It was shown that a number (h’)of perfect-mix permeators in series approaches the performance of one no-mix permeator with the same total area. I n the extreme as N + the performance of the series of perfect-mix permeators should be identical to the no-mix model. However, with N = 8 the correspondence was quite close, and this indicates that proper baffling on the high-pressure side of the permeator would result in a performance approaching the no-mix model. Finally, the study of a two-stage permeator with both stages operating a t the same temperature indicates that the area requirements are much greater, the interstage stream that must be compressed is greater and enrichment is less for the perfect-mix case than the no-mix case, further emphasizing the need for correct flow characteristics in a permeator. Acknowledgment
The assistance of Chemical Projects International, Rome, Italy, and its Xanaging Director, R. G. Minet, in carrying out this research is gratefully acknowledged. Nomenclature
membrane area, ft2 flux through membrane, scfd/ft2 volume rate of high-pressure stream, scfd L volume rate of low-pressure stream (permeate), scfd P H ~PCO , = pure gas permeability coefficients, cm3(STP) . cm/sec . cm2 cm H g
A F H
= = = =
+
W = mole fraction CO in high-pressure stream Y = mole fraction CO in low-pressure stream or permeate dZ/dA = flus of CO only through membrane, a t any point mole fraction H2 product/ mole fraction CO
j
mole fraction Hz feed = actual local separation factor mole fraction CO (rrhere the compositions are measured at the same position coordinate) literature Cited
Atkinson, R., “‘Permasep’ Permeators for Hydrogen Separation,’’ Tech. Bull., E. I. DuPont de Nemours & Co., Inc., 1970. Franks, R. G. E., “ A Digital Computer Program for Simulating Unsteady State Processes,” Engineering Dept., E. I. DuPont de Nemours & Co., Inc., 1970. Levenspiel, 0. “Chemical Reactor Engineering,” Chap. 6, Wiley, Kew York, N. Y., 1962. RIichaels, H. S., Chem. Eng. Progr., 64 (12), 31-43 (1969). Oishi, J., Rlatsumura, Y., Hagashi, K., Ike, C., U.S. Atomic Energy Commission, Report No. AEC-TR-5134, 1961. Stern, S.A., Sinclair, T. F., Gareis, P. J., Vahldieck, K. P., Mohr, P. H., Ind. Eng. Chem., 57 (2), 49-60 (1965). Stern, S. A., Gareis, P. J., Sinclair, T. F., Alohr, P. H., J . A p p l . Polym. Sci., 7, 2035-51 (1963). Stern, S. A,, Walawender, W. P., Separ. Sci., 4(2), 129-59 (April 1969). Weller, S., Steiner, W. A , , Chem. Eng. Progr., 46 ( l l ) , 585-90 (November 1950). RECEIVED for review September 1, 1971 ACCEPTEDJuly 11, 1972
Kinetics of Silver-Catalyzed Ethylene Oxidation Peter 1. Metcalf,l and Peter Harriott Cornell University, Ithaca, N.Y. 14860
The kinetics of silver-catalyzed ethylene oxidation were studied using a differential reactor. Behavior of various inhibitors including ethylene oxide, carbon dioxide, water, and dichloroethylene was investigated at a number of reactant partial pressures. The rates of both ethylene oxide and carbon dioxide formation passed through maxima with increasing oxygen pressure. The inhibiting effects of carbon dioxide and water appeared to be rapidly reversible and to follow a noncompetitive rate law. The inhibiting effects of ethylene oxide and dichloroethylene were only slowly reversible. These results were not completely consistent with Langmuir-Hinshelwood type rate expressions.
E t h y l e n e is easily oxidized with a silver catalyst. Two products are produced in this reaction: ethylene oxide and carbon dioxide.
0
CH?=CH?
+ 302
2C02
+ 2H20
(11)
Earlier studies by Bolme (1957), Buntin (1961), and Klugherz and Harriott (1971) have shown t h a t the reaction 1 Present address, E. I. du Pont de Xemours & Co., Inc., Wilmington, Del. 19898. To whom correspondence should be addressed.
478
Ind. Eng. Chem. Process Des. Develop., Vol. 1 1 , No.
