REACTION RATE STUDIES OF CATALYTIC OXIDATION OF ETHYLENE LARRY G. NAULT,’ DONALD W.
BOLME,B AND LENNART N. JOHANSON
Cniversity of Washington, Seattle, L’ash.
The kinetics of the reactions of oxygen and ethylene to form ethylene oxide and carbon dioxide were studied using a silver catalyst on an impervious granular support. Catalyst behavior was sensitive to pretreatment conditions, but good stability and activity were attained with the specific pretreatment used for the kinetic runs, conducted a t 200” C. Ethylene, oxygen, and carbon dioxide feed partial pressures were varied independently using an inert make-up gas (nitrogen) to hold flow rates constant. Differential and integral packed bed reactor data could b e fitted by the same postulated rate equations. No correspondence was found between reported chemisorption behavior on silver surfaces and apparent adsorption behavior of reactants and products on the catalyst, as deduced from kinetic studies. The reaction product carbon dioxide strongly inhibited the reaction rates.
THE general field of heterogeneous kinetics and catalysis is still hampered by a lack of understanding of the exact nature of a catalytic surface. I n some cases the catalyst acts a t specific sites and the microscopic identity of the catalytic site is probably known-for instance, when a n ion exchange resin is used as a n acid catalyst. The “clean” metallic surface has also been extensively studied. T h e reaction takes place with no apparent preference as to microscopic location on the surface. I n both types of catalyst the adsorption on the catalytic site in some cases has been closely associated with directly measured adsorption of the individual compounds. Rate equations based upon the Langmuir isotherm model as developed by Hougen and Watson ( 9 ) and Schwab, Taylor, and Spence (77) have been used with some success to interpolate and extrapolate the effects of temperature and concentration. The general class of promoted (or promotable) catalysts cannot be treated in so direct a manner. T h e fact that promoters are present and are, presumably, involved in the microscopic catalytic areas means that these areas or sites may not behave like the gross surface. Reactants and products may therefore adsorb differently on the catalytic site than on the gross surface. Direct measurement of adsorption is therefore not particularly useful. The major catalytic activity may be associated with sites of a narrow range of activation energies as proposed by Constable (7), in which case the Langmuir model may still apply. Few, if any, such systems have been proved to exist, hoicever, and the use of such a model for a promoted catalyst system is open to question. The work reported here, the reaction between ethylene and oxygen on a silver catalyst, involves a promotable catalyst. It was undertaken in part to explore the relationship between adsorption for the gross catalytic surface as measured directly and adsorption on the catalytic site as inferred by the effect of concentration upon catalytic rate. The work is divided into two parts. The first, more extensive study utilized a differential reactor technique, which allows the rate to be determined directly, in contrast to the integral
Present address, Southern Dyestuffs Co., Charlotte, N. C. Present address, Washington State University, Pullman, Wash.
reactor method, in which the rate is usually measured by obtaining the slope of a plotted line. The effect of reaction products upon rate may be determined with a differential reactor by adding individual products to the feed stream. This technique also minimizes the complications normally arising from the accumulation of reaction products and by-products, and from the heat of reaction. The second part involved integral rate data, which were obtained to verify the form of the equations developed with the differential data and to obtain a direct comparison between differential and integral rate data collected under similar conditions. Such comparisons are seldom made and are valuable in evaluating the importance of temperature and concentration gradients in integral reactor beds.
Previous Work
Kinetics. Much of the reported work on the kinetics of this system has been done with industrial application as a goal. Reaction mechanism, rate-controlling step, adsorption effects, etc., were either not determined or were “guessed at,” using data admittedly ill suited for this purpose. Investigations undertaken to establish rate-controlling steps in the diffusion-reaction sequence include those of Twigg (20-23), Kummer (70), Orzechowski and MacCormack (75), Roginskii and Margolis (76), Kurilenko, Kul’kova, Rybakova, and Temkin (729, and Gorokhovatskii, Rubanik, and Kholyavenko (8). Orzechowski and MacCormack concluded that the kinetics of the reaction are controlled by the rate of the reaction a t the catalyst surface substantially independent of transport, adsorption, and desorption rates. Subsequent authors have not disagreed with this view. Twigg points out that his kinetic data indicate that ethylene is moderately strongly adsorbed on the catalyst and oxygen is only slightly adsorbed (20). H e rejects this logical basis for postulating rate equations, as it appeared to him inconsistent with observed adsorption behavior of ethylene and oxygen on silver surfaces. Oxygen is known to be strongly adsorbed on silver (as the oxide), and ethylene is known not to adsorb on silver. I t is apparent that the effective catalyst site is not fully oxidized silver, since the presence of ethylene in the reactant stream results in a much lower state of oxidation than when VOL.
1
NO.
4
OCTOBER 1 9 6 2
285
Table 1. Comjound
Hydrogen Oxygen Carbon dioxide Ethylene Ethylene oxide Acetaldehyde Water a
Summary of Adsorption Information Adsorption Adsorption on Silver Adsorption on on Silvera Oxidea Catalytic Sit&
None Chemical None None Chemical Chemical None
Based on available literature.
