalyst appears such that this transient can be ignored. Conclusions On the basis of the foregoing analyses we conclude that manganese-alumina is a satisfactory catalyst for use in hydrogen peroxide decomposition either alone or in conjunction with immobilized catalase. The catalyst, though Ind. Eng. Chem., Prod. Res. Dev., Vol. 14, No. 1 , 1975
39
not as active as catalase, demonstrates superior resistance to the effects of peroxide at high concentrations and deactivates slowly enough to be considered for long term usage. The kinetics of the system appear to be first order over the range of variables studied, but corrections for the rate constant are required due to particle size or feed buffer ions. An activation energy of about 8600 cal/mol is probably satisfactory for particle sizes of the present study, but should be used with 'caution for other particle sizes.
Altomare, R. E., Greenfield, P. F , Kittrell, J. R.. Biotechno/, Bioeng., accepted for publication (1974a). Altomare, R . E . , Kohler, J.. Greenfield, P. F., Kittrell, J. R . , Biotechno/. Bioeng., accepted for publication (1974b) Rosell, J. M.. Can. Dairy I c e Cream J . , 40, 50 (1961) Roundy, 2. D . , J . Dairy Sci.. 41, 1460 (1958). Rovito. 6.J . , Kittrell, J. R . , Biotechnol. Bioeng., 15, 143 (1973) Scott, D . , U.S. Patent 2,758,934 (Aug 14, 1956) Scott, D , Hammer. F . , fnzym., 22, 229 (1960) Smith, J. M.. ' Chemical Engineering Kinetics," 2nd ed. McGraw-Hill, New York, N.Y., 1970 Traher. A D . , Kittrell, J. R.. Biotechno/. Bioeng.. 16, 413 (1974). Whitaker. J. R , ' Principles of Enzymoiogy for the Food Sciences," Marcel Dekker, New York. N. y . , 1972
Literature Cited Receiced for reuzew J u n e 12, 1974 Accepted D e c e m b e r 9, 1974
Altomare. R . E., M.S. Thesis, Department of Chemical Engineering, University of Massachusetts, Amherst, Mass.. 1974.
Levulinic Acid from Sucrose Using Acidic Ion-Exchange Resins Richard A. Schraufnagel and Howard F. Rase" Department of Chemical Engineering, The University of Texas, Austin, Texas 78772
The production of levulinic acid and 5-hydroxymethylfurfural was studied using highly acidic ion-exchange resins as catalysts and sucrose as a convenient soluble polysaccharide. Four commercial resins were compared at 100°C. Resin pore size had a major effect on selectivity. Ion-exchange resins are attractive for this process because of the selectivity possible by pore-size control and the reusable character of any solid catalyst. Higher temperature resins, however, are needed to take advantage of faster rates possible with modest temperature increases.
Processes developed in the 1920's to 1940's for producing useful chemicals from farm products proved in most cases to be uneconomical because low cost natural gas and oil assured the success of a rapidly growing petrochemical industry. Recent events, however, have dramatized the need for replenishable raw materials such as farm products or, more particularly, waste materials from such products as alternate sources of industrial chemicals and polymers. It would seem desirable therefore, to reexamine older studies in this area with the view of adapting newer techniques and materials to producing useful products. One such product, levulinic acid (4-oxopentanoic acid, CH3COCH2CH2COOH), has been known since 1840 to result from the reaction of carbohydrates with mineral acid (Mulder. 1840). It has great potential as a chemical intermediate in producing various pharmaceuticals, pesticides, and plasticizers. A solution of its sodium salt is a less corrosive permanent anti-freeze than ethylene glycol (Aries and Copulsky, 1949), and levulinic acid itself is used as a dye mordant in the textile industry. Most of its extensive possible uses, which have been documented by Leonard (1956) and Morton (1947) have not been developed because of low yields of existing processes and, until recently, lower costs of petroleum derived intermediates. The many new, highly efficient, and rugged ion-exchange resins now marketed offer possibilities for more effective and economical means for converting carbohydrates to levulinic acid. In order to assess such possibilities sucrose was used as a convenient soluble polysaccharide for evaluating four commercial ion-exchange resins described in Table I as catalysts in producing levulinic acid and an intermediate product, 5-hydroxymethylfur40
Ind. Eng. Chem., Prod. Res. Dev., Vol. 1 4 , No. 1. 1975
fural (5-HMF), which has potential use as a raw material for adipic acid manufacture (Aries and Copulsky, 1949). Because of its major importance in foods and concomitant price volatility, sucrose would not be a logical raw material source, but it serves as a useful model compound since it hydrolyses into the well-known monosaccharides, glucose and fructose. Previous Work Early work on levulinic acid involved studying different carbohydrate sources including glucose, fructose, lactose, sucrose, starch, cellulose from wood pulp, and chitin (Wiggins, 1948). Yields in these studies, performed at atmospheric pressure, were usually less than 25%. In later studies by McKibbins (1962) autoclaves were used to increase reaction temperature to 160-200°C and thereby increase yields. The effects of acid concentration and type were also studied. Commercial production using an autoclave began in the United States in 1940 by A . E. Staley Co. using dextrose as feed and HC1 as the acid (Meyer, 1945). Four additional patents exist (Redmon, 1956; Dunlop, 1957; Carlson, 1962; Sassenrath, 1966). The process by Redmon employs an ion-exchange catalyst. The advantages of an ion-exchange process are : (1) little humin (or solid waste) byproduct, (2) reaction temperature can be low, and (3) the catalyst can be readily separated from the reacted mixture, and regenerated. The patent by Dunlop (1957) describes an atmospheric pressure process for producing levulinic acid from any hexose-yielding substance ranging from sucrose to cellulosic wastes such as corn cobs, bagasse, grain hulls, and wood products.