4, 1972
is approximately first order in ethylene and three-halves order in oxygen a t low partial pressures of ethylene and oxygen. The rates of both reactions increase as the pressure of ethylene is increased, pass through maxima, and then decrease. A maximum in the rate of carbon dioxide formation with increasing oxygen pressure was suggested by Harriott (1971), but no maximum in the rate of ethylene oxide formation with increasing oxygen pressure has been reported. Buntin (1961) and Klugherz and Harriott (1971) found that the reaction products-ethylene oxide, water, and carbon dioxide-inhibit the reactions. liargolis et al. (1962) has shown t h a t various other elements more electronegative than silver-such as bismuth, sulfur, and chlorine-also inhibit the reactions.
Klugherz and Harriott (1971) explained these results b y a kinetic model in which the reaction takes place between a n ethylene species and a n oxygen species, both of which are competitively adsorbed on top of a layer of oxygen chemisorbed by the silver catalyst. I n this model, the reaction products are also adsorbed by the chemisorbed oxygen layer thus inhibiting the reaction. I n this study, a differential reactor was used to measure the effects of carbon dioxide, water, ethylene oxide, and dichloroethylene on the rates of ethylene oxide and carbon dioxide formation. ;Ilso, the effect of oxygen was determined over a much wider range of partial pressures than in a n y previous work. The data were fitted to Langmuir-Hinshelwood rate expressions, although some inconsistencies were noted. Further discussion of the results and alternate theories is given in the thesis byhletcalf (1971).
-
I
I
I
I
15.0
W
0) L
5.0 \
9
c
10% Conversion) I
0.01
I
I
0.02
I
0.03 0.04
2, (Atm) Figure 1. Effect of conversion on rate E t 0 formation
Apparatus
A Cop formation
The ethylene oxidation reaction was carried out in a packed-bed reactor consisting of a in. o.d., 0.430 in. i d . , piece of stainless steel tubing packed with a mixture of 10.0 grams of 20-28 mesh a-alumina-supported silver catalyst and 15.0 grams of 20-30 mesh Ottawa sand. T h e length of the bed was approximately 71/2 in. long. A '/8 in. thermowell was held in the center of the reactor, and the reactor centerline temperature was measured by a thermocouple which fit tightly inside. Reaction temperature mas controlled by immersing the reactor in a n electrically heated fluidized bed of sand. The fluidized bed temperature was controlled a t 219" 0.2"C by a temperature controller. T h e average temperature rise between the fluidized bed and the reactor center was usually less than 1-2"C, although peak temperature rises as high as 8°C were occasionally observed. The feed to the reactor contained ethylene, oxygen, helium, and various trace gases. The helium was used as a diluent. The feed gases were metered from their cylinders with precision needle valves. The gas flow rates mere measured with either rotameters (at high flow rates) or capillary flow meters packed with sand (at low flow rates). K a t e r was added to the feed with a saturator. The total feed rate to the reactor varied between 25 and 1000 scc/min. At 1000 sccimin, the pressure drop across the reactor was about 0.03 atm. The rates of Reactions I and I1 were determined by measuring the composition of the product stream with a gas chromatograph equipped with a thermal conductivity detector. A '/8 in. 0.d. 6-ft-long column packed with 100-120 mesh Poropak Q was used to separate the components. The catalyst used in this study was approximately 8.1% silver supported on Alcoa T-71 a-alumina with a surface area of 0.5 m2/g. The catalyst was prepared in the same manner as t h a t used by Klugherz and Harriott (1971).