Reacts Little or none Strong Strong Forms part of site?
Chemical Reacts Reacts Reacts Strong b
Based on kinetics of this work.
oxygen is present alone. The possibility that the catalyst site was something other than pure silver or silver oxide was not explored by Twigg. Kummer (70) used single silver crystals and achieved incremental reaction by recycling his reactant stream, removing the products in cold traps. and making up reactants as needed. His kinetic results. in general, supported those of Twigg. Orzechowski and MacCormack (75) found that catalyst activity varied with recent catalyst history but that this variation could be either avoided or minimized. Similar difficulties are reported by Kurilenko et a / . (72). Several authors have reported inhibition of the reaction by the reaction products. Roginskii and Margolis (76),Kurilenko et al. (72), and Gorokhovatskii et ai. (8) found that ethylene oxide inhibited the reaction. Kurilenko and Gorokhovatskii, et a / . , in publications available after this work was completed, also state that carbon dioxide and water inhibit the reaction. Gorokhovatskii, Rubanik, and Kholyavenko give the order of inhibition as C 2 H 4 0 > COt > H 2 0 . Since the reactions are essentially irreversible, such inhibition cannot be due to the
k --
To Sfandpipe
Ressure Gauge Gas Inlet Thermocouple
h"S/oinless
9 4 . Stainless
Steal Pipe S t e e l Shell
Dou f h erm Mica Lead to 3 4 R Resisfance L e a d fa 60R Resistance Cafalysf
Bed
Loose Diafamaceous Earth Gas Ouile? Thermocouple B a t t o m Thermocouple (safety suilch bahindl
Tronsife Disks
Tripod
Figure 1. 286
(5). Ethylene is adsorbed very weakly, if at all, on silver (27). Ethylene oxide and acetaldehyde (the latter a possible shortlived intermediate in the oxidation of ethylene oxide) are adsorbed with decomposition to form a nonvolatile layer, composed of carbon, hydrogen, and possibly oxygen, combined in indefinite and varying proportions (27). "Interdependence" is used in this work to denote the possible adsorption of one species (example, oxygen), forming a new surface (silver oxide) upon which other species may be adsorbed. There is definite indication that interdependence in adsorption occurs on silver. Carbon dioxide adsorbs on silver oxide, forming a (surface) carbonate ( 3 ) ; water adsorbs on silver oxide (5); ethylene, ethylene oxide, and hydrogen all react with oxygenated surfaces (27),and the possibility of their adsorption for catalytic purposes cannot be discounted. The resulls of previous work on directly measured adsorption are summarized in the first two columns of Table I. Equilibrium. The free energy changes of the oxidation of ethylene to ethylene oxide, of ethylene to carbon dioxide and water, and of ethylene oxide to carbon dioxide and water have been calculated by standard methods from free energy of formation of the compounds at 25' C. (6). All give large negative values indicating essential irreversibility of the reactions, in agreement ivith TIvigg's results at a higher temperature (20).
Steel Pipe
Temperature Confrol Swifch (thermocwple w.dl behind)
Mild
reverse reaction but must result from reversible deactivation or coverage of the catalyst. Adsorption. Oxygen is chemisorbed on silver, forming silver oxide. Physical adsorption of oxygen in a molecular as well as an atomic form has also been suggested (7, 2, 4, 7 7 , 73). Lewis (73) has demonstrated that reduced silver catalyzes the decomposition of silver oxide, and Benton and Elgin (4)concur. Hydrogen undergoes a weak (physical) adsorption on silver ( 3 ) . It apparently does not dissolve in silver below about 40OoC. (78). Carbon dioxide is not adsorbed on silver above 0' C. and exhibits only a weak physical adsorption below this temperature (3> 27). LVater is adsorbed on silver, obeying a Langmuir isotherm
Leg
Reactor design
I&EC PROCESS D E S I G N A N D DEVELOPMENT
Experimental The reactor (Figure 1) consisted of a 3/4-inch stainless steel pipe 21 inches long, surrounded over its entire length by a Dowtherm bath. The catalyst bed was located in the bottom portion of the tube with about 11/2 inches of 4-mm. Kimble brand glass beads between it and the bed support. When used as a differential reactor, the bed occupied the next 9 inches and the top portion of the tube was filled with 4-mm. beads to serve as a preheater and temperature stabilizer for the entering gas. LVhen used as an integral reactor, the top portion of the tube was filled with blended layers of catalyst and glass beads. Iron-constantan thermocouples were located a t the bed exit, the top and bottom of the Dowtherm bath, and the top of the reactor pipe. Analysis was by means of gas chromatography. The column used for the separation of ethylene oxide consisted of 1.5 feet of 6-mm. glass tubing containing 80- to 100-mesh Johns-Manville (2-22 firebrick with 4 grams of cetane (Du Pont n-hexadecane) per 10 grams of firebrick. The column used for the separation of ethylene and carbon dioxide consisted of 3 feet of 6-mm. tubing containing 80- to 100-mesh silica gel (which separated carbon dioxide and ethylene from air) followed by 6 feet of silver nitrate-saturated glycol on 80- to 100-mesh firebrick. The silver nitrate separated the ethylene from the carbon dioxide. Gas flow rates were closely controlled by using standard pressure reducers, needle control valves, and capillary tubemanometer flow indicators. The pressure on the upstream