*
Experimental Procedure
The reactor was operated with lorn ethylene conversion, but the inhibiting effects of the products plus changes in the catalyst activity required a special procedure to get consistent kinetic data. The effect of the products was evident from a large decrease in reaction rate t h a t followed immediately after a decrease in feed rate. The change could not be explained by the slight decrease in average reactant pressures or b y evternal mass transfer. Figure 1 shows t h a t the rates of formation of ethylene oxide and of carbon dioxide follow the simple kinetic laws
PEO = 0.259 atm pori
= 0.250
otm
where 2 , is the sum of the average product concentration generated by the reaction. The catalyst activity gradually changed with time. The selectivity of the reaction ordinarily decreased as the catalyst activity increased. The catalyst activity usually increased during operation a t low conversions and high temperatures, and decreased during operation a t high conversions and low temperatures. These changes in catalyst activity did not appear to affect the coefficients a1 and b1 in Equations 1 and 2. They had the values 50 atm-I and 100 atrn-l for standard runs regardless of catalyst activity. The fact that al and bl were constant simplified measurement of the catalyst activity. The rate could be calculated from the conversion of ethylene to form ethylene oxide and carbon dioxide under standard operation conditions: pEt = p o 2 = 0.263 atm; feed rate = 400 sccm. The measured rate was then corrected using Equations 1 and 2 to the rate which would have been measured if the feed rate had been adjusted so that 2, = 0.01 atm. The experimental procedure consisted of alternating kinetic runs made a t the feed composition and flow rate being studied with standard runs made a t standard operating conditions. The rates measured a t the standard conditions rrere corrected to 2, = 0.01 a t m using Equations 1 and 2. The rates measured for the kinetic runs were divided by the average of t h e corrected rates obtained from the bracketing standard runs to obtain the relative rates (denoted R E ~ O and R c ~ ? , respectively). Relative rates and corresponding 2,'s mere measured a t a t least four different flow rates for a feed gas with a constant composition. The reciprocals of the measured relative rates were plotted against 2, and a straight line was drawn through the data to obtain the relative rates corresponding to 2, = 0.01 a t m (denoted ROE~O and ROco,) I
Experimental Results
Effect of High Oxygen Pressures. T h e oxidation of ethylene was studied at high oxygen (1.31-13.1 a t m ) a n d low ethylene (0.026-0.062 a t m ) pressures t o test for t h e existence of a maximum r a t e a n d t o see if t h e selectivity is markedly affected b y these extreme operating conditions. Ind. Eng. Chem. Process Der. Develop., Vol. 11, No. 4, 1 9 7 2
479
1 .o
.,
RoEtO
0.5
P 0 p ( At m.) R O ,
Figure 2. Effect of high oxygen pressures on rate of ethylene oxide formation
.
-\
1
1
,
I
I
I
I
I
I
,_1
0.2 0.3 PCO, ( A t m . )
0.1
I
1
I
I
I
I
Figure 4. Effect of carbon dioxide on oxidation rates atm
atm
0.250
0.004
Table 1. Interaction Between Carbon Dioxide and Water: Ethylene Oxide Rate Dataa
0 004 0 014 0 024 a
Data for pEt
10.0
Pg,(Atm.) PH~O, atm
\
=
0 80 i 0 07 0 83 i 0 07 0 81 k 0 07 0.259 atm,
RoEto(pco2= 0.024 atm) R0~to(p~o2 = 0.004 atm)
0 71 0 74 0 74
PO? =
=t0
06
* 0 07
* 0 07
0.250 atm
Rocozlpco?= 0.014 otm) RaCo2(pCo2= 0.004 atm)
R0co2(pco2= 0.024 atm) Rocol(pco, = 0.004 otm)
0 55 i 0 05 0 50 i 0 05 0 52 i 0 05 0 2.59 atm, PO?
0 42 i 0 04 0 39 L 0 0.1 0 39 0 04
0 004 0 014 0 024
Figure 3. Effect of high oxygen pressures on rate of carbon dioxide formation PEt, a t m
Ind. Eng. Chem. Process Des. Develop., Val. 11, No. 4, 1972
RoEto(pco2= 0.014 atm) RO~to(p~o, = 0.004 a t m )
Table 11. Interaction Between Carbon Dioxide and Water: Carbon Dioxide Rate Data.
111115.0 1111111111
A 0.062
PH20t
atm
0.259
The possibility of interaction between water and carbon dioxide was studied by measuring the rate of oxidation as a function of carbon dioxide pressure a t various levels of water vapor pressure. The results of this study are summarized in Tables I and 11. The dat'a in these tables indicate that, \Tithin the limits of experimental error, the ratios of the relat'ive rates a t two different' levels of carbon dioxide pressure are independent of water pressure. This would indicate noncompetitive inhibition, but the results are not conclusive. The competitive adsorption model, Equation 9, predicts that the ratio in the last column of Table I would change from 0.73 to 0.78 over the range of water pressures given, and this difference is about the same as the standard deviation of the measured ratios. There is better evidence against competitive inhibit,ion of carbon dioside formation, with a predicted change in ratio from 0.48 to 0.56, compared to a decrease from 0.42 to 0.39 as shown in Table 11.
I
0 0.026
PO?,
PEt,
p ~ ? oatm ,
480
I
0.30
,
The results of these experiments are shown in Figures 2 and 3. The rates of oxidation to ethylene oxide and to carbon dioxide pass through maxima in the range of 2-5 a t m of oxygen pressure. These maxima shift to higher oxygen pressures as the pressure of ethylene is increased. A t the highest oxygen pressures, both reactions are about first order in ethylene and -0.3 order in oxygen. The selectivity for ethylene oxide formation also goes through a maximum in the range of 2-5 a t m osygen pressure and then starts to decrease. Effect of Carbon Dioxide Pressure. Carbon dioxide was added in amounts varying from 0.004 t o 0.36 a t m to a feed with constant partial pressure of 0.263 a t m ethylene and 0.263 a t m oxygen. T h e rate of oxidation both t o ethylene oxide and to carbon dioside decreased rapidly with the addition of increasing amounts of carbon dioxide (Figure 4). T h e selectivity increased rapidly a t first, but approached a constant value a t high carbon dioxide pressures. The rate of ethylene osidation t o form ethylene oxide fell off with increasing carbon dioxide pressure to the -0.25 power a t low carbon dioxide pressures and to the -0.5 power a t high pressures. The rate of oxidation to form carbon dioxide fell off ivith the -0.45 power of carbon dioxide pressure throughout the range of pressures studied. The lines in Figure 4 are based on Equation 9, which predicts much too low values a t high carbon dioxide pressures. I
I
0.20
I':
PEt, O t m 0 0.026 A 0.062
2.01
I
0.10
a
Data for p~~
=
*
=
0.250 atm
I
I
I
I
I
I
Table 111. Effect of Carbon Dioxide on Rate of Ethylene Oxidation at Various Reactant Pressures. p ~ t atm ,
0.062 0.062 0.259 1.32 a Data
RE~O(PCO* = 0.024 atm) po2, atm RE&CO~ = 0.004 atm)
0.052 1.31 0.250 0.055 for
PH?O
0.72 i 0 . 0 6 0.70 0.06 0.71 i 0.06 0 . 9 2 + 0.08 0.004 atm.
*
=
F!coI(pco2 = 0.024 atm) Rco,(pco2= 0.004 otm)
0.48 i 0 . 0 6 0 . 4 2 =k 0 . 0 5 0.41 =k 0.05 0.95 0.11
0.01
*
I
The interaction among carbon dioxide and the reactants was studied by measuring the rate of oxidation with 0.02 a t m carbon dioxide added to the feed at various levels of oxygen and ethylene partial pressures. The results of these experiments are tabulated in Table 111.Increased ethylene pressures appear to reduce the relative effect of carbon dioxide, suggesting t h a t ethylene and carbon dioxide compete for the same sites. Increased oxygen pressures do not affect the relative effect of carbon dioxide addition significantly, though little effect is expected at these pressures, xhere the reactions are nearly first order to oxygen. Effect of Water Vapor Pressure. T h e effect of water on t h e oxidation of ethylene was tested by adding small amounts (0.010 a n d 0.020 atni) of water t o t h e reactor feed. F o r feed compositions of 0.263 a t n i ethylene and 0.263 a t m oxygen, t h e rate of ethylene oxide formation declines in proportion t o t h e -0.25 power of t h e average water vapor pressure, \vhile t h e rate of carbon dioxide formation falls off with t h e -0.20 power of the average water vapor pressure (see Figure 5). The interaction among water and the reactants n-as tested b y measuring the rate of oxidation with 0.02 a t m water added to the feed at various levels of oxygen and ethylene pressures. The results of these experiments are tabulated in Table IV. The effect of oxygen and ethylene is similar to that found in the study of the interaction among carbon dioxide, oxygen, and ethylene. The addition of ethylene reduces the inhibiting effect of water. Greater oxygen pressure has little effect on the inhibition of ethylene oxide formation by water, but may even enhance the effect of water on carbon dioxide formation. Effect of Ethylene Oxide Pressure. T h e addition of 0.005 a t m of ethylene oxide t o t h e standard feed gas caused a n immediate 50% reduction in t h e r a t e of oxidation t o ethylene oxide, b u t almost no reduction in t h e rate of oxidation t o carbon dioxide. T h e catalyst activity then trended steadily downward over a period of 3 h r until almost no ethylene oxide or carbon dioxide was being formed. K h e n t h e ethylene oxide addition was stopped, t h e rate of ethylene oxide formation jumped almost back to the original value while there mas little immediate change in the rate of
PEt, atm
PO*,atm
0.062 0.062 0.259 1.32 a Data
0.052 1.31 0.250 0.055 for
PCO, =
RE&H~O = 0.024 atm) R C O ~ ( P H ~ O= 0.024 atm) REto(pH20= 0.004 atm) R C O * ( P H ~ O= 0.004 otm)
0.60 i 0.05 0.54 i 0.05 0.66 0.06 0.88 i 0.08 0.004 atm.