side of the needle valve was kept high (about 4 atm. absolute) to take advantage of the damping action of limiting (Mach 1) velocity a t the throat of a nozzle. T h e needle valves, except that for nitrogen, had 20 turns from full open to full closed. T h e manometer fluid was di-n-butyl phthalate. T h e catalyst was designed with the goal of eliminating, as far as possible, any mass or temperature gradients both in the bed as a whole and within individual catalyst particles, and making it similar to that used in commercial processes. T h e catalyst was made by silver plating 6- to 8-mesh silicon carbide granules six times and slurrying the plated granules with silver oxide using the method of Wan ( 2 4 ) , but adding no promoter. Prior to use for a run the catalyst was loaded into the reactor and the reactor heated as rapidly as possible to 250" C. with oxygen flowing over the catalyst at about 1 to 2 ml. per second. T h e oxygen was turned off and the reactor flushed with nitrogen for 5 minutes a t 50 ml. per second. The catalyst was then treated with 2% ethylene in nitrogen a t a combined flow rate of 50 ml. per second until the catalyst was reduced (about 10 minutes). Following this the catalyst was treated with 10% ethylene for 5 minutes a t 50 ml. per second and 100% ethylene a t 5 ml. per second for 30 minutes. T h e reactor was then cooled to operating conditions under ethylene at a reduced flow rate.
-
4 1 ~
jit. P
*----a
Run 8 1200°C) Formation of Ethylene Oxide Formotion of Carbon Dioxide
2 3
-0
0.01
Figure
0.02
0.03 0.04 0.65 0.06 Mole fraction Cop hole! Boutlet1
0.07
0.68
2. Effect of carbon dioxide on reaction rate
*---e
Run 5 1200°Cl Formation of Ethylene Oxide Formation of Carbon Dioxide
It was hoped, by this procedure, to produce an unpromoted catalyst. Performance indicated that the catalyst was promoted, however, probably by potassium present in the silver plating step. P
Differential Rate
B
Studies
3 005
Orzechowski and MacCormack (75) found that the catalyst activity varied with the pretreatment and was, at least partially, a function of the concentration of the reactants and products as well as of temperature. The rate of change of catalytic activity was slow, however. I t was therefore decided. in studying the effect of reactant and product concentration upon rate, to vary only one component at a time, so that the first and sixth, the second and seventh, etc., points were duplicates. A typical run (run 5) would then be as shown in Table 11. The catalyst in this run is out of contact with the original reactant mixture only 6 hours, 2:05 to 8:05. I n most cases the catalyst did not change activity during a run, and when it did (run 6) the discrepancy was detected and explained by later work which investigated the effect of changes in pretreatment. I n all runs the standard conditions were 20 mole % oxygen, 10 mole yo ethylene, and 70 mole % nitrogen at a total flow rate of 50 ml. per second a t room temperature. The first and last points in every run were taken at these standard conditions, as were all points in the pretreatment runs. I n each run the partial pressure of a given reactant or product component was
Table II.
0 13 0 15 Mole fraction Oxygen (outlet 8 inlet1
Figure 3.
0 20
Effect of oxygen on reaction rate
varied while partial pressure of all other reactants or products, as well as space velocity, was held essentially constant, by appropriate adjustments in flow rate of the inert carrier gas, nitrogen. All runs depicted were made with the same catalyst charge. Figures 2 through 4 show typical effects of varying the concentration of the inlet carbon dioxide, oxygen, and ethylene upon reaction rate. Each data point is represented by two triangles or dots connected by a straight line. The two points, in each case, represent the inlet and outlet concentrations of the variable in question. Separate tests in which two operating sequences were used, one for increasing and the other for decreasing initial concentrations of carbon dioxide, were in good agreement, indicating reproducibility of data and stability of catalyst during the period of operation (7 hours). Changes in reaction rate as a result of changes in feed gas composition are apparently rapid and reversible.
Summary of Data from Typical Run
(Run 5) C
D
E
0:oo
B 2:05
3:05
4:05
5:05
F 6:05
G 7:05
8:05
20.2 69.7
10.06 79.9
5.02 84.9
3.19 86.9
3.20 86.8
4.91 85.0
10.02 80.0
20.15 69.65
10.07 99.97
10.06 100.02
10.01 99.93
9.96 100.05
10.0 100.0
10.07 99.98
10.04 100.06
_10.12 _
0.548 0.259 49.48 199.0 211.2 2:oo
0.293 0.145 49.7 199.6 205.4
0.225 0.085 49.79 200.2 204.0 4:OO
0.183 0.047 50.2 199.4 203.4 5:OO
0.219 0.047 50.0 202.4 205.6
0.227 0.076 49.88 200.4 204.4 7:OO
0.313 0.144 49.91 199.0 204.4 8:OO
0.485 0.282 49.50 199.4 205.8 9 :00
Sample
A
Time niix introduced Inlet comDosition. mole %
FI
I.