*
0 . 7 3 =& 0.07 0 . 4 7 =t0 . 0 5 0.72 I 0 . 0 7 0 . 8 8 =t= 0 . 0 9
I
I
I
I
I
Figure 5. Effect of water on oxidation rates 0 pco, = 0.004 atm
A pco? = 0.014
atm
pcog = 0.024 atm P E ~= 0.259 atm PO, = 0.250 atm
carbon dioxide generation. N o data on the relative rate of ethylene oxidation as a function of ethylene oxide partial pressure was gathered since the ethylene oxide apparently reacted slowly to inactivate the catalytic sites rat'her than being reversibly adsorbed on the sites as in the case of the reactants and the other product gases. Holyever, the reversible effect of the combined products could be accounted for by the presence of carbon dioxide and water. Effect of Dichloroethylene Pressure. Dichloroethylene (DCE) is a strong inhibitor for t h e silver-catalyzed osidation of ethylene. Figure 6 illustrates t h e effect of adding several different concentrations of DCE to t h e standard feed. T h e rate of ethylene oxide formation falls rapidly. T h e rate of carbon dioxide (not, shou-n) follows a siniilar pattern. As the Oxidation rates fall, selectivity is gradually increasing. Selectivities as high as 75% were obtained as t h e activity approached zero, compared with the normal
1 t
!i0.02
v)
v
02
Table IV. Effect of Water on Rate of Ethylene Oxidation at Various Reactant Pressures.
0.02
PH,O ( A t m . )
8 .-
O.O' 0.005
e
DCE Added
0.002
2
(Standard Operating
4
6
Time
(Hrs.)
8
Figure 6. Transient response of t h e rate of ethylene oxide formation to dichloroethylene addition 0160ppbDCE B320 ppb DCE A 640 ppb DCE Ind. Eng. Chem. Process Des. Develop., Vol. 1 1, No. 4, 1972
48 1
Table V. Effect of 80 Ppb DCE Modification on Rate of Ethylene Oxidation
--.
R'rtn
ROco,
80 p p b
80 PPb p ~ t atm ,
0.062 0.062 0.062 0.062 0.258 0.062 0.259 0.259 0.259
pox, atm
pco2, atm
Normal
DCE
Normal
Increasing oxygen pressure 0.052 0.0042 0.24 0.47 0 . 3 6 0.249 0.0036 0 . 5 3 0 . 8 6 0 . 7 6 1.31 0.0043 1 . 3 0 1 . 9 6 1.10 Increasing ethylene pressure 0.052 0.0042 0.24 0 . 4 6 0 . 3 6 0.054 0.0046 0.21 0.24 0 . 2 8 0.249 0.0036 0 . 5 3 0 . 8 6 0.76 0.250 0.0037 1.00 1 . 0 0 1 . 0 0 Increasing carbon dioxide pressure 0.250 0.0037 1 . 0 0 1.00 1 . 0 0 0.250 0.0244 0.71 0.71 0.42
DCE
0.80 1.07 1.33 0.80 0.13 1.07 1.00 1.00 0.43
selectivity of 45%. T h e catalyst activity could be restored by feeding the standard feed overnight while raising t h e reactor temperature to 240OC. The rate of decline of catalyst activity varied considerably depending on the concentration of D C E in the feed. A t 640 ppb D C E , the rate of carbon dioxide formation fell a t a rate of approximately 90'%/hr. A t half that concentration, 320 ppb, the rate fell a t 30%/hr; a t 160 ppb, the rate of decline was hard to detect. Several factors may have contributed to this highly nonlinear effect. The amount of D C E in these tests was so low that little of the surface mas covered. Adsorbed chlorine atoms undoubtedly deactivated sites adjacent to the sites on which they were adsorbed. The degree of this type of inhibition would not be expected to be a linear function of the amount of chlorine adsorbed. The rate of oxidation of ethylene was determined for several pressures of ethylene, oxygen, and carbon dioxide on a D C E modified catalyst (Table V). The catalyst was modified by feeding 100 sccm of gas containing 0.263 a t m ethylene, 0.263 a t m oxygen, and 80 ppb D C E overnight. The activity declined until 2 , = 0.01. Since the feed rate was 25y0 of the usual standard feed rate, the activity of the modified catalyst was about 25'%.of t h a t of the unmodified catalyst. It should be noted that it is not clear how much of this reduction in activity is due to the addition of D C E , and how much is due to the prolonged operation of the reactor a t high product pressures. The relative rates of reaction for the various feed and product conditions were determined in the same manner as with the unmodified catalyst except that the alternate runs a t standard conditions were replaced with alternate runs a t the modified standard feed conditions of 100 scc/min of gas with 0.263 a t m ethylene, 0.263 a t m oxygen, and 80 ppb D C E . T h e standard conversion was still defined so that Z P = 0.01 atm. A flow of 80 ppb of D C E was maintained both for the kinetic runs and for the modified standard runs. The selectivity of the modified standard runs varied from 45 to SO%, but was usually not much higher than the selectivity for the normal standard runs. I t is apparent from Table V t h a t D C E modification of the catalyst reduces the activity of the catalyst relatively more a t high ethylene and oxygen pressures than a t low pressures: the rate maxima would be shifted to the lower pressures. Carbon dioxide has the same relative effect on the reaction rate with either the modified or the unmodified catalyst. 482 Ind. Eng. Chem. Process Des. Develop., Vol. 11, No. 4, 1972
This behavior indicates that the sites which remain active after inhibition adsorb oxygen and ethylene more strongly than the sites on the unmodified catalyst. Such behavior is contrary to that which would be expected from a catalyst which obeyed a Langmuir-Hinshelwood type rate equation. A Langmuir-Hinshelwood type rate equation would predict that the rate maxima would be shifted to higher ethylene and oxygen pressures on the addition of a n inhibitor if the reactants and the inhibitor are competing for the same catalytic sites. Analysis of Results. T h e Langmuir-Hinshelwood theory of heterogeneous catalysis was used to summarize much of the kinetic d a t a gathered in this study. A model was proposed which explained the following observations : (a) The rate goes through a maximum with increasing oxygen pressure and with increasing ethylene pressure. Klugherz and Harriott (1971) showed that the maxima shift to higher ethylene pressures as the pressure of oxygen increases. This indicates that the rate-determining step involves a bimolecular surface reaction with competitive adsorption of oxygen and ethylene. (b) Klugherz and Harriott (1971) also showed that the reaction order with respect to oxygen is greater than one a t high ethylene pressures and low oxygen pressures. (e) The inhibiting effect of water and carbon dioxide decreases as the ethylene pressure increases, suggesting t h a t ethylene and the products compete for the same sites. (d) Ethylene, carbon dioxide, and water are adsorbed only on oxygen-covered silver-not on bare silver (Bolme, 1957). (e) Flank (1965) demonstrated that most of the oxygen adsorbed by silver does not evaporate a t ordinary reaction temperatures and pressures, but can be removed by reaction with ethylene. (f) Changes in the reactant and product pressures have very nearly the same effect on the rate of formation of carbon dioxide as they do on the rate of formation of ethylene oxide. The selectivity of the reaction for the formation of ethylene oxide varies in the range of about 3@-60% throughout the wide range of pressures studied. The mechanisms of the two reactions seem to be closely related, and it seems appropriate to use the same kinetic model in the analysis of each. The fraction of the silver surface consisting only of bare catalytic sites may be denoted 8,. Three types of oxygen are assumed to be adsorbed on the silver surface: Species of oxygen
Fraction of s u r f a c e covered
Schematic representation
/"\
atomic
e1
molecular
e2
Ag
Ag
o=o adsorbed These species are assumed to be a t equilibrium with each other so t h a t their concentrations on the surface are related to each other by 03 822
Kip0201
(3)
K@iOg
(4)
=
Ethylene is also reversibly adsorbed by the oxygencovered surface:
s
h
v
h d
v r3
so t h a t the concentration of ethylene on the surface is given by 5~ = KEPE&
(5)
Atomic oxygen may gradually be removed from the surface by reaction with adsorbed ethylene as well as by recombination to form molecular oxygen CH2=CH2 Ag/"\Ag
+ Ag/"\Ag
k7c
and
.. ..