0 2
N* CO, &HI
Total Products in outlet, mole yo
coz
CzHO Flow rate, ml./se: Bath top temp., C. Bed bottom temp., C. Time sample taken
3:OO
6:OO
VOL.
1
NO. 4
99.92
OCTOBER 1 9 6 2
287
P 2-
.
B E
*--e ----e-*--
*-e
-
k
9
+----e
9
4
:o
002
001
Figure 4.
Run 6 (200°C.) Formation of Ethylene Oxide Formotion of Corbon Dioxide
004 005 006 007 Mole froction Ethylene (outlet-inlet)
003
008
009
OiO
Effect of ethylene on reaction rate
Rate-Controlling Step. The following brief discussion leading to a plausible rate-limiting reaction step is in part qualitative, in part quantitative. Two features of the differential rate data obtained were particularly helpful for rapid assessment of possible rate equations: The differential form of the data avoided the need to work with the more complex integrated equations, and independent variation of partial pressures simplified analysis. The rate of a catalytic reaction is controlled by the rate of the sloivest step or sequence of steps in the reaction. The rate of reaction may be controlled by 1. Any of a number of transport processes. 2. The rate of adsorption of one or more of the reacting species upon the catalyst. 3. The reaction rate of the slowest step or steps in the surface reaction. 4. The rate of desorption of product or products. 5. Some combination of the above. By referring to Figure 2, it may be seen that the addition of less than 2% carbon dioxide to the reactant mixture halved the rate of ethylene oxide formation (and markedly reduced the rate of carbon dioxide formation), even though the reactions involved are substantially irreversible. There is in the flow stream a large quantity of gas not taking part in the reaction and the addition of such a small amount of carbon dioxide would not be expected to effect gross changes in transport bet\veen the main stream and the catalytic site. Also, the catalyst, impervious silicon carbide coated with silver, was made with the elimination of pore diffusion difficulties as a
specific objective. I n view of the above: step 1 cannot be considered a major rate-controlling step. If an irreversible reaction, such as either of those being studied, is controlled by the rate of adsorption of one of the reactants, increasing the concentration of the other reactant either would not affect the rate, if it were not adsorbed on competitive sites, or would lower the rate if it Lvere adsorbed on competitive sites. In no case would raising the concentrations of the noncontrolling reactant increase the rate. Figure 3 indicates that increasing oxygen concentration increases the rate of both reactions. This could not happen if the adsorption of ethylene were rate-controlling. I t follows that the adsorption of ethylene is not rate-controlling in either reaction at the ethylene concentration used (0.1 0 atm.). Similarly> Figure 4 indicates that the adsorption of oxygen is not ratecontrolling in either reaction a t the oxygen concentration used (0.20 atm.). In the Langrnuir adsorption-reaction theory, the assumption is made that the rate of adsorption of a reactant is directly proportional to the concentration of that reactant a t the interface and the concentration of the unoccupied adsorption sites (more exactly, the concentration of adsorption sites not denied to it). 4 reaction in which adsorption of both reactants \\-as rate-controlling would necessitate a catalytic surface almost completely free of reactants and products. In this case the rate would be a linear function of all reactant concentrations. Run 6. Figure 4. indicates that the rates are not linear functions of the ethylene concentration and therefore the rate of adsorption of both reactants is not rate-controlling. The reaction rate has been shown not to be dependent upon the rate of adsorption of ethylene, oxygen, or both under the conditions used. Therefore, step 2 is not a logical rate-controlling step. Other adsorption models would result in similar conclusions. The possibility that the rate of desorption of a product of the reaction is rate-controlling can be discounted as follows. In an irreversible catalytic reaction where desorption of a strongly adsorbed product is rate-controlling, the catalytic sites are all quickly occupied by that adsorbed product, whereupon the rate is determined by the number of catalytic sites and the rate of desorption (.9), The rate in this case is not affected by reactant concentrations. Figures 3 and 4 indicate that reactant concentrations do affect the rate of both reactions. Therefore, step 4 is not considered a possible rate-determining step for both reactions.
O-----O
-
Run 12 (200°Cl Formotion of Ethylene Oxide Formotion of Corbon Dioxide Overnight oxidation 15 min. dilute H2 105 min. pure Hz 60 min. pure C 2 H q N p added then 02
Run 1 1 (200'(;1 Formotion of Ethylene Oxide of Corbon Dioxide Overnight oridotion(Pur8 0,ot N p ond then C2H4 introduced
c----O Formation
2OPC )
A
\
i0
Figure 5. catalyst 288
60
90
120
Time, minutes
160
Effect of pretreatment with ethylene on reduced
I&EC PROCESS DESIGN A N D DEVELOPMENT
^0
30
Figure 6.
I
60
90 o
a
Effect of pretreatment with oxygen
o
The above considerations leave step 3 as the most plausible rate-controlling step under the conditions of 0.10 mole fraction ethylene, 0.20 mole fraction oxygen, 0.0 mole fraction initial product concentrations. This is in agreement with Orzechowski and MacCormack. Adsorption on Catalytic Sites. Figure 2 shows the rate of ethylene oxidation to ethylene oxide and to carbon dioxide as functions of carbon dioxide concentration. Space velocity and reactant concentrations were held constant. Since all reactions in the system are thermodynamically irreversible, the carbon dioxide must be inhibiting the reactions by inhibiting the rate-controlling step. If the surface reaction is ratecontrolling, this must be accomplished by deactivating sites. This deactivation was found to be reversible and therefore due not to poisoning but to a relatively rapid, reversible occupancy of the site. This indicates that carbon dioxide is strongly adsorbed on active sites for both reactions. Figure 3 shows the rate of ethylene oxidation to both ethylene oxide and carbon dioxide as functions of oxygen concentration. During the run the space velocity and the ethylene concentration were held constant. Nitrogen was varied to compensate for the variation in oxygen flow. The rate c j . concentration curves for both reactions are essentially linear, which means that oxygen is weakly adsorbed on the active catalyst site if it is adsorbed a t all. Small deviarions from a straight line could be explained by the effects of slightly increased carbon dioxide concentration and decreased reactant concentrations in the latter part of the bed, the effect of which Ivould increase with increasing rates: causing the line to cusve downward. Figure 4 shoFvs the rate of ethylene oxidation to ethylene oxide and to carbon dioxide, both as a function of ethylene concentration. As in Figure 3, only the one reactant concentration xias varied, and the space velocity and the other reactant concentration Ivere held constant by varying nitrogen flolv. The ethylene oxide formation sho1z.s a slow change in catalyst activity with ethylene concentration change. iVhether one takes the descending or ascending data, hoivever, the ability of a n increase in ethylene concentration to increase the reaction rates depends strongly on the amount of ethylene already present. This behavior is characteristic of strongly adsorbed reactants. As the ethylene concentration increases? a n increasing percentage of the catalytic sites is
-
*----a
t
Run 13 C2CPC.l Formation o f Ethylene Oxide Formation of Carbon Dioxide Overnight oxldotion Flush with Ne 110 min pure H p Mix added without flushing
taken up by adsorbed ethylene and the catalyst approaches saturation. This curvature of the curves of rate us. ethylene concentration cannot be assigned to increased carbon dioxide concentration, decreased oxygen concentration, or temperature rise in the latter part of the bed, because all of the above variations were present to the same extent in experiments reported in Figure 3, in which practically no curvature resulted. I t is concluded that a t 200' C. carbon dioxide and ethylene are adsorbed on the catalytic site or sites, and oxygen is not adsorbed or is adsorbed very slightly. These conclusions in regard to adsorption of ethylene and oxygen were suggested by Twigg based on his kinetic data. He, however, rejected this explanation of kinetic behavior and instead adopted a mechanism consistent with the adsorption characteristics of metallic silver surfaces. Effect of Pretreatment. Some observed effects of pretreatment of the catalyst are shown in Figures 5 through 8. One test was of a pretreatment with ethylene of a catalyst soaked overnight in pure oxygen at 200' C. After an initial surge of products, possibly owing to the extra ethylene adsorbed on the catalyst, the activity fell s1owly, almost leveling off after 120 minutes. Plotted curves of reaction rate us. time Lvere similar in shape to those of Figure 5, which shows the effect of pretreatment with ethylene on a catalyst reduced with hydrogen. After the initial surge of products, the catalyst assumed a higher activity than with the oxygen "presoak," and the activity fell more rapidly. Both the higher activity and the fact that the activity fell faster Ivith hydrogen reduction seem to indicate that the extra reduction in Figure 5 produced a greater excess of sites when compared to steady-state conditions than ethylene reduction only, and that the catalyst was less than completely oxidized. The reduced state of a similar catalyst has been demonstrated by Twigg (23). Figure 6 depicts the effect of overnight oxidation (pure oxygen a t 200' C.) upon the catalvst. The initial reduction and the gradual rise in catalytic activity both support the theory that the catalytic activity is at least partially diminished by the oxidation. Figure 7 depicts the effect, upon the catalyst, of a simple reduction with hydrogen, and indicates that the catalytic site is not produced by a simple reduction. Treatments \vith ethvlene resulted in higher catalyst activity, and therefore
-
)---*
z
.