0 I
Ag/"\Ag
-t
Ag/"\Ag
-%
Atomic oxygen is added to the surface by this s l o ~reaction between molecular oxygen and the bare sites
..
0
The direct reaction of oxygen with bare sites t o form atomic or molecular oxygen is neglected, although it would be important at low surface coverages. These nonequilibrium reactions for removal of oxygen are assumed to be much slower than the equilibrium reactions or the rate-determining steps. By making the steady state approximation for the atomic oxygen species, one obtains k&30 = ki5&
+ k&2
(6) Both carbon dioxide and water are reversibly adsorbed on the atomic oxygen: O\
C
P
I
/"\
H
H
\
/"\
H , O + Ag
Ag
K h
\\
I /
/"\
Ag
H
I
H
P-"\ Ag
Ag or Ag
The fractions of the surface covered by adsorbed carbon Ind. Eng. Chem. Process Des. Develop., Vol. 1 1 , No.
4, 1972 483
dioxide and by adsorbed water are given by ec and eh, respectively. These fractions may be computed with the equilibrium relationships:
(7)
Be = KcpCO2el
and
The fraction of the surface covered by each of these molecular species may be written in terms of the concentration of atomic oxygen and the partial pressures of the reactants and products. These equations may be combined with the condition that the sum of the fractions of the surface covered by each of these species is equal to 1.0 to obtain a series of equations giving the fraction of the surface covered b y each of the species in terms only of the partial pressures of the reactants and the products. ii kinetic model in which the rate-determining step was a bimolecular surface reaction between adsorbed ethylene and molecular oxygen fit most of the kinetic data quite well for both ethylene oxide formation and the formation of carbon dioxide.
One significant difference between the experimental data and the predictions occurred at the highest carbon dioxide pressures where the model predicted that the ratio would fall off more rapidly with increasing carbon dioxide pressure than was actually observed (see Figure 4). These data were not included in the regression analyses. Another difference was that the relative inhibiting effects of water and carbon dioxide were almost independent of each other, whereas the model predicts t h a t they compete for sites with each other as well as with the reactants. Finally, the model does not describe the slow changes in reaction rate t h a t occurred when ethylene oxide and D C E were added to the feed. These shortcomings may be due to heterogeneity of the catalyst surface or to changes in the electronic properties of the catalyst. Nomenclature a,
k,
= =
constant rate constants
K1, K z , K B , K,, K h = equilibrium constants pco2 = average carbon dioxide pressure, atm p~~ = average ethylene pressure, atm PEto = average ethylene oxide pressure, a t m pn20 = average water pressure, atm po2 = average oxygen pressure, atm %to = average rate of conversion of ethylene to form ethylene oxide, scc min-l g-l T ~ O , = average rate of conversion of ethylene to form carbon dioxide, scc min-’ g-1 UzpO2 U31)02312 arpEt (1 REto = relative rate of ethylene oxide formation a E P E t p 0 ~ ~ ’ ~~ C P C O ~ P OahpH~OP02~’~)~ ~~~~ Reo, = relative rate of carbon dioxide formation e, = fraction of surface covered by ith molecular species The values for the coefficients in this equation were esti2 , = sum of average product pressures, a t m mated by regressing 0 = superscript denoting value a t 8, = 0.01 atm 0 = subscript denoting value a t 2 , = 0.00 atm
+
+
+
+
+
+
+
Literature Cited
against
Benton, A. F., Drake, L. C., J . Amer. Chem. Soc., 54, 2186-94 Po2liZ,Po2, P023i2,PEt, PEtPO2li2,P O ~ ~ ~ ~and P CPOO ~~ ,~ / ~ P H ? O (1932). Bolme, D. W., PhD thesis, University of Washington, Seattle, Wash., 1957. These estimates are tabulated along with the corresponding Buntin, R. B., PhD thesis, Purdue University, Lafayette, Ind., t values of the estimates in Table VI. The values predicted 1961. Flank, W.H., PhD thesis, University of Delaware, Newark, Del., by these equations are plotted in Figures 2-5. 1 Q6.i The parameters estimated for a, were zero, indicating t h a t Hiri%t, P., J . Catal., 21, 56-65 (1971). ke