$ e 6
6
RunP14 (2oO'Cl Formation o f Ethylene Gxlde Forrnatlon of Corbon Dloxlde Overnight oxldatlon Flush with N2 110 min pure HZ 6 0 min pure C 2 H 4 0 M I X added without flushlng
E
P
*
n
9
9
P
E
6 P
t P
$
ro'
9
s
P
0
30
60
90
T i m e , minuter
Figure 8. Effect of pretreatment with ethylene oxide on reduced catalyst VOL. 1
NO. 4
OCTOBER 1 9 6 2
289
ethylene or a similar compound appears necessary for production of a catalytic site. Figure 8 depicts the effect of pretreatment with ethylene oxide upon the reduced catalyst. Activity is unexpectedly high. Ethylene oxide is adsorbed on silver to form a thick carbonaceous layer (27). During the first 15 minutes, however, 1.1 grams of carbon dioxide and 0.94 gram of ethylene oxide were produced for every gram of catalyst present. This is equivalent to about 1.5 grams of ethylene oxide released per gram of catalyst, considerably more than could be stored in a packed catalyst bed such as that used. I n view of this and the fact that the ethylene in the off-gas was greatly depleted, the high rate of appearance of ethylene oxide and carbon dioxide in the product stream must be explained in some other way, such as hyperactivation of the catalyst. The run was terminated after 30 minutes, as it became obvious that the heat of reaction was not being removed by the Dowtherm bath and the bed was greatly overheated. A later run a t a lower temperature of 180' C. indicated that the catalyst activity deteriorates approximately as a first-order reaction. The pretreatment studies seem to indicate that: The catalytic site is a t least partially reduced. T h e catalytic site involves an ethylene oxide molecule or molecules or molecular fragments. The treatment of the hydrogen-reduced catalyst with ethylene oxide produces either very many sites or very active ones. I t might also be concluded that some sites may be rapidly destroyed by removing the ethylene which is necessary for their constitution and/or maintenance, but that once so destroyed the sites are slow to rebuild (Figure 6). This would suggest that in Figure 4 the ascending curves are obtained with a constant number of catalyst sites and are therefore the ones to use in any mechanistic determination. Directly a n d Indirectly Measured Adsorption. Adsorption on catalytic sites as determined by catalytic experiments and direct adsorption data as normally measured are compared in Table I . This table indicates that there is no correlation between adsorption as measured directly on silver and adsorption as it affects the kinetics of the catalytic oxidation, but that there may be some qualitative correlation between adsorption on silver oxide and the catalytic reaction. Twigg and others have demonstrated, however, that the catalyst does not consist solely of silver oxide. Therefore, it cannot be assumed that such correlation will hold for all reactants, products, or contaminants-for example: water. In retrospect this does not seem surprising; the silver-silver oxide catalyst is one for which successful promoters have been claimed and the effect of promoters is to produce abnormal regions in the system. To account for the increased catalytic activity. these abnormal areas must be associated with catalytic sites. It bvould therefore seem illogical to expect these abnormal areas to show the adsorption characteristics one would measure for the gross surface. The correspondence of adsorption on catalytic sites and on silver oxide could indicate that the catalytic site involves a promotor oxide. Adsorption on a stable oxide such as potassium oxide might easily be of the same nature as that shown by the catalytic site. The authors therefore are not in agreement with earlier interpretations proposed by Tlvigg and Kummer, which were based on the assumption that the catalytic area must necessarily exhibit the same adsorption characteristics as metallic silver. For silver-catalyzed oxidation of ethylene the inhibition of the reaction by carbon dioxide is in itself sufficient to cast doubt on 290
l & E C P R O C E S S DESIGN A N D DEVELOPMENT
this assumption. This inhibition has been reported in a number of studies, and since carbon dioxide does not adsorb on silver at the reaction temperature, the catalytic site must have adsorption characteristics differing from those of metallic silver. Since the above work was finished, Tamaru (79),who studied the catalytic decomposition of formic acid on silver and nickel and on copper, concluded that adsorption on the gross catalyst surface and adsorption on the catalytic site did not correspond. H e measured the effect of formic acid pressure upon adsorption on the catalyst and upon the catalytic rate simultaneously. I n the case of silver he obtained a correlation between the fraction of the surface covered by adsorption and apparent fraction of catalytic sites covered. In the case of copper, however, adsorption was first order with formic acid partial pressure but the catalytic rate was independent of it, indicating that the catalytic site shows different adsorption characteristics than the gross surface and that the catalytic sites of the copper catalyst occupy a small part of the surface and have a much higher heat of adsorption than the gross surface. Probable Catalytic Mechanism. The treatment of these data to test various one- and two-site Langmuir-type catalytic rate equations has been discussed in detail (6, 74). The three mechanisms for the produclion of ethylene oxide which are not eliminated by the data are all two-site mechanisms:
1, Two molecules of ethylene each occupy one site. 2. One site is occupied by an ethylene molecule, the other by an oxygen molecule, and the sites are competitive. 3. One site is occupied by an ethylene molecule, the other by an oxygen molecule, and the sites are noncompetitive. Quantitative Evaluation of Rate Equations and Comparison with Integral Reactor Data
Following the qualitative conclusions reached above, constants of the resulting equations were evaluated by using the differential data. These equations were then tested against integral reactor data obtained using the same apparatus. Comparison was also made between the equation forms and the data of Orzechowski and MacCormack, extrapolated to zero contact time. Carbon dioxide lowers the rate of oxidation of ethylene both to ethylene oxide and to carbon dioxide plus water. It is therefore necessary to predict the rate of formation of carbon dioxide from ethylene in order to predict the rate of formation of ethylene oxide. Equation 1 was developed for this purpose. I t is not suggested as representing the actual mechanism but only as a convenient expression which makes it possible to allow for rate suppression by carbon dioxide. No attempt was made to include in the calculations the loss of ethylene oxide due to oxidation, nor the carbon dioxide formed from this reaction.
Of the three plausible mechanisms for ethylene oxide formation, equations were developed for the two which best fit the experimental data. Equation 2 represents Mechanism 1, in which two adjacently adsorbed ethylene molecules react with gaseous oxygen to form ethylene oxide. Equation 3 represents Mechanism 2, reaction between adjacently adsorbed ethylene and oxygen molecules. 4980 (CzH4)' (
Rate CnHaO Rate CnHtO
=
=
[l
0 2 )
+ 164 (CzHa) + 269 (C02)lz
(2)
14.7 (CzHa) (On)
[I
+ 15.3 (C?H4) f
1.26 (
0 2 )
f 35 (COz)]' (3)
I
0
002
004
006
008
002
010
0'04
0'06
0.b8
% \.
10
0.10
E T H Y L E N E PARTIAL P R E S S U R E , A T M
E T H Y L E N E PARTIAL PRESSURE, A T M
0 0
0.1
0.2
0.3
0.4
0
O X Y G E N PARTIAL PRESSURE, A T M
Figure
9.
Test of Equations 2 and
3 with
data
of Bolme
In each of these equations no product concentrations are in the numerator because of irreversibility. With the exception of carbon dioxide, the reaction products were assumed to have no effect upon the rate of the reaction such as would occur if they were adsorbed on the catalyst. The concentrations are in atmospheres and the rates in terms of gram-moles of ethylene reacted to the product shown per bulk cubic centimeters of catalyst per hour. The data of Orzechowski and MacCormack were also used to test the form of Equations 1, 2, and 3. No attempt was made to use the equation constants obtained in this work because of differences in catalyst activity and temperature level. The three equations are tested with data of Bolme in Figures 9 and 11 and with data of Orzechowski and MacCormack in Figures 10 and 11. Equations 2 and 3 plotted in linear form appear to fit both sets of data equally well. Integral Rate Data. Integral rate data were obtained (74) using essentially the technique and apparatus of the differential rate work. Conversion levels were of course much higher. Equations 1, 2, and 3 were fitted to these data as follows:
Figure 10. Test of Equations 2 and 3 with data of Orzechowski and MacCormack a t 2 7 4 " C.
. 2
2
\
::~
a 200 \
5
6
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Integration of Equations 1, 2, and 3 using the form d ( W / F ) = dx/r integrated between appropriate limits of reciprocal
space velocity, W/F, and conversion, x . Units for W / F (catalyst weight over flow rate) must be consistent with units of x and r (conversion and reaction rate) for such integrations. Solving appropriate pairs of Equations 1 and 2 or 1 a n d 3 simultaneously, by incremental trial and test methods, by assuming a conversion for a n increment, adjusting mean concentration levels for this conversion, and calculating the conversion from the integrated equations to check with the assumed value. Typical results for Equations 2 and 3 are shown as solid lines in Figures 12 and 13. Experimental data are shown as points. The results for low levels of COz (Figure 12) are in reasonable agreement with the data u p to hourly reciprocal space velocities of about 2.5. At higher conversions (higher W / F ) , however,
0
0
004
008
OXYGEN PARTIAL
012
016
IO 028
P R E S S U R E ,ATM
Figure 1 1 . Test of Equation 1 with data of Orzechowski and MacCorrnack a t 274" C. and of Bolrne
the experimental yield of ethylene oxide is less than predicted. This is believed to be caused by further oxidation of the ethylene oxide a t these higher concentrations, not considered in this study. Agreement between experimental results and the integrated equations is much poorer for runs in which a large quantity of carbon dioxide was added with the feed (Figure 13). This suggests that the carbon dioxide suppression correction is too VOL. 1
NO. 4
OCTOBER 1962
291
$ 30‘
3
RUN 15. 115% CtH, , 2 5 % 0, I a Percent ethylenr oxidized to Cot 0 Percmt ethylene oxidized to E T 0
0
Percent #lhylene oxidired to ET0
/
w/F
~
large a t these levels of carbon dioxide concentrations. In addition, Equations 1, 2, and 3 should probably contain terms for suppression of reaction rate by the other reaction products, water and ethylene oxide. Such suppression has been reported by other workers (8, 72, 76). These terms would allow a much better fit of the equations to the data of Figure 13 (as found by qualitative test) without necessarily altering the fit to data such as shown in Figure 12. Since no experimental work was done to test the possible suppression of reaction rate by these products, no further discussion of this point appears justified here. Catalyst Stability. Catalyst stability was tested to ensure that activity remained reasonably constant during the time required for a run. An extended run was made with conditions held constant for 28 hours. With inlet feed composition of 8.8% ethylene, 31.8% oxygen, remainder nitrogen, and a W)‘F ratio of 2.50 hours, concentration of ethylene oxide a t the reactor outlet remained between 1.5 and 1.65%. Carbon dioxide concentration remained between 1.7 and 1.85%. The slight fluctuations in yield appeared random. Conclusions
A comparison of published adsorption data and reaction rate data of this work and previously published results suggests no correspondence between adsorption of reactants on silver surfaces and apparent adsorption of the same reactants on the catalytically active areas. In this the present interpretation differs from that proposed by Twigg. There is much closer correspondence betwen adsorption on silver oxide surfaces and behavior of reactants as reflected in the reaction rate equations. Ot!ier evidence also suggests that the active catalyst is partially oxidized, possibly a t the silver-silver oxide or silver-promoter o?-ideboundaries. Differential reactor rate data are particularly convenient for the straightfonvard testing of possible rate-controlling steps and reaction rate mechanisms. Reaction rate equations evaluated in differential form with such data are consistent with integral reactor data below the conversion level at which ethylene oxide oxidation becomes appreciable. This suggests that gradients established in small (3/c-inch diameter) integral reactor beds are not so great as to invalidate such data for reactor design purposes. I n particular initial rate data obtained from such
292
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PROCESS DESIGN A N D DEVELOPMENT
/’
(REACTION TIME COORDINATE), h&
reactors should be reliable. The experimental data of Tivigg and of Orzechowski and MacCormack are consistent with the data and semitheoretical equations proposed in this \vork. The form of rate equation suggested by Kurilenko, Kul’kova. Rybakova, and Temkin is similar to those of this liork. literature Cited
Benton, A. F.,Drake, L. C.,J . A m . Chem. Sor. 54,2186 (1932). Ibid., 56, 255 (1934). Ibid.. D. 506. Bentoh, A. F.,Elgin, J. C., Ibid., 48,3027 (1926). Ibzd., 51, 7 (1929). Bolme, D. TV., Ph.D. thesis in chem. eng., University of Washington, Seattle, TVash., 1957; Dissertation Abst. 13, 1362 (1958) : microfilm. Universitv Microfilms. Ann Arbor. Mich. 17)‘ Constable. F. F..’Proc. Rov. koc. 108A. 355 11925). (8j GorokhovatskiY,’Ya. B.: Rubanik,’ M. ‘Ya., ’Kholyavenko. K. LV., Doklady A k a d . .\hukS.S.S.R. 125,83-6 (1959). (9) Hougen, 0. A,, Watson, K. M., “Chemical Process Principles,” Vol. 111, “Kinetics and Catalysis,” Wiley, New York,
(1) (2) (3) (4j (5) (6)
1947. (10) Kummer, J. T., J . Phys. Chem. 60, 666 (1956). 111) Ibid.. 63. 460 11959). (12j Kurilenio, A ‘ I . , KLl‘kova. N. V., Rybakova. N..A,. Temkin, M. I.. Zhur. F u . Khzm. 32, 797, 1043 (1958). (13) Lewis, G. N., Z. phys. Chem. 52, 310 (1905); Proc. A m . Acad. Arts Sa. 40, 719 (1905); J . A m . Chem. Sac. 27, 601 (1905) . , .
I
(review).
(14) Nault, L. G., M. S. thesis in chem. eng., University of TVashinzton. Seattle. TVash.. 1959. (15) OrzGhowski, A,,’ MacCormack, K. E., Can. J . Chem. 32, 388, 415, 432, 443 (1954). (16) RoginskiY, S. Z., Margolis, L. Ya., Doklady A k a d . ‘Va’auk 89, 515 (1953). (17) Schwab. G. M., Taylor, H. S.,Spence, R., “Catalysis,” Chau. 12. Van Nostrand. New York. 1937. (18) Sieacie, E. TV. R., Johnson, F. M. G., Proc. Roy. Sac. 112A, 542 (1926). (19) Tamaru, K., Trans. Faraday Sac. 55,824, 1192 (1959). (20) Twigg, G. H., Proc.Roy. Soc. 188A,92 (1946). (21) Zbzd., p. 105. (22) Zbzd., p. 123. (23) Twigg, G. H.. Trans. Faraday SOC. 42, 284, 657 (1946). (24) TVan, S. IV., IND.ENG.CHEM. 45,234 (1953).
RECEIVED for review July 21, 1961 ACCEPTED March 12. 1962 Division of Industrial and Engineering Chemistry, 137th Meeting. ACS, Cleveland, Ohio, April 1960. Major portion of work supported by Office of Ordnance Research, United States Army. Texas Co. provided fellowship support of a portion of Nault‘s work.
